Chemistry Notes 2010
Core Module 1: Production of Materials
Contextual Outline Humans have always exploited their natural environment for all their needs including food, clothing and shelter. As the cultural development of humans continued, they looked for a greater variety of materials to cater for their needs. The twentieth century saw an explosion in both the use of traditional materials and in the research for development of a wider range of materials to satisfy technological developments. Added to this was a reduction in availability of the traditional resources to supply the increasing world population. Chemists and chemical engineers continue to play a pivotal role in the search for new sources of traditional materials such as those from the petrochemical industry. As the fossil organic reserves dwindle, new sources of the organic chemicals presently used have to be found. In addition, chemists are continually searching for compounds to be used in the design and production of new materials to replace those that have been deemed no longer satisfactory for needs. This module increases students’ understanding of the implications of chemistry for society and the environment and the current issues, research and developments in chemistry. 1.1 Construct word and balanced formulae equations of all chemical reactions as they are encountered in this module: • Acid reactions: o acid (aq) + base (aq)  salt (aq) + water (l) o acid (aq) + active metal (s)  salt (aq) + hydrogen (g) o acid (aq) + metal carbonate (s)  salt (aq) + water (l) + carbon dioxide (g) • Cracking of long chain alkanes (e.g. decane) to give a shorter chain alkene & alkane o Thermal cracking (heated with steam in absence of O2 to 750-900°C) Initiation: C10H22  2C5H11· Propagation: C5H11·  C3H7· + C2H4 Termination: 2C3H7·  C6H14 o Catalytic cracking zeolite (heated in absence of O2 to 500°C)

C10H22 (l)  C2H4 (g) + C8H18 (l) • • Bromine water Br2 (l) + H2O (l) ↔ HOBr (aq) + HBr (aq) Cyclohexene + bromine water addition reaction C6H10 (l) + Br2 (aq) + H2O (l)  C6H10BrOH (l) + HBr (aq)

1

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Alkane substitution reactions hv C 2 H 6 + X 2  C 2 H 5 X + HX X = F, Cl, Br, I; hv = light • Addition reactions of ethylene o Hydrogenation:
Ni or Pt

C 2 H 4 (g) + H 2 o Halogenation:

(g)

C 2 H 6 (g)

Ni or Pt

C 2 H 4 (g) + Cl 2 (g)  C 2 H 4 Cl 2 (l) o Hydrohalogenation:
Ni or Pt

C 2 H 4 (g) + HF (g)  C 2 H 5 F (g) o Oxidation:
Ag

2C2H4 (g) + O2 (g)  2C2H4O • Addition polymerisation of ethylene

(ethylene oxide, high T & P)

(Initiator driven: activation, initiation, propagation, termination – high T & P, LDPE) (Catalyst driven: Ziegler-Natta – low T & P, HDPE)

•

Condensation polymerisation o alcohol (l) + carboxylic acid (l)  alkyl alkanoate (l) + water (l) methanol (l) + ethanoic acid (l)  methyl ethanoate (l) + water (l)
H2SO4

•

Formation of a polyester o dicarboxylic acid + diol  ester dimer + water

•

Dehydration of ethanol: conc. H 2 SO 4 or H 3 PO 4

C 2 H 5 OH (l) ↔ C 2 H 4 (g) + H 2 O (l) • Catalytic hydration: dilute H 2 SO 4

C 2 H 4 (g) + H 2 O (l) ↔ C 2 H 5 OH (l)
(indirect: via reaction mechanism with dilute H 2 SO 4 , mild T & P) (catalytic: H 3 PO 4 or tungstic acid impregnated into zeolite, high T & P)

2

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Photosynthesis: chlorophyll 6CO 2 (g) + 6H 2 O (l)  C 6 H 12 O 6 (aq) + 6O 2 (g) • Fermentation of glucose: yeast C 6 H 12 O 6 (aq)  2C 2 H 5 OH (aq) + 2CO 2 (g)
(~37°C, anaerobic, alcohol concentration Ca > Na > Mg > Al > Zn > Fe > Pb > H2 > Cu > Ag > Au o Active metals are strong reductants, inactive metals are weak reductants o Oxidants: Au3+ > Ag+ > Cu2+ > H+ > Pb2+ > Fe2+ > Zn2+ > Al3+ > Mg2+ > Na+ > Ca2+ > K+ o Cations of inactive metals are strong oxidants, cations of active metals are weak oxidants o Displacement reaction will only occur if the metal & metal ion can successfully donate & accept electrons 4.2 identify the relationship between displacement of metal ions in solution by other metals to the relative activity of metals

•

The greater the difference in activity between the two metals, the more vigorous the displacement reaction o Order of decreasing reductant strength – ease of oxidation o Strongest reductants at top right, strongest oxidants at bottom left o Reversible arrows – direction of reaction can be reversed under different circumstances o A spontaneous redox reaction will occur if the reductant is higher in the reduction half-equation table than the oxidant o A more active metal will displace a less active metal ion from solution

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4.3 Account for changes in the oxidation state of species in terms of their loss or gain of electrons • Oxidation state: a number given to an atom to indicate (theoretically) the number of electrons it has lost or gained (that is, its state of oxidation) also called oxidation number o Used to describe the oxidation & reduction processes in redox reactions o In a redox reaction the OS of the reductant increases and the OS of the oxidant decreases o Elements assigned oxidation states (written as Roman numerals) according to assumption that all bonds in compounds are ionic – although it is not true for many compounds  Oxygen has an oxidation state of –II (except in peroxides (O22-) where its OS = -I)  Hydrogen has an oxidation state of +I (except in metal hydrides (e.g. NaH) where its OS = -I) o Using these rule the oxidation states of elements in compounds can be found o Thus it can be identified whether an element is oxidised or reduced when not obvious

o An increase in oxidation number means that electrons have been lost – oxidation number moving towards the “positive” numbers due to the loss of negative electrons o A decrease in oxidation number means that electrons have been gained – oxidation number is moving towards the “negative” numbers due to a gain of negative electrons. 4.4 Describe and explain galvanic cells in terms of oxidation/reduction reactions • Galvanic cell: an arrangement of electrodes & electrolytes in which a redox reaction causes a flow of electricity; also called an electrochemical cell • Electricity is simply a flow of electrons. Redox reactions are electron-transfer reactions; if this electron flow can be exploited, electricity can be produced. • In a galvanic cell, to utilise the electron flow, the redox reaction is physically split into its two half-reactions called half-cells: o Oxidation occurs at the anode (An Ox) – negative electrode in a galvanic cell o Reduction occurs at the cathode (Red Cat) – positive electrode in a galvanic cell  Each half cell contains an electrode in an electrolyte • A conducting wire and salt bridge connects the two half-cells and completes the circuit; as electrons have to flow from the oxidation cell to the reduction cell, a flow of electrons is produced in the wire, and hence electricity is produced. • Salt bridge: an electrolyte or electrolyte gel that joins two half-cells in a galvanic cell & allows movement of ions to maintain a balance of charges in the electrolytes; also called an ion bridge o Prevents build up of positive & negative charges in each half-cell o Solutions of nitrate salts (e.g. KNO3) commonly used – ions don’t form precipitates with other ions o Alternatively, a porous partition can be used through which ions can slowly diffuse • When the cell is complete & operating, electrons leave negative anode & travel to positive cathode through a conducting wire in response to a potential difference between the half-cells 31

o In internal circuit cations move towards reduction half-cell & anions move towards oxidation half-cell

4.5 Outline the construction of galvanic cells and trace the direction of electron flow

• • • • •

Two half cells are set-up, each containing an electrode in an electrolyte solution A wire connects the two electrodes to allow the electrons to flow from the anode to the cathode A salt bridge, e.g. filter paper soaked with KNO 3 is set up so that it dips into both electrolytes The salt bridge allows the flow of ions from one solution to the other Cell diagram – e.g. one half-cell is zinc electrode in zinc nitrate solution & other is copper electrode in copper (II) nitrate solution: Zn (s) | Zn(NO 3 ) 2 (aq) || Cu(NO 3 ) 2 (aq) | Cu (s) (salt bridge denoted by || ) e.g. Zn | Zn2+ and Cu | Cu2+ Nitrate ions are ‘spectator’ ions – involved in maintaining the charge balance in the electrolytes Oxidation half-cell shown at left side of cell diagram Negative terminal of voltmeter connected to anode & positive to cathode – positive voltage will then register

• Redox couple: the oxidant-reductant pair in a half-equation,
• • •

4.6 Define the terms anode, cathode, electrode and electrolyte to describe galvanic cells Word
Electrode Electrolyte

Definition
A metallic conducting plate which carries electric current into or out of a galvanic half-cell A substance that conducts an electric current by movement of its ions  Anolyte: the electrolyte present in the anode compartment  Catholyte: the electrolyte present in the cathode compartment The negative electrode in a galvanic cell where oxidation occurs The positive electrode in a galvanic cell where reduction occurs

Anode Cathode

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4.0.1 Perform a first-hand investigation to identify the conditions under which a galvanic cell is produced 4.0.2 Perform a first-hand investigation and gather first-hand information to measure the difference in potential of different combinations of metals in an electrolyte solution RISK ASSESSMENT:  KNO 3 is moderately toxic if ingested  Explosive mixtures with active metals such as Mg, Na, Al  Copper sulfate, zinc sulfate slightly toxic if ingested  Lead nitrate highly toxic by all routes of exposure, cumulative poison  Eye & skin protection, wash hands thoroughly after use of Pb  CuSO 4 , PbNO 3 , ZnSO 4 should be returned to lab assistant for precipitation or disposal  Use lead nitrate in small quantities in well ventilated area METHOD: 1. Half cell = beaker containing 50mL of 0.1 M solution of cations of metal ‘A’ and a ‘cleaned’ electrode made of metal ‘A’ – metals used: zinc, copper, lead 2. Salt bridge = U tube filled with KNO 3 (electrolyte) 3. A galvanic cell requires 2 half cells (connected by a salt bridge) with an external circuit (including a voltameter) connecting the 2 electrodes by alligator clips 4. More reactive metal (oxidised) connected to negative terminal of voltmeter, less reactive metal (reduced) connected to positive terminal, potential difference recorded RESULTS: Galvanic cell Potential difference (V) 2+ 2+ [ Zn | Zn || Cu | Cu ] 0.4 2+ 2+ [ Fe | Fe || Cu | Cu ] 0.5 3+ 2+ [ Al | Al || Cu | Cu ] 0.2 2+ 2+ [ Mg | Mg || Cu | Cu ] 0.95  Adding more electrolyte did not affect the results DISCUSSION:  Relative differences (rather than reference values) between half cells were apparent  Reliability could be improved through use of more precise digital multimeter  Validity – standard conditions require 1 M solutions & 25°C, used 0.1 M solutions & ambient temperature not 25°C  Metal strips cleaned of any oxide/coating  KNO 3 used as salt bridge electrolyte as it does not form precipitates with any of the electrolytes  Electrodes used had same dimensions to give good comparability CONCLUSION: A complete circuit is required for the operation of a galvanic cell 4.0.3 Gather and present information on the structure and chemistry of a dry cell or lead-acid cell and evaluate it in comparison to one of the following: button cell, fuel cell, vanadium redox cell, lithium cell, liquid junction photovoltaic device (eg the Gratzel cell) in terms of: chemistry, cost and practicality, impact on society, environmental impact • Batteries – portable sources of electrical energy, more than one cell o Not all electric cells are batteries, e.g. dry cells are ‘cells’ as they consist of only one galvanic cell o When batteries are connected in series a higher voltage is obtained 33

Structure

Lead-acid Cell • Secondary – recharged by application of an external current • Used to provide energy for car starter motors • Gradually recharged during driving using electrical energy from car’s alternator • Structure: 6 cells, each supplying 2V arranged in series to produce a 12V battery o Anode: porous lead sheet o Cathode : lead sheet coated in compressed lead (IV) oxide o Electrolyte: H2SO4 (35% w/w) o Electrodes in the form of a grid to increase surface area for electrode reactions o Thin, perforated fibreglass sheets separate each electrode

Oxygen-Hydrogen Fuel Cell (PEMFC) • Secondary – can be combined with electrolysers • Fuel cells – fuel (reductant) & the oxidiser (oxidant) are constantly replenished from a reservoir outside the cell • More modern proton exchange membrane fuel cells operate at much lower temps than early hydrogen-oxygen fuel cells • Structure o Cathode: platinum catalyst on graphitecoated, ridged paper o Anode: platinum catalyst on graphitecoated, ridged paper o Electrolyte: Fluorocarbon polymer with sulfonic acid functional groups

Chemistry

• •

Voltage = 12 V Anode half-equation: Pb (s) + SO42- (aq)  PbSO4 (s) + 2e• Cathode half equation: PbO2 (s) + SO42- + 4H+ + 2e-  PbSO4 (s) + 2H2O (l) • Life of battery limited by: o Lead sulfate disintegrating from electrode surface affects ability of battery to recharge as it ages o Slow corrosion of lead anode o Internal short circuiting • During discharge the density of electrolyte drops from 1.26 g/cm3 to 1.1 g/cm3 as H+ ions are consumed o Battery hydrometer can be used to measure density of electrolyte & this is used as a measure of the state of charge of the battery

• • • • • • • •

Voltage = 0.7 V Anode half-equation: H2 (g)  2H+ + 2eCathode half-equation: O2 (g) + 4H+ + 4e-  2H2O (l) Hydrogen diffuses to anode catalyst where it dissociates into protons & electrons Protons are conducted through membrane but electrons forced to travel through external circuit Oxygen molecules react on cathode catalyst with electrons & protons to form water The water produced by the cell reaction is ejected in the oxidant gas flow Pure hydrogen must be used otherwise the platinum catalyst is deactivated

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Cost &
Practicality

• • • • • •

Expensive due to lead content Lasts many years & can be recharged over & over again Can be recharged externally to the car using a suitable transformer Too-rapid recharging can cause explosive hydrogen gas to form Heavy – limits portability Lowest energy density of most commonly used rechargeable batteries

• • • • • • • • • •

Impact on Society

• • • • •

First rechargeable battery for commercial use Important battery for car start-up motors – provides high currents over short periods Useful storage battery for remote locations – can be recharged by connecting to solar panels or electric generators Useful for emergency lighting Reliable Recycled to retrieve lead – toxic to organisms in the environment, causes anaemia in humans Electrolyte (sulfuric acid) is highly acidic & can cause severe damage if spilled Sealed lead-acid cells prevent acid fumes from causing corrosion

• • • • • • •

Environm ental Impact

• • •

Expensive due to high developmental costs Only a little thermal shielding is required to protect personnel as the cell operates at 85°C Could be ideal for commercial applications Compact, lightweight, no major moving parts Use of a solid polymer electrolyte eliminates corrosion & safety concerns associated with liquid electrolyte fuel cells Water must be evaporated at same rate as it is produced Steady ratio of fuel & oxidant required for efficiency, same temp must be maintained Storage may be a risk because the cell uses hydrogen Higher power density than many other fuel cells Less efficient than lead-acid cell at recharging but better suited for long term storage of power Not likely to be a short-term impact as this is a developing technology The long operating lifetime (in excess of 50 000 hours) is impressive May be one of the technologies that eventually replaces fossil fuel energy sources Instant start-up, reliable Used as power sources in remote locations – e.g. spacecrafts, remote weather stations, In the long term such a cell that utilises nonfossil fuels may be an important power source No pollutants produced during operation – only water produced

4.0.4 Solve problems and analyse information to calculate the potential Eө requirement of named electrochemical processes using tables of standard potentials and half-equations • The total voltage (or electromotive force) a galvanic cell can produce is determined by the substances taking part in the redox reaction • Electrode/electrolyte couples have a FIXED voltage, no matter how many moles of each substance is present – reduction potential • Standard electrode potentials – measured under standard conditions  Electrolyte concentration = 1.0 mol/L, standard temp (25°C) and pressure (100 kPa) o Standard hydrogen half-cell (Pt, H2 | H+) – standard reference half-cell to compare other half-cells  Inert platinum electrode covered in ‘platinum black’ – highly porous form of platinum metal powder 35

Electrode placed in 1.0 mol/L solution of H+ ions H2 gas bubbled over surface of electrode under standard temperature & pressure Reduction half-equation: H+ (aq) + e½H2 (g) Assigned a half-cell potential of zero volts Thus when a test half-cell is connected to the hydrogen half-cell in a galvanic cell the measured voltage = the half-cell potential of the test couple o Cell potential (Eө): the difference between the electrode potentials of the half-cells of a galvanic cell; also called voltage  Eө = Eө reduction + Eө oxidation  Positive or negative depending on whether test cell acts as site of oxidation or site of reduction o Reduction potential: the potential of a reduction half-cell relative to the standard hydrogen electrode; also called electrode potential  Reduction potential table arranged with strongest oxidants in lower left & strongest reductants in upper right  Table used to predict whether redox reaction is spontaneous in forward direction  Predictions may not be accurate if standard conditions aren’t used o To predict a spontaneous redox reaction:  Reduction: examine table & select correct reduction half-equation. Record reduction potential for this half-equation  Oxidation: examine table & select oxidation half-equation by reversing appropriate reduction half-equation. Record oxidation potential by changing sign of the half-cell potential  Balance electrons in the half-equations by multiplying  Add the reduction half-equation & oxidation half-equation  Add the reduction potential and the oxidation potential to obtain the cell potential  A positive cell potential is indicative of a spontaneous redox reaction in the forward direction o The more positive the reduction potential the greater the tendency for reduction to occur o The more positive the oxidation potential the greater the tendency for oxidation to occur o Spontaneous redox reactions occur in the direction in which the cell potential is positive     

5. Nuclear chemistry provides a range of materials • Isotopes: atoms with the same atomic number but different mass number – same number of protons but different numbers of neutrons • In nuclear chemistry (chemistry dealing with nuclear reactions), isotopes are shown in the following form:

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•

A stable isotope is one that doesn’t disintegrate. An unstable isotope is one that continuously emits alpha, beta or gamma radiation from its nucleus in effort to become stable There are three types of radiation: α (alpha), β (beta) and γ (gamma) radiation: Alpha Particles (α) Helium nucleus ( 4 He)
2

Consist of: Charge Mass Ionising ability Penetration Deflection in electric fields

Beta Particles (β) 1 electron ( 0 e)
-1

Gamma Radiation (γ) High frequency EM radiation No charge No mass Poor Very good No deflection Accompany many nuclear decay reactions – means of shedding excess energy as nuclear particles rearrange o Release does not change conservation of atomic mass & mass numbers

2+ 4 amu Good Poor (2-10cm in air) Towards negative plate Ejected from heavy, unstable nuclei to remove a surplus of protons & neutrons o Tend to occur in nuclei with Z > 83 o Nucleus that forms may still be radioactive – further decay reactions will occur

-1 0.00055 amu Fair Fair (100cm in air) Towards positive plate Released from nucleus when a neutron decays into a proton and an electron o Occurs when n:p ratio is too high due to a surplus of neutrons o Mass number does not change o Atomic number increases by 1 Beta decay of zirconium-97

Example of decay

Alpha decay of radon-222

Gamma emission from cobalt-60

•

Positrons have same mass an electron but possess positive charge, form when a proton decays into a neutron • Electron capture: some unstable nuclei have a surplus of protons (n:p is too low) & to achieve stability they capture an inner shell electron & convert a proton into a neutron o E.g. beryllium-7 can transmute into litium-7 • Radioactive atoms decay at random – during this process, the nucleus emits particles and/or energy in older to attain a stable structure. The time for half of a radioisotope to decay is called the half-life – activity of the sample decreases by 50%. The decay of a radioisotope is a nuclear reaction. o In some cases the decay of a radioactive nucleus into a stable nucleus can occur in one step, in others (e.g. U) many decay steps involved before nucleus becomes stable o Atomic numbers & mass numbers are conserved

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5.1 Distinguish between stable and radioactive isotopes and describe the conditions under which a nucleus is unstable • Radioactivity is the spontaneous emission of RADIATION from certain atoms • The stability of the nucleus is related to the strength of the forces that hold the nuclear particles together o Nucleons – the protons & neutrons of the nucleus o Strong nuclear force – the force that holds the nucleons together  Operates over small nuclear distances and independent of charge  Much stronger than the electrostatic repulsion force between protons • There are 2 conditions used to predict whether an atom will be radioactive: o Atomic Number: All atoms (including all their isotopes) with Z > 83 (above bismuth) are radioactive as their nucleus is unstable o Proton-Neutron Ratio: The ratio of protons to neutrons determines whether an atom will be stable or not. Anything outside the ratios below is radioactive:  Z < 20, the stable ratio of protons to neutrons is 1:1  Around Z = 30, the stable ratio is about 1:1.3  Around Z ~ 73-80, the stable ratio is 1:1.5 5.2 Describe how transuranic elements are produced • Transuranic elements are elements with atomic numbers greater than uranium; that is Z > 92 • All transuranic elements are artificially produced: o Transuranic elements with Z = 93, 94 and 95, neptunium, plutonium, and americium are produced by neutron bombardment in a nuclear reactor. o Transuranic elements with Z > 95 are produced using particle accelerators. • Neutron Bombardment (in nuclear reactors): o In nuclear reactions, the fission chain-reaction (of uranium-235 or other elements) produces large amounts of neutrons. The nuclear reactor allows the uranium chain reaction to occur safely, releasing neutrons at a slow and controlled rate. A target is bombarded with neutrons to produce a radioactive species with extra neutrons. Thus, nuclear reactors produce neutron-rich isotopes. o Neutrons excellent projectiles as they have no electric charge so there is no problem in overcoming electrostatic repulsion with the target nucleus  E.g. unstable U-239 formed decays by beta emission to form neptunium-239 & plutonium-239 0 1 239 238 239 92 U + 0 n  92 U + γ  93 Np + −1 e
239 93

Np 

239 94

Pu +

0 −1

e

•

Fusion Reactions (in particle accelerators): o The production of larger transuranic elements is achieved by colliding heavy nuclei with high-speed positive particles (such as helium or carbon nuclei) o The positive particles are accelerated to near the speed of light in particle accelerators using electric and magnetic fields in order to overcome the positive repulsive force of the heavy nuclei and fuse with them o Particle accelerators produce neutron-deficient isotopes.  E.g formation of Curium-243: 4 239 242 1 94 Pu + 2 He  96 Cm + 0 n 38

5.0.1 Process information from secondary sources to describe recent discoveries of elements • The first element to exist only in the lab was Technetium (atomic number 43), created in 1937 by bombarding Molybdnum with Deutenium nuclei. During and after WW2, an American team created 10 new elements using a particle accelerator, including neptunium, the first element heavier than uranium, and plutonium, the element used in the atomic bomb. • Since the mid 1970s, synthesis of ever-heavier new elements has depended on new generations of particle accelerators. • Four new elements have been discovered in the past 10 years. They are listed below. Note that their strange names are just temporary until the IUPAC decides on permanent names: o Ununhexium: Also known as “eka-polonium”, element 116 was synthesised in December, 2000, by the Joint Institute for Nuclear Research (Dubna, Russia). It was produced in a particle accelerator through the fusion of curium-248 and calcium-48. The atom decayed 48 milliseconds later. o Ununpentium: Also known as “eka-bismuth”, element 115 was synthesised in February, 2004, by the scientists from the Joint Institute for Nuclear Research (Dubna, Russia) and the Lawrence Livermore National Laboratory (America). It was produced through the fusion (partle accelerator) of americium-243 and calcium-48. The atom then underwent ALPHA decay, forming element 113, a new element o Ununtrium: Also known as “eka-thallium”, element 113 was also synthesised in February, 2004, through the alpha decay of ununpentium. o Ununoctium: Also known as “eka-radon”, element 118 is the most recently produced, and the heaviest element known to man. It was produced by the fusion of californium-249 atoms and calcium-48. 5.3 Describe how commercial radioisotopes are produced • Radioisotope: a radioactive isotope of an element; some radioisotopes are natural (e.g. carbon-14) whereas others are synthetic • Commercial radioisotopes are isotopes that are produced on a regular basis for medical, industrial or other use • Many are produced by neutron bombardment within nuclear reactors, as explained above; at the Lucas Heights nuclear reactor in Sydney, the Australian Nuclear Science and Technology Organisation (ANSTO) produces a range of neutron-rich isotopes for commercial use: o Technitium-99m (an important medical radioisotope) is produced by neutron bombardment of molybdenum-98 o Cobalt-60 (used in industry and medicine) is produced by neutron bombardment of the stable cobalt-59 o Americium-241 (a domestically used isotope; in smoke alarms) is produced by neutron bombardment of plutonium-241 • Other isotopes are produced in particle accelerators, such as the National Medical Cyclotron, near the Royal Prince Alfred hospital. Particle accelerators accelerate nuclei to incredible speeds, and which are then 39

collided with heavy nuclei. This produces neutron-deficient radioisotopes. Linear Accelerators accelerate particles in a straight line, while cyclotrons accelerate particles in a spiral. Radioisotopes produced include: o Iodine-131 (used to diagnose thyroid disorders) o Carbon-11, nitrogen-13, oxygen-15 (all used in PET scans) 5.4 identify instruments and processes that can be used to detect radiation • Detecting radiation – properties of nuclear radiation used to detect their levels o Radiation dangerous to humans so important to monitor exposure o Measured in several different units – the Becquerel (Bq) is a unit of radioactivity = one nuclear disintegration per second o Key properties used to detect nuclear radiation:  Darkens photographic film  Causes electron excitation into energy ‘traps’ in the lattice of certain crystals  Penetrates materials to different extents  Causes ionisation in gases & vapours  Deflected (or not deflected) by electric or magnetic fields • Geiger-Muller Counter: o Nuclear radiation called ionising radiation – able to ionise materials o GM tube & counter measures strength of ionising radiations o Beta radiation readily detected although can be designed to detect other types of radiation o Radiation enters GM tube through a mica window at one end o Inside tube is a low-pressure inert gas e.g. argon o The high-energy particles cause electrons to be ejected from the neutral atoms  Ar (g) + high energy particle/rays  Ar+ (g) + eo Inside tube is a cylindrical copper cathode & a central positive anode – high voltage maintained between these electrodes o Ionisation releases electrons which are attracted to the anode, stronger radiation = more ionisation o As electrons accelerate due to high voltage they cause more ionisations of gaseous atoms o Leads to a cascade of electrons that arrive at anode – amplified electrical pulse created at anode & detected by digital counter o Positive ions attracted to negative casing and accept electrons to complete the circuit

•

Radiation dosage badges – worn by workers in nuclear industry to monitor their radiation exposure o Film badges – monitor high-energy beta rays as well as gamma & X radiation  On one side of film is a sensitive (or ‘fast’) silver halide emulsion, on other a ‘slow’ emulsion  Low radiation doses will blacken the sensitive emulsion but not the slow emulsion  Higher doses will blacken the slow emulsion  Analysed regularly to assess radiation dose  Alpha rays have very low penetrating ability & are not normally monitored 40

o Thermoluminescent dosimeters (TLDs) – certain crystals (e.g. Al2O3 or LiF) absorb strong beta or gamma radiation  This absorbed energy causes electrons to be excited to higher energy states where they become trapped in ‘trapping centres’  Radiation dose determined by releasing the trapped electrons using laser radiation or heating  Visible photons of light emitted & their energy is related to the ionising radiation dose  Also produced as finger badges to record doses to the hands  Have an open window that allows detection of beta rays which won’t penetrate plastic casing • Cloud chamber – contains a cold saturated vapour (e.g. alcohol) o As ionising radiation travels through the air & vapour it ionises air molecules o Vapour molecules condense onto these ions creating small droplets or cloud trails that reveal path of ionising radiation o Alpha particles strongly ionising but have low penetration – form thick but short cloud trails o Beta particles less ionising but more penetrating – form thinner, longer trails that may show some zig-zag effects o Gamma rays – very weakly ionising but highly penetrating – form long, wispy trails o Magnetic or electric fields can also be used with cloud chambers – positive alpha particles deflected in different directions from those of negative beta particles, gamma rays not deflected

5.5 identify one use of a named radioisotope: in industry, in medicine • Industry: Cobalt-60 to inspect metal parts & welds for defects • Medicine: Technitium-99m is used to pin-point brain tumours 5.6 describe the way in which the above named radioisotopes are used and explain their use in terms of their chemical properties 60 • Industry: 27 Co  60 Ni + −0 e + γ 1 28 o Cobalt-60 is used in industrial radiography to inspect metal parts and welds for defects  Beams of radiation are directed at the object to be checked  Radiographic film on the other side of the metal part detects the amount of radiation passing through the metal part  More radiation will pass through if there are cracks, hence making the film darker in sections. 41

o Co-60 is used because it emits gamma rays which penetrate metal parts. It has a half life of 5.3 years and can be used in a chemically inert form. This enables the equipment to have along ‘lifetime’ and not require regular maintenance. m • Medicine: 9943 Tc  99 Tc + γ 43 o Technitium-99m (Tc in a metastable state – excited but not highly unstable) is used in over half the current nuclear medicine procedures as a tracer/diagnosing tool – e.g. pin pointing brain tumours  The Tc-99m is attached to a biological molecule that concentrates in the organ being investigated. o Tc-99 can be changed into a lot of different oxidation states, enabling the production of a wide range of biologically active chemicals o Tc-99m is used because it has a short half life of 6 hours so it quickly leaves the patients system. It emits low energy gamma radiation that minimises damage to healthy tissue. It is reasonably reactive and can form a compound with chemical properties that leads to concentration in the organ of interest. 5.0.2 Use available evidence to analyse benefits and problems associated with the use of radioactive isotopes in identified industries and medicine • BENEFITS of radioisotopes: o Industrial benefits include the ability to make monitoring equipment that is more sensitive, precise and reliable than previously possible. It allows for more efficient processes (such as sterilisation and food irradiation) and previously impossible things (examining faults in construction and machinery) o Medical benefits include a new wide range of non-invasive diagnostic techniques that would not be possible on sensitive organs (brain, etc). Radiation therapy is also a greatly positive new treatment. • PROBLEMS with radioisotopes: o Nuclear reactors, which are the source of neutrons, produce considerable amounts of nuclear waste, which we have no way of disposing safely, and which last for thousands of years. o The storage of radioactive material presents a problem, as they must be kept in shielded containers to prevent radiation leakages. o Doses of radiotherapy must be extremely carefully controlled, to balance between the benefits of killing cancer cells, and the risk of harm. o Nuclear technicians and other workers must be continually protected and avoid any form of irradiation, as disease such as cancer or radiation poisoning can result from this.

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Chemistry Notes 2010
Core Module 2: The Acidic Environment
• • • Acidic oxides of non-metals: SO3 (g) + 2NaOH (aq)  Na2SO4 (aq) + H2O (l) Basic oxides of metals: CuO (s) + 2HCl (aq)  CuCl2 (aq) + H2O (l) Amphoteric oxides: o Al2O3 basic oxide: o Al2O3 acidic oxide: • Al2O3 (s) + 6HCl (aq)  2AlCl3 (aq) + 3H2O (l)

Al2O3 (s) + 2NaOH (aq)  2NaAlO2 (s) + H2O (l) Solubility of carbon dioxide in water CO2 (g) ↔ CO2 (aq) CO2 (aq) + H2O (l) ↔ H2CO3 (aq) H2CO3 (aq) ↔ H+ (aq) + HCO3- (aq)

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Sulfur dioxide o Natural source (bacterial decomposition): 2H2S (g) + 3O2 (g)  2SO2 (g) + 2H2O (g) o Artificial source (roasting of zinc sulfide ore): 2ZnS (s) + 3O2 (g)  2ZnO (s) + 2SO2 (g) Oxides of nitrogen o Natural & artificial sources (lightning & high temperature combustion): N2 (g) + O2 (g)  2NO (g) o Slow oxidation of NO 2NO (g) + O2 (g)  2NO2 (g) Formation of acid rain SO2 (g) + H2O (l)  H2SO3 (aq) 2H2SO3 (aq) + O2 (g)  2H2SO4 (aq) SO3 (g) + H2O (l)  H2SO4 (aq) 2NO2 (g) + H2O (l)  HNO2 (aq) + HNO3 (aq) 2HNO2 (aq) + O2 (g)  2HNO3 (aq) o Chemical weathering of marble CaCO3 (s) + H2SO4 (aq)  CaSO4 (s) + H2O (l) + CO2 (g)

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Hydrochloric acid HCl (g) + H2O (l)  H3O+ (aq) + Cl- (aq) Sulfuric acid H2SO4 (l) + H2O (l)  H3O+ (aq) + HSO4- (aq) HSO4- (aq) + H2O (l) ↔ H3O+ (aq) + SO42- (aq)

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Acetic (ethanoic) acid CH3COOH (aq) + H2O (l) ↔ H3O+ (aq) + CH3COO- (aq)

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Citric acid (2-hydroxypropane-1,2,3-tricarboxylic acid – C6H8O7) C3H5O(COOH)3 (aq) + H2O (l) ↔ H3O+ (aq) + C3H5O(COOH)2COO- (aq) C3H5O(COOH)2COO- (aq) + H2O (l) ↔ H3O+ (aq) + C3H5O(COOH)(COO)22- (aq) C3H5O(COOH)(COO)22- (aq) + H2O (l) ↔ H3O+ (aq) + C3H5O(COO)33- (aq) Acidic salts: o strong acid + weak base  acidic salt + water NH4OH (aq) + HCl (aq)  NH4Cl (aq) + H2O (l) NH4+ (aq) + H2O (l) ↔ NH3 (aq) + H3O+ (aq) Basic salts: o weak acid + strong base  basic salt + water 2NaOH (aq) + H2CO3 (aq)  Na2CO3 (aq) + 2H2O (l) CO32- (aq) + H2O (l) ↔ HCO3- (aq) + OH- (aq)

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Neutral salts: o strong acid + strong base  neutral salt + water HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)

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Amphiprotic substances o Water as an acid: H2O (l) + OH- (aq)  OH- (aq) + H2O (l) o Water as a base: H2O (l) + H3O+ (aq)  H3O+ (aq) + H2O (l) o HCO3- as an acid: HCO3- (aq) + OH- (aq)  CO32- (aq) + H2O (l) o HCO3- as a base: HCO3- (aq) + H3O+ (aq)  H2CO3 (aq) + H2O (l)
(NaHCO3 used to neutralise both acidic & basic spills)

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Natural buffer – carbonic acid/hydrogen carbonate ion equilibrium for maintenance of pH of blood H2CO3 (aq) + H2O (l) ↔ HCO3- (aq) + H3O+ (aq)

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Esterification o alkanol (l) + alkanoic acid (l) ↔ alkyl alkanoate (l) + water (l) CH3OH (l) + CH3COOH (l) ↔ CH3OOCCH3 (l) + H2O (l) methanol ethanoic acid methyl ethanoate water + H2SO4  + H2O
H2SO4

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Esterification practical CH3COOH (l) + C4H9OH (l) ↔ acetic acid 1-butanol
H2SO4

CH3COOC4H9 (aq) + H2O (l) 1-butyl ethanoate water

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1. Indicators were identified with the observation that the colour of some flowers depends on soil composition. 1.1 Classify common substances as acidic, basic or neutral • Substances can be grouped according to their ability to change the colour of a natural dye or acid-base indicator o Acid-base indicator: a solution of a pigment or dye that changes colour in the presence of acids & bases o Historically, classification was related to other observable properties of these substances, e.g. taste & ability to attack & corrode other materials such as iron or limestone Acidic Vinegar (ethanoic acid) Lemon and citric juices Stomach juices (HCl) Soda water Car battery acid Lactic acid Aspirin Vitamin C (ascorbic acid) Neutral Pure water Glucose solution NaCl solution (salt water) Alcohol-water solutions Lactose solution Basic Drain cleaner Washing soda solution Baking soda solution Ammonia solution Lime water Oven cleaners Antacid tablets

1.2 – Identify that indicators such as litmus, phenolphthalein, methyl orange and bromothymol blue can be used to determine the acidic or basic nature of a material over a range, and that the range is identified by change in indicator colour • pH: a logarithmic scale used to measure the acidity or alkalinity of a solution o acidic solution < 7, 7 = neutral solution, basic solution > 7, • In addition to natural plant extracts, a wide range of synthetic acid-base indicators have been developed o Allow chemists to determine the range of acidity or basicity within very narrow limits Methyl orange – changes in acidic range (3.1-4.4) from red through orange to yellow Litmus – changes in neutral range (5.0-8.0) from red through purple to blue Bromothymol blue – changes in neutral range (6.0-7.6) from yellow through green to blue Phenolphthalein – changes in basic range (8.3-10.0) from colourless through pink to red

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1.3 – Identify and describe some everyday uses of indicators including the testing of soil acidity/basicity • Acid-base indicators have the following everyday uses: o Testing the acidity of soil – some plants prefer slightly acidic soil, others prefer slightly alkaline soils  Soils too acidic can be partly neutralised by adding bases such as crushed calcium carbonate or dolomite  Soils too basic can be partly neutralised with fertilisers such as sulfate of ammonia or by adding rotting plant material (compost) & manures containing natural acids  Soil pH can be measured using electronic instruments, universal indicator or suitable narrow-range indicators  Test tube is one-third filled with soil, distilled water added, tube stoppered & shaken & soil allowed to settle  Small amount of white barium sulfate suspension can be added to aid settling process  Pasteur pipette used to withdraw a sample of the supernatant water into 2 clean tubes  Universal indicator added to 1st tube & colour compared with a pH chart  A narrow range indicator then selected to test the other sample to determine a more accurate soil pH o Chemical research – used to:  Determine whether solutions are acidic, basic or neural  Monitor the change in acidity during volumetric analysis when an acid is being used to neutralise a base o Testing the acidity of water – universal indicator suitable for field testing water samples as its extensive colour range allows small changes in acidity to be noted  Acidity levels must be monitored in swimming pools – sodium hypochlorite (NaOCl) added to kill microbes  Hypochlorite ion reacts with water to produce unstable hypochlorous acid (HOCl) & hydroxide ions OCl- + H2O (l)  HOCl + OH HOCl is the active form which kills microbes but is relatively unstable  Hydroxide ions raise basicity of the water, HCl added to return water to near neutrality (pH 7.2-7.6)  If water too acidic will cause irritations to the eyes & skin & attack metal fittings of pool circulation system  If water too basic green algal scum will grow in the pool  At pH of 7.5 ratio of active HOCl to inactive OCl- ≈ 1:1 & maximum chlorination efficiency is achieved  Samples of pool water tested using a pool test kit – phenol red used, colours compared with colour chart o Waste water monitoring – ensuring alkaline solutions are neutralised before discharged into sewer. o Testing water in aquariums – fish are sensitive to the pH of water so must be controlled and maintained

1.0.1 – Perform a first-hand investigation to prepare and test a natural indicator In the laboratory, various commercial indicators can be used e.g. litmus, methyl orange, bromothymol blue, phenolphthalein and universal indicator. However natural indicators can also be used in the laboratory e.g. cabbage leaves, flower petals. You may have to chop these finely, boil them and drain of the liquid. This acts as an indicator. We did an experiment with the red cabbage leaves. METHOD:  Gather red cabbage leaves as your natural indicator and chop them.  Put into 800mL beaker and add 500mL distilled water to red cabbage leaves in beaker.  Heat using Bunsen burner for 10 minutes, continuously stirring with stirring rod.  Draw up purple solution (natural indicator) using small pipette and add 5 drops into 5 test tubes.  Add 3 drops 1M HCl into test tube 1, 3 drops 2M HCl into test tube 2, 3 drops distilled water into test tube 3, 3 drops 1M NaOH into test tube 4 and 3 drops 2M NaOH into last test tube  Observe colour changes and record this change in colour.  Repeat experiment 10 times for reliable results. RESULTS:  Cabbage juice (originally purple) had a pH range of 5-8 with it been green in basic solution and red in acidic solution 4

1.0.2 – Identify data and choose resources to gather information about the colour changes of a range of indicators Indicator Methyl orange Methyl red Bromothymol blue Litmus Phenol red Phenolphthalein Universal indicator pH range 3.1 ↔ 4.4 4.4 ↔ 6.0 6.0 ↔ 7.6 5.0 ↔ 8.0 6.8 ↔ 8.4 8.3 ↔ 10.0 0.0 ↔ 14.0 Colour change red → orange → yellow pink → yellow yellow → green → blue red → purple → blue yellow → red colourless → pink → red red → yellow → green (neutral) → blue/violet

1.0.3 – Solve problems by applying information about the colour changes of indicators to classify some household substances as acidic, neutral or basic A problem with using indicators to test the pH of a green detergent is that the colour of the detergent would mask any colour change of the indicator. This problem can be overcome by diluting the detergent or using a pH meter or pH probe. 2. While we usually think of the air around us as neutral, the atmosphere naturally contains acidic oxides of carbon, nitrogen and sulfur. The concentrations of these acidic oxides have been increasing since the Industrial Revolution 2.1 – Identify oxides of non-metals which act as acids and describe the conditions under which they act as acids Oxides are chemical compounds formed when an element reacts with oxygen, e.g. MgO, P4O6, P4O10 Some oxides of non-metals can act as acids when dissolved in water or when neutralising bases When oxides of phosphorous are dissolved in warm water containing universal indicator the mixture turns from green to red thus phosphorus oxides are acidic P4O10 (s) + 6H2O (l)  4H3PO4 (aq) o Presence of water is essential o The phosphoric acid releases hydronium (H3O+) ions that cause the change in colour of the indicator • Some oxides are not soluble in water but can act as acids or bases in neutralisation reactions acid + base  salt + water o Salt: the ionic compound formed when a base is neutralised by an acid, metal ion of base named first followed by name of anion of parent acid, e.g. Barium hydroxide + phosphoric acid  barium phosphate + water 3Ba(OH)2 (s) + 2H3PO4 (aq)  Ba3(PO4)2 + 6H2O (l) Acidic oxides Basic oxides Amphoteric oxides Oxides that dissolve & react with Oxides that dissolve & react Oxides that exhibit both acidic & basic strong basic solutions, e.g. sulfur with strong acids, e.g. copper (II) properties trioxide oxide, magnesium oxide o React with both strong acids & SO3 (g) + 2NaOH (aq)  Na2SO4 (aq) + H2O (l)) CuO (s) + 2HCl (aq)  CuCl2 (aq) + H2O (l) strong bases, e.g. aluminium oxide MgO (s) + H2O (l)  Mg(OH)2 (aq) Al2O3 basic oxide:
Al2O3 (s) + 6HCl (aq)  2AlCl3 (aq) + 3H2O (l))

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Al2O3 acidic oxide:

Al2O3 (s) + 2NaOH (aq)  2NaAlO2 (s) + H2O (l)) or Al2O3 (s) + 2NaOH (aq) + 3H2O (l) 2NaAl(OH)4 (same as NaAlO2 when aq)

(aq)

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o Generalisations:  Most metallic oxides are basic oxides  Most non-metallic oxides are acidic oxides  Some oxides are neutral  Some oxides have both basic & acidic properties Acidic oxides Cl2O7 SO3 SO2 NO2 CO2 Basic oxides Na2O K2O CaO BaO MgO Amphoteric oxides Al2O3 ZnO SnO PbO Sb2O3 Neutral oxides CO NO N2O

2.2 – Analyse the position of these non-metals in the Periodic Table and outline the relationship between position of elements in the Periodic Table and acidity/basicity of oxides The metals of Groups I & II all form basic oxides Most non-metals (other than the noble gases) form acidic oxides Across a period: decreasing basicity  amphoteric  increasing acidity until group VII Down a group basicity of oxides increases with increasing metallic character The non-metallic oxides with the highest oxidation states in the non-metal tend to be more acidic  E.g. SO3 is more acidic than SO2 • In many cases an element may have a basic or acidic or amphoteric oxide due to the different oxidation states of the element  The higher the oxidation state of the metal or semi-metal, the more amphoteric or acidic the oxide • Some oxides behave in different ways in different chemical environments so have more than one classification such as amphoteric to basic) o Related to the decreasing metallic character of the elements across each period o Also related to the type of crystal lattice each oxide forms – differences in bonding affect the behaviour of each oxide in water  The basic oxides of period 3 (Na2O & MgO) are ionic compounds  Al2O3 lattice is also ionic and the compound in quite insoluble in water  SiO2 lattice is covalent network & is very insoluble in water  Remaining oxides form covalent molecular lattice  High solubility of Na2O allows oxide ion to interact with water & form OH- ions  High solubility of SO3 allows formation of acidic H+ ions  The two oxides that are very insoluble in water (Al2O3 & SiO2) exhibit amphoteric & weakly acidic properties respectively, only in the presence of strong acids & bases • • • • •

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2.3 – Define Le Chatelier’s principle “If a system is at equilibrium & a change is made that disturbs the equilibrium, then the system responds in such a way as to counteract the change & eventually a new equilibrium is established”. • • Many reactions don’t proceed to completion – final reaction mixture consists of both products & reactants Chemical equilibria are dynamic & not static o E.g. in a saturated solution rate of dissolution = rate of crystallisation o Chemical reactions are reversible – shown by using reversible arrows  Arrow pointing right designates forward reaction & arrow pointing left designates reverse reaction • All systems in chemical equilibria exhibit the following characteristics: o The system is closed – atoms, molecules or ions cannot leave or enter the system o The observable or measurable properties (macroscopic properties) of the system are constant – e.g. colour, electrical conductivity, concentration, gas pressure o The concentration of reactants & products are constant once equilibrium is achieved o The rate of the forward reaction equals the rate of the reverse reaction, thus the system is dynamic 2.4 – Identify factors which can affect the equilibrium in a reversible reaction • The factors that may cause a disturbance to a chemical system in equilibrium are: o Temperature – reactions classified as endothermic or exothermic on the basis of whether heat is absorbed or released in the forward reaction, using Le Chatelier’s principle:  Endothermic equilibria: an increase in temperature causes the equilibrium to shift to favour the products  Exothermic equilibria: an increase in temperature cause the equilibrium to shift to favour the reactants o Concentration  Increasing the concentration of reactants or decreasing the concentration of products shifts the equilibrium to favour the products  Decreasing the concentration of reactants or increasing the concentration of products shifts the equilibrium to favour the reactants o Gas partial pressure – in a mixture of gases, the total gas pressure is the sum of the individual pressures of each component gas or the partial pressures of each gas  Partial pressure is therefore a measure of the concentration of the gas within the total gas mixture  Increasing the partial pressure of a reactant gas shifts the equilibrium to favour the products  Increasing the partial pressure of a product gas shifts the equilibrium to favour the reactants o Total gas pressure (in gaseous equilibria) – decrease in volume = increase in pressure  Increasing the total gas pressure shifts the equilibrium to favour the production of less gaseous particles  Decreasing the total gas pressure shifts the equilibrium to favour an increase in gaseous particles o Catalysts do not alter the position of a system already in equilibrium  If a system is not at equilibrium catalysts reduce the time to reach equilibrium  Final equilibrium position is the same whether the catalyst is present or not • In equilibrium systems that involve solid or liquid species, changing the quantity of the solid or liquid present has no effect on the position of equilibrium (‘concentrations’ of solids & liquids unchanged) 2.5 – Describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and explain in terms of Le Chatelier’s principle • CO2 is an acidic oxide, dissolves in water to produce a solution of carbonic acid (H2CO3) o Carbonic acid is in equilibrium with H+ ions & CO32- ions as shown in the following reactions

o The dissolution reaction is exothermic 7

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Carbonated soft drinks are manufactured by supersaturating water with CO2 under pressure o System will remain this way only if bottle is sealed so as to maintain high CO2 gas pressure under the cap  Increase in CO2 gas pressure shifts equilibrium (1) right to reduce gas pressure = more CO2 dissolved  Pushes equilibrium (2) to the right to make more H2CO3 (aq)  Increase in [H2CO3] shifts equilibrium (3) right thus acidity increases as more H+ ions are formed - Thus soda water tastes sour or tart – pH ~ 4 o Changing the pressure – when the cap of a soft-drink bottle is unscrewed there is rapid effervescence  Soft drink eventually goes ‘flat’ & tastes less sour due to the reversal of these equilibria  Escape of gaseous CO2 shifts equilibrium (1) to the left in the open system, other 2 equilibria shift left o Changing the pH – e.g. addition of an acid increases [H3O+], equilibria shift to the left to counteract change, soda water degasses o Changing the temperature (at constant pressure) – e.g. warming soda water in a gas syringe, equilibria shift left to counteract change as reaction is exothermic, soda water degasses

2.6 – Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen SO2 • Acidic oxide • About 50% of SO2 that enters the atmosphere is derived from the oxidation of hydrogen sulfide (H2S) produced from bacterial decomposition 2H2S (g) + 3O2 (g)  2SO2 (g) + 2H2O (g) • Volcanic eruptions & geysers are major source • Formed from combustion of organic matter during bushfires NOx • NO & NO2 are most common atmospheric oxides of nitrogen found in urban air • NO is a colourless gas formed when N and O react at high temperatures o Neutral oxide • NO2 is a brown gas produced by oxidation of NO – acidic oxide o Rate of NO2 formation depends on concentration of NO in air o Up to 10% of total NOx in air is NO2 • Lightning is a major source of NOx • N2O (nitrous oxide) produced by bacteria acting on nitrogenous material in soils • Combustion of fuels such as coal & diesel oil • Produced in high temperature combustion o Fossil fuels contain small quantities of sulfide • In Sydney about 86% of total emissions of minerals (e.g. FeS2) NOx comes from engines of motor o These minerals are oxidised during combustion vehicles & other transport & SO2 is released • NOx derived from other sources including: • Metal smelters that convert sulfide minerals into o Industries metals are also a major source of SO2 pollution o Electrical power production o E.g. smelting of chalcopyrite (CuFeS2) during o Oil refining production of copper results in release of SO2 o Increased use of fertilizer increasing nitrogenous material in soils

Formation & Natural Sources

Industrial origins

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2.7 – Describe, using equations, examples of chemical reactions which release sulfur dioxide and chemical reactions which release oxides of nitrogen SO2 • Combustion of sulfide minerals in fossil fuels (e.g. FeS2) 4FeS2 (s) + 11O2 (g)  2Fe2O3 (s) + 8SO2 (g) • • Roasting zinc sulphide ore 2ZnS (s) + 3O2 (g)  2ZnO (s) + 2SO2 (g) Smelting of chalcopyrite (CuFeS2) during production of copper 2CuFeS2 (s) + 5O2 (g) + 2SiO2 (s)  2Cu (l) + 4SO2 (g) + 2FeSiO3 (l)

NO • Nitrogen and oxygen reacting in lightning strike to form nitric oxide N2 (g) + O2 (g)  2NO (g) NO2 • Slow oxidation of NO to form nitrogen dioxide 2NO (g) + O2 (g)  2NO2 (g) 2.8 – Assess the evidence which indicates increases in atmospheric concentration of oxides of sulfur and nitrogen • Monitoring SO2 & NOx levels o Strong evidence shows atmospheric concentrations of SO2 & NOx increased by 25% over last 200 years  Evidence from: - Analysis of trapped air bubbles in Antarctic ice core samples - Measurement of carbon isotopes in old trees - Increases in photochemical smog - Increased damage due to acid rain  However, difficult to obtain accurate evidence as concentration of sulfur & nitrogen oxides are present in very small concentrations (about 0.001 ppm) - Instruments capable of measuring these small quantities have only been available since the 1970s - SO2 & NO2 form SO42- & NO3- ions which are soluble in water & circulate in the biosphere & hydrosphere but don’t precipitate out in forms that can b readily studied  Although do not have accurate measurements in concentrations of atmospheric oxides of sulfur & nitrogen taken over a long period of time enough evidence to suggest significant increases in these concentrations have taken place, especially since the Industrial Revolution o During Industrial Revolution (1800s) large amounts of coal burnt – vast quantities of SO2 produced o Iron smelters generated large volumes of SO2 as they produced growing quantities of steel required for industry o The atmosphere of large industrialised cities in Europe & USA became highly polluted with acidic gases o Increased use of motor vehicles in 1900s increased oil consumption & production of nitrogen oxides o After many deaths in London in 1952 due to heavy acidic smogs, pollution controls were introduced o In 1970s the development of more sensitive gas analysis technologies allowed chemists to monitor the global increase in SO2 emissions due to the expansion of industries in Asia o Recently, air quality has improved in most westernised countries however due to increasing population & motor vehicle usage levels of pollutants have stabilised rather than continuing to decrease o Rapid industrialisation of Asia has led to huge increases in SO2 emissions – predicted that SO2 emissions in Asia will triple in the 20-year period from 1990 o In Western countries acidic oxide pollution is controlled whereas in some poorer nations or emerging economies such as China, the amount of gaseous acidic oxides entering the air is increasing 9

2.9 – Calculate volumes of gases given masses of some substances in reactions, and calculate masses of substances given gaseous volumes, in reactions involving gases at 0°C and 100kPa or 25°C and 100kPa Avogadro’s Law: Equal volumes of gases at the same temperature & pressure contain equal numbers of molecules Gas volumes vary as temperature & pressure change Molar weight varies from one substance to another whereas molar volume is constant for all gases Most of space occupied by a gas is empty & actual volume of each gas molecule is very small compared with the total gas volume Only when gases are highly compressed that this rule is not strictly true Molar volume (VM) of any gas has the following value at the stated temperatures & pressures At 0°C and 100 kPa: VM = 22.71 L/mol At 25°C and 100 kPa: VM = 24.79 L/mol The relationship between the number of moles (n) of a gas & its volume (V) is: n = V/VM E.g. The Apollo-13 spacecraft suffered an explosion on the outward journey to the moon in April 1970. Critical levels of carbon dioxide began to build up in the cabin. Above 2% CO2, the astronauts would feel nausea and giddiness. The astronauts constructed a temporary air purifier using lithium hydroxide to remove carbon dioxide from the air. Lithium hydrogen carbonate was formed as the carbon dioxide was absorbed. (a) Write a balanced equation for the reaction between carbon dioxide gas and lithium hydroxide. LiOH (s) + CO2 (g)  LiHCO3 (s) (b) Calculate the number of CO2 molecules that can be removed from the air by 100g of lithium hydroxide nCO2 = nLiOH = 4.176 mol MLiOH = 23.949 g/mol no. of carbon dioxide molecules = (4.176) x (6.022 x 1023) nLiOH = 100/23.949 = 2.51 x 1024 molecules = 4.176 mol (c) Determine the volume of this carbon dioxide at 25°C and 100kPa V = nVm = (4.176)(24.79) = 103.6 L 2.10 – Explain the formation and effects of acid rain • • When all gases are removed from pure water its pH is 7.0 (neutral) at 25°C Natural water contains dissolved gases including CO2 making the water weakly acidic – typical pH: 6.0-6.5 CO2 (g) + H2O (l) H2CO3 (aq) • Acid rain: rain that has a pH lower than 5 • Formation: o When atmosphere is polluted with acidic oxides such as SO2 & NO2 rainwater can become quite acidic (pH 4.0-5.0) due to high solubility of these gases in water o When oxides of sulfur & nitrogen dissolve they produce solutions of various acids  SO2 forms weak sulfurous acid (H2SO3) whereas SO3 produces strong sulfuric acid (H2SO4)

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NO2 produces weak nitrous acid (HNO2) and strong nitric acid (HNO3)

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Effects: o Chemically weathers marble statues & building facades  When calcium carbonate of marble is attacked by sulfuric acid in acid rain, the surface of the marble is converted into insoluble calcium sulfate CaCO3 (s) + H2SO4 (aq)  CaSO4 (s) + H2O (l) + CO2 (g)  The wet calcium sulfate crystallises to form a porous & crumbly mineral called gypsum (calcium sulfate dehydrate)  Over centuries soot & dust collect in pores of gypsum & turn the marble black o Attacks metallic structures composed of iron & steel – iron oxidised by H+ ions & is chemically weathered Fe (s) + H2SO4 (aq)  FeSO4 (aq) + H2 (g) o Devastating effect on many northern hemisphere forests – pine needles lose their waxy coating & turn brown, trees lose their foliage o Affects soil – acidified soils inhibit growth of plant seedlings  Basic minerals in soil (such as dolomite & limestone) are attacked & dissolved  Dissolves calcite (calcium carbonate) that holds sandstone grains together so causes significant chemical & weathering erosion  Mineral nutrients (e.g. K, Ca, Mg) required for plant growth removed when acid rain soaks into ground  Some insoluble minerals dissolved by the acidified water causing a release of toxic levels of metal ions, affecting plants & bacteria o Can significantly acidify lakes – populations of aquatic organisms become stressed  Presence of H+ ions interferes with CO2/CO32- equilibrium in the water – amount of dissolved CO2 in water drops as CO32- are removed – stresses photosynthetic organisms CO32- (aq) + 2H+ (aq) H2CO3 (aq) H2O (l) + CO2 (aq) CO2 (aq) CO2 (g)  Many aquatic invertebrates cannot reproduce in an acidic environment, fish eggs & fish die if too acidic  Creates high levels of toxic heavy metals from leaching of bedrock & soil – kills fish, affects food chain

2.0.1 – Identify data, plan and perform a first-hand investigation to decarbonate soft drink and gather data to measure the mass changes involved and calculate the volume of gas released at 25°C and 100kPa AIM: To decarbonate a soft drink & measure mass changes involved. To calculate the volume of gas released at 25°C and 100kPa Equipment: Electronic balance, 500mL beaker, glass rod, 1 bottle (unopened) soft drink, 6g salt PROCEDURE: 1. Unopened soft drink bottle weighed, clean dry beaker & glass rod were weighed 2. Soft drink carefully poured into beaker, approx 6g salt accurately weighed 3. Salt gradually added to beaker & stirred 4. After 10 mins beaker, glass rod, empty bottle & lid were reweighed DISCUSSION:  Validity of design questionable – uncertain whether solution was completely degassed, solution lost  No control used (e.g. pure water) to compare evaporative loss over time  Accuracy uncertain - % error reasonable but assumption made about temp & pressure  Reliability could be improved by repeating & averaging results  Allowing 30 mins, using a magnetic stirrer, larger quantities would reduce % error VARIABLES: time allowed (independent), quantity of soft drink, mass change (dependent), temperature & pressure, mass of salt used (controlled) 2.0.2 – Analyse information from secondary sources to summarise the industrial origins of sulfur dioxide and oxides of nitrogen and evaluate reasons for concern about their release into the environment • Industrial Origins: o SO2: burning of fossil fuels containing sulfide impurities, smelting of metal sulfide ores o NOx: high temperature combustion, e.g. internal combustion engines, power generation 11

Reasons for concern: o SO2 – irritating, poisonous gas – long term exposure has serious impact on respiratory system  Affects people with respiratory problems such as asthma & emphysema o NOx – irritating to skin, eyes, respiratory system, harmful in [low], toxic in [high]  Contributes to photochemical smog o Soil pH – acid deposition decreases pH of soils  Decreasing availability of some essential elements e.g. N, P  Increasing release of some elements to toxic levels e.g. Al  Stunts plant growth & decreases productivity o Acid rain  Natural environment – changes pH of water decreasing the survival of aquatic organisms, attacks waxy coating on leaves  Built environment – attacks marble and metal structures • Evaluation: o Increasing levels of SO2 & NOx are leading to permanent changes in the natural and built environment o Also have harmful effects on humans o Huge cause for concern – if left unchecked will cause further damage 3. Acids occur in many foods, drinks and even within our stomachs 3.1 – Define acids as proton donors and describe the ionisation of acids in water • Acid: a proton (hydrogen ion) donor • Base: a proton acceptor • Monoprotic acid: an acid that can donate one proton per molecule of acid • Diprotic acid: an acid that can donate two protons per molecule of acid • When an acid is placed in water it is ionised – covalent bond breaks, H+ ion that forms from the breakage donated to water molecule to form hydronium ion (H3O+) • The negative ion remains solvated by the water • Importance of water in ionisation process o When HCl is dissolved in a non-polar organic solvent (e.g. benzene or hexane) no ionisation occurs o Molecules of HCl remain, if active metal (e.g. Mg) added to these organic solutions there is no reaction o Not true for a water solution of HCl – therefore HCl is acidic only because of the presence of water o For this reason water is often called an ionising solvent • Alkalis – bases that are soluble in water & contain OH- ion, must be manufactured – not common in nature 3.2 – Identify acids including acetic (ethanoic), citric (2-hydroxypropane-1,2,3-tricarboxylic), hydrochloric and sulfuric acid • Hydrochloric acid – monoprotic acid, produced by passing HCl gas into water, conc. solution 11-12 mol/L HCl (g) + H2O (l)  H3O+ (aq) + Cl- (aq) o Strong acid – because all of HCl molecules ionise in water use straight arrow o Solution is a strong electrolyte because of the high concentration of ions – high electrical conductivity • Sulfuric acid – diprotic acid, concentrated solution around 18 mol/L, not quite as strong as HCl o Concentrated solution contains some unionised molecules o Ionisation effectively complete, particularly in dilute solutions o Ionises in 2 stages – 1st complete, 2nd effectively complete in only very dilute solutions ( in hot water • The pH (potential hydrogen) scale is used to figure out if a particular substance is acidic, basic or neutral o Compares the concentration of hydrogen ions in solutions o For most common purposes pH scale between 0 and 14  E.g. 1.0 mol/L HCL solution has pH of 0, 1.0 mol/L NaOH solution has pH of 14  Strong acid solution greater than 1 mol/L has pH < 0, strong base solution > 1 mol/L has pH > 14

o pH scale useful for comparing acidity and basicity of household substances & some laboratory chemicals

As H3O+ ions are added to the water, the system is no longer in equilibrium. The system readjusts to remove some of the added H3O+ ions and thus the concentration of OH- ions decreases

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3.4 – Describe acids and their solutions with the appropriate use of the terms strong, weak, concentrated and dilute • Strong acid: an acid that completely ionises in water solution, e.g. H2SO4, HCl, HNO3 • Weak acid: an acid that is incompletely ionised in water solution, equilibrium established, e.g. H2CO3, acetic, citric o Often less than 10% ionisation in water – most of solution consists of un-ionised molecules • Concentrated acid – an acid solution that has a high concentration of acid particles • Dilute acid – an acid solution that has a low concentration of acid particles • Strong & weak – describes degree of ionisation of acid molecules o Nothing to do with concentration – acids are strong or weak no matter how much water is added • Concentrated and dilute – describes amount of acid dissolved in the solution

A concentrated solution of a strong acid has a larger number of ions than a dilute solution of the same acid. A concentrated solution of a weak acid has a high number of un-ionised molecules and few ions. A dilute solution of a weak acid has fewer particles than the concentrated solution but the number of ions compared with molecules has increased slightly

3.5 – Identify pH as –log10 [H+] and explain that a change in pH of 1 means a ten-fold change in [H+] • The pH scale – logarithmic scale based on powers of 10

pH = -log10[H+] o o o o

[H+] = concentration of hydrogen ions in mol L-1

Or [H+] = 10-pH A change of 1 unit of pH is equivalent to a 10-fold change in the hydronium ion concentration A change of 2 units of pH is equivalent to a 100-fold change in the hydronium ion concentration In diprotic acids the molarity of the acid has to be doubled to find [H+] (assuming complete ionisation)

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3.6 – Compare the relative strengths of equal concentrations of citric, acetic and hydrochloric acids and explain in terms of the degree of ionisation of their molecules • Degree of ionisation o For the reaction: HA (aq) + H2O (l) H3O+ (aq) + A- (aq) degree of ionisation = [H3O+] / [HA] x 100% o Degree of ionisation of a weak acid is specified at a constant temperature (25°C) and at a known concentration of the acid HA o Acid concentration must be specified as dilution will affect equilibrium position – will shift to the right

E.g. Hypochlorous acid (HOCl) is a weak monoprotic acid. This acid is important in swimming pools as it kills microbes. In a 0.0010 mol/L solution its degree of ionisation is 0.54% a) Write a balanced equation for the ionisation of hypochlorous acid in water HOCl (aq) + H2O (l) ↔ H3O+ (aq) + OCl- (aq) b) Identify which species (HOCl or OCl-) is in higher concentration in water. Justify your answer. HOCl is in higher concentration as the degree of ionisation is very low (0.54%). This means that only 54 molecules in 10 000 will ionise. c) Calculate the concentration of hydronium ions in this 0.0010 mol/L HOCl solution Degree of ionisation = [H3O+]/[HOCl] x 100% [H3O+] = (Degree of ionisation)( [HOCl])/100 [H3O+] = (0.54)(0.0010)/100 = 5.4 x 10-6 mol/L • Comparing the degree of ionisation of acids of the same molarity

o HCl – strongest acid as it is completely ionised o Citric acid – weaker acid, degree of ionisation mainly due to step one of the three ionisation steps  Higher value for citric acid shows degree of ionisation (at first step) is greater than for acetic acid o Acetic acid – weakest acid, degree of ionisation less than citric acid o Weak acids incompletely ionised, difficult to compare directly as one is triprotic and the other monoprotic o Only strong monoprotic acids with a molarity of 0.01 mol/L will have a pH of 2 o Other two acids must be weaker as their pH is higher 15

3.7 – Describe the difference between a strong and a weak acid in terms of an equilibrium between the intact molecule and its ions • As mentioned above, hydrochloric acid is a strong acid because when it is dissolved in water all the HCl molecules ionise o 100% degree of ionisation, no equilibrium as it is a complete forward reaction. HCl (g) + H2O (l)  H3O+ (aq) + Cl- (aq) However a weak acid like acetic acid does not completely ionise in the water o Degree of ionisation is much lower (around 4% in 0.01 mol/L solution) – only a very small proportion of the molecules ionise in the water o Unlike in strong acids, the molecular form of weak acids (e.g. acetic acid) is in equilibrium with its ions and the equilibrium lies to the left (hence the double arrow in the equation below) CH3COOH (aq) + H2O (l) H3O+ (aq) + CH3COO- (aq)

•

3.0.1 – Solve problems and perform a first-hand investigation to use pH meters/probes and indicators to distinguish between acidic, basic and neutral chemicals AIM: To use pH meters and indicators to distinguish between acidic, basic and neutral chemicals METHOD: 1. Substances to test: HCl, H2SO4, NaCl, Na2CO3, NaHCO3, lemonade, orange juice, washing powder, rainwater, alcohol, 2. Calibrate pH meter, put each solution in a beaker, measure the pH of each using a pH meter or probe, repeat 3 times 3. Each substance was tested with universal indicator by adding 1 drop to each solution on a spot plate DISCUSSION:  Using a pH meter or probe is a non-destructive way of testing whether a chemical solution is acidic, basic or neutral  Using indicator solution is a destructive way of testing, as the indicator will contaminate the portion of solution tested  Validity – cleaned pH meter with distilled water before use, calibrated with substances of a known pH (buffers)  pH meter or probe is more precise than an indicator as it provides an actual pH reading rather than a range in which a pH falls, does not depend on judging the colour change. 3.0.2 – Plan and perform a first-hand investigation to measure the pH of identical concentrations of strong and weak acids AIM: To measure the pH of equimolar solutions of strong and weak acids. METHOD: 1. 0.0010 M solution of hydrochloric acid, 0.0010 M solution of acetic acid, 0.0010 M solution of citric acid 2. Place narrow range (pH 3-5) paper in clean beaker & add drops of solution, compare with colour chart 3. Calibrate pH meter, measure pH of solutions, compare with other groups & average results DISCUSSION:  HCl should give pH of 3 – compare  HCl most acidic (lowest pH), citric acid weaker, acetic acid weakest 3.0.3 – Gather and process information from secondary sources to write ionic equations to represent the ionisation of acids   Strong acids, when ionised, denoted by single arrow Weak acids denoted by reversible arrow, indicating that it will reach equilibrium – not completely ionised 16

3.0.4 – Use available evidence to model the molecular nature of acids and simulate the ionisation of strong and weak acids   Molecular model kits can be used to simulate the ionisation of strong and weak acids. Complete ionisation of strong acid – construct 4 HCl molecules & 4 molecules of water (with 2 lone pairs of electrons on O), remove hydrogen atom from each HCl and attach it to 1 lone pair of O  4 hydronium ions & 4 chloride ions left over Incomplete ionisation of weak acid – construct 4 HF molecules & 4 water molecules, remove only one hydrogen atom from one molecule of hydrofluoric acid and attach it to one O, 3 unionised HF molecules & one hydronium ion.

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3.0.5 – Gather and process information from secondary sources to explain the use of acids as food additives • o o o o Weak acids often added to food for: Inhibition of the growth of microbes such as bacteria and moulds (pH control) Prevention of spoilage by oxidation (antioxidant) Improvement of flavour of foods and drinks Leavening agent (leavening agents react with NaHCO3 to produce CO2 gas) Chemical formula CH3COOH HOOCCH2COH(COOH)CH2COOH H3PO4 CH3CH2COOH Information Used as vinegar (4% solution) to preserve foods (e.g. pickling); flavour enhancer Flavouring and preservative (anti-oxidant), especially in soft drinks; antacid ingredient Acidulation of soft drinks (particularly colas); manufacture of cottage & processed cheese; pH control in diet jellies Controls bacteria & mould growth, particularly in bread, potato crisps and cake mixes

Acidic food additive Acetic acid Citric acid Phosphoric acid

Propanoic acid

3.0.6 – Identify data, gather and process information from secondary sources to identify examples of naturally occurring acids and bases and their chemical composition Natural acid Citric acid Structural Formula Information Found in all citrus fruits and some vegetables

Hydrochloric acid Formic acid (methanoic acid) Malic acid

H–Cl HCOOH

Found in stomach acid Found in ant stings

Found in apples and cherries

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Common natural bases – acidic environment causes many bases to be neutralised o Some however are found in nature because they are insoluble in water & so have persisted over time Chemical formula CaCO3 CaCO3.MgCO3 Information Present in limestone & marble; limestone can be decomposed to form calcium oxide (lime) that is used to de-acidify soils Found as dolomite

Natural base Calcium carbonate

Calcium magnesium carbonate

3.0.7 – Process information from secondary sources to calculate the pH of strong acids given appropriate hydrogen ion concentrations E.g. An aqueous solution has a pH of 2.8 at 25°C. Calculate the hydronium ion concentration. pH = -log10[H+] [H+] = 10-pH = 10-2.8 [H+] = 1.58 x 10-3 mol/L E.g. An aqueous solution at 25°C has a hydronium ion concentration of 1.0 x 10-4 mol/L. Calculate the pH of the solution. The hydronium ion concentration is the same as the hydrogen ion concentration. A hydronium ion is a hydrated hydrogen ion. Thus: [H+] = 1.0 x 10-4 mol/L pH = 4 (solution is mildly acidic) 4. Because of the prevalence and importance of acids, they have been used and studied for hundred of years. Over time, the definitions of acid and base have been refined 4.1 – Outline the historical development of ideas about acids including those of • Early experiments on acids and bases showed they had a number of properties o Acids have a sour taste; bases (alkalis) have a bitter taste o Acids are corrosive o Acids and bases change the colour of vegetable dyes o Acids and bases neutralise one another on mixing – Lavoisier • Antoine Lavoisier (1743-1794) – The oxygen theory of acids o First chemist to put forward a theory of acidity o Showed that many non-metal compounds containing oxygen (e.g. CO2, SO2) produced acids when dissolved in water o 1776 – suggested the presence of the oxygen in these compounds gave them their acidic properties o Did not explain why oxides of metals were not acidic – Davy • Humphry Davy (1778-1829) – The hydrogen theory of acids o 1772 – Joseph Priestley had generated HCl gas by reacting sea salt with concentrated H2SO4  Readily dissolved in water to form a very acidic solution – ‘marine acid’ o Davy electrolysed samples of the solution & showed it produced H gas and Cl gas when electrolysed  No oxygen formed – named it hydrochloric acid & stated it was a compound of H and Cl only  Other chemists showed that other acids (e.g. HCN) also contained no oxygen but did contain hydrogen 18

o Proposed that the presence of H in acids gave them their acidic properties o Did not explain why many compounds of H were not acidic (e.g. methane) o Von Liebig extended theory – stated acids contained ‘replaceable hydrogen’ – reasoned that when acids attack metals the metals replace the hydrogen in the acid to form a salt  Only compounds with replaceable hydrogen would attack these metals (metal + acid  salt + hydrogen)  Failed to account for production of gases such as NO2 (rather than H) when conc. HNO3 attacked metals – Arrhenius • Svante Arrhenius (1859-1927) – The hydrogen ion theory of acids o During electrolysis of aqueous solutions of all common acids, H gas produced at cathode ( –tive electrode) o Applied his theory of electrolytes to observation and proposed that cause of H gas was conversion of H+ ions in the water into H2 molecules 2H+ (aq) + 2e-  H2 (g) o 1884 stated that all acidic solutions are formed when acids ionise into ions when they dissolve in water  Used the word ‘dissociate’ incorrectly E.g. HCl (g)  H+ (aq) + Cl- (aq) H2SO4 (l)  2H+ (aq) + SO42- (aq) o Recognised some acids (e.g. acetic) were weaker than others – proposed that they did not ionise as completely as strong acids did in water o Important in development of pH scale as it recognised importance of [H+] ions in water solution o Stated that a base is a substance that produces OH- ions when dissolved in water to form an alkaline solution E.g. NaOH (s)  Na+ (aq) + OH- (aq) Ba(OH)2 (s)  Ba2+ (aq) + 2OH- (aq) o Arrhenius acid: a substance that produces H+ ions in water solution o Alkaline solution: a solution containing OH- ions & having a pH > 7 o Arrhenius base: a substance that produces OH- ions in aqueous solution o Proposed that when an acid neutralises a base it is the H+ ions & OH- ions that react to form neutral water E.g. HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l) Net ionic: H+ (aq) + OH- (aq)  H2O (l) o Did not explain why are metallic oxides & carbonates basic  Contain no OH- ions  Many carbonates are insoluble yet they do react with acids and neutralise them o Did not explain why are solutions of various salts acidic or basic rather than neutral  Even though a solution of NaCl is neutral, solution of ZnCl2 is acidic, solution of sodium sulfide is basic 4.2 – Outline the Bronsted-Lowry theory of acids and bases • Acids are proton donors • Bases are proton acceptors • Acid-base behaviour thought of as proton exchange o Substance cannot act as an acid without another behaving as a base • Assigned a role to the solvent – not just an inert liquid in which solutes dissolved, water an ionising solvent o Molecular acids (e.g. HCl gas) dissolve in water to produce ions – according to theory occurs because a proton is donated from the molecular acid to the water molecule to produce the hydronium ion HCl (g) + H2O (l)  H3O+ (aq) + Cl- (aq) proton donor proton acceptor

o HCl is a B-L acid, water is a B-L base and not just a solvent o Presence of hydronium ions gives the solution its acidic properties o Molecular bases (e.g. NH3 gas) dissolve in water to generate ions – here water behaves as a B-L acid NH3 (g) + H2O (l)  NH4+ (aq) + OH- (aq) proton acceptor proton donor

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o Water now seen to have a dual nature – can behave as either a proton donor or acceptor • Allowed chemists to venture outside aqueous chemistry and apply these concepts to non-aqueous solvents or in gas-phase reactions 4.3 – Describe the relationship between an acid and its conjugate base and a base and its conjugate acid • B-L theory develops the concept of acid-base pairs o When a B-L acid donates its proton to the base, the anion of the acid is proton deficient  Can act as a base as it can accept a proton and reform the original acid HF (aq) + H2O (l) H3O+ (aq) + F- (aq)

- HF/F- acid base pair HF = B-L acid F- = conjugate base of HF - H2O/H3O+ acid base pair H2O = B-L base H3O+ = conjugate acid of H2O o In the Bronsted-Lowry theory each acid has its conjugate base and each base has its conjugate acid • Comparative strength of conjugate acid-base pairs o Strong acids have very weak conjugate bases  E.g. hydrochloric acid is a strong B-L acid, its conjugate base (Cl-) is a very weak base – it has a very poor ability to accept protons o Weak acids have relatively strong conjugate bases  E.g. acetic acid is a weak B-L acid, its conjugate base (CH3COO-) is a stronger base than Cl- ions o Strong bases have very weak conjugate acids  E.g. OH- ions are strong B-L bases, their conjugate acid (H2O) is a very weak acid o Weak bases have relatively strong conjugate acids  E.g. HCO3- ions are weak B-L bases, their conjugate acid (H2CO3) is a stronger acid than water • When acids and bases are combined in aqueous solution the position of the equilibrium is determined by the relative strengths of the conjugate acid-base pairs o Strong acids react with strong bases to form weak conjugate bases and weak conjugate acids o Equilibrium lies so far to the right that equation written as going to completion E.g. H3O+ (aq) + OH- (aq)  H2O (l) + H2O (l) o Reactions between a stronger acid & a weaker conjugate base produce an equilibrium that lies to the right E.g. HSO4- (aq) + NH3 (aq) SO42- (aq) + NH4+ (aq)  HSO4- is a stronger acid than NH4+, NH3 is a stronger base than SO42-, thus equilibrium lies to the right o Reactions between a weaker acid and a stronger conjugate base produce an equilibrium that lie to the left E.g. CH3COOH (aq) + F-(aq) CH3COO- (aq) + HF (aq) HF is a stronger acid than CH3COOH, CH3COO- is a stronger base than F-, thus equilibrium lies to left o Reactions between a stronger base & a weaker conjugate acid produce an equilibrium that lies to the right E.g. S2- (aq) + H2CO3 (aq) HS- (aq) + HCO3- (aq) 2o S is a stronger base than HCO3 , H2CO3 is a stronger acid than the HS-, thus equilibrium lies to the right  4.4 – Identify a range of salts which form acidic, basic or neutral solutions and explain their acidic, neutral or basic nature • B-L theory explains why various salt solutions are acidic or basic o Basic salts form basic solutions, acidic salts form acidic solutions, other salts are neutral o The reaction of a salt with water to produce a change in pH is called hydrolysis Weak Acid + Strong Base  Basic Salt + Water o E.g. Na2CO3 – carbonate ion acts as a B-L base, water acts as a B-L acid CO32- (aq) + H2O (l) HCO3- (aq) + OH- (aq)  OH- is a stronger base than the CO32-, HCO3- is a stronger acid than water thus equilibrium lies to left  Sufficient OH- ions formed however to produce a basic solution 20

Weak Base + Strong Acid  Acidic Salt + Water o E.g. Cu(NO3)2 – acidic salt; copper (II) ion in solution surrounded by a primary hydration shell of 4 water molecules  Bonded to the central copper ion by coordinate covalent bonds  Hydrated copper (II) ions act as B-L acids in water Cu(H2O)42+ (aq) + H2O (l) Cu(H2O)3OH+ (aq) + H3O+ o E.g. NH4Cl is the salt formed from NH3 (or NH4OH) and HCl, NH4 acts as a weak B-L acid NH4+ (aq) + H2O (l) NH3 (aq) + H3O+ (aq)  Equilibrium lies to left as H3O+ is a stronger acid than the NH4+ ion & NH3 is a stronger base than water Strong Acid + Strong Base  Neutral Salt + Water o Salts do not react with water to any appreciable extent – pH remains unchanged e.g. NaCl, KBr, NaNO3  Salts of weak bases and weak acids may be neutral if the relative strengths of the acid and base are similar 4.5 – Identify conjugate acid/base pairs

Conjugate acid-base pairs in decreasing order of their strength as acids or bases

4.6 – Identify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutions • Some chemical species can behave as either acids or bases depending on other chemical substances present o Amphiprotic – not same as ‘amphoteric’ – amphiprotism is a classification only within the B-L theory o E.g. HCO3- acts as a proton acceptor when added to a solution of strong acid HCO3- (aq) + H3O+ (aq)  H2CO3 (aq) + H2O (l) H2CO3 (aq) H2O (l) + CO2 (g) The H3O+ ions react with the HCO3- ions to form carbonic acid which decomposes to form H2O & CO2 Presence of the strong acid & loss of CO2 drives the acid-base reaction to the right o HCO3- ions act as proton donors when added to a solution of sodium hydroxide HCO3- (aq) + OH- (aq)  CO32- (aq) + H2O (l)  Presence of the strong base drives the reaction to the right o Other examples are HPO42- and water Proton donor: H2O (l) + HS- (aq) OH- (aq) + H2S (aq) Proton acceptor: H2O (l) + NH4+ (aq) H3O+ (aq) + NH3 (aq)   4.7 – Identify neutralisation as a proton transfer reaction which is exothermic • Neutralisation reactions are proton transfer reactions which are exothermic o Amount of heat liberated per mole when a strong base is neutralised by a strong acid is almost the same no matter what acid or base 21

Same reaction occurs in each case – neutralisation of a hydronium ion & hydroxide ion to form water H3O+ (aq) + OH- (aq)  2H2O (l) o Neutralisation reactions involving weak acids or weak bases produce slightly less heat per mole • Diluting strong acids safely – small quantities of acid added to large volume of water with constant stirring 4.8 – Describe the correct technique for conducting titrations and preparation of standard solutions Volumetric analysis involves the determinations of the volume of a standardised solution that is required to react with the substance being analysed • In acid-base analysis the reaction is a neutralisation reaction • Titration: a technique used in volumetric analysis in which one reactant (dissolved in a solvent) is added to another (also dissolved in a solvent) from a burette until an end point is reached o Standard solution must have an accurately known concentration – may be primary or secondary standard o Primary standard: a substance used in volumetric analysis that is of such high purity & stability that it can be used to prepare a solution of accurately known concentration  Prepared using chemicals which satisfy a special list of criteria: - High level of purity - Accurately known composition - Free of moisture - Stable & unaffected by air during weighing – e.g. doesn’t absorb moisture from air or react - Readily soluble in pure water - High molar weight solid to reduce % error in weighing - Reacts instantaneously & completely Acid standards Potassium hydrogen phthalate (KHC8H4O4) Benzoic acid (C7H5O2) Oxalic acid (H2C2O4.2H2O) Base standards Sodium carbonate (Na2CO3) Sodium hydrogen carbonate (NaHCO3) Borax (Na2B4O7.10H2O) •

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o Secondary standards – solutions whose concentrations have been determined using primary standards  E.g. HCl solution can become a secondary standard by reacting it first with a known primary standard such as NaCO3 solution o Preparation of a standard solution  Calculations – mass required of primary standard  Rinsing procedure – once with tap water, twice with distilled water  Step 1: mass required of primary standard weighed in a clean, dry tray  Step 2: transferring the solid into a clean volumetric flask – use clean funnel & wash bottle  Step 3: dissolving the solid – add small amount of distilled water and swirl  Step 4: adding water to the mark – drop by drop when near mark until the bottom of the meniscus is aligned with the line, avoid parallax error, invert flask several times, label flask o Titration techniques and calculations – standard solution used to analyse unknown solution  Neutralisation chemical equation  Step 1: preparing burette - Rinse burette then rinse with unknown solution – space below tap must be rinsed - Clamped to retort stand, small funnel (rinsed with water & unknown solution) used to fill burette with unknown solution (ensure solution has filled space below tap) - Position of base of meniscus recorded - Liquid in burette is called the titrant  Step 2: preparing pipette and transferring standard solution into conical flask - Pipette rinsed with water & standard solution - Filled with standard solution till bottom of meniscus is on engraved line – aliquot formed - Aliquot: a known volume of a liquid 22

- Conical flask rinsed – aliquot of standard solution transferred to flask with tip of pipette resting against the inside of the glass wall  Step 3: performing the titration - Suitable indicator chosen for neutralisation reaction – 3 drops added to conical flask - Solution from burette added to flask, flask swirled, wash bottle use to rinse sides - Continue to add until colour of indicator just changes – record volume change in burette - First titration called ‘rough’ titration – approximate end-point located so more accurate result achieved in subsequent titrations - Titre: the volume of solution delivered from the burette that achieves an end-point - End-point: the stage in a titration at which the indicator just changes colour  Step 4: calculating the concentration of the unknown acid - Stoichiometric ratio, number of moles of primary standard used, concentration of unknown solution • Choosing a suitable indicator – depends on the strength of the acid and base involved o Chosen so that pH at end-point matches as closely as possible pH at the equivalence point of the titration  Because indicators change colour over a narrow pH range rather than an exact pH this is important o Equivalence point: the point where the acid has stoichiometrically reacted with the base  Indicators are weak acids HIn (aq) + H2O (l) In- (aq) + H3O+ (aq)  Equilibrium influenced by pH of solution  Colour changes that we observe occur because the un-ionised molecule has a different colour from that of the conjugate base Indicator Methyl orange Bromothymol blue Litmus Phenolphthalein pH range 3.1 ↔ 4.4 6.0 ↔ 7.6 5.0 ↔ 8.0 8.3 ↔ 10.0 Colour of un-ionised molecule (HIn) red yellow red colourless Colour of ionised molecule (In-) yellow blue blue red

o Matching pH graphs with indicator pH ranges  Strong acid – strong base titration – salt that is formed does not hydrolyse to any appreciable extent - pH at equivalence point = 7 – centre of inflection - Each indicator may be used but bromothymol blue best matches pH at equivalence point  Strong acid – weak base titration – pH at equivalence point 7 due to hydrolysis of the basic salt produced in the neutralisation reaction - Phenolphthalein suitable indicator - Buffering effect on pH as base is added – well before end point solution contains both weak acid & its conjugate base – pH rise only slowly in this region 23



Weak acid – weak base titration – pH at equivalence point = 7 - No rapid change in pH even at equivalence point - No single indicator will achieve a sharp end point – best to avoid these titrations

4.9 – Qualitatively describe the effect of buffers with reference to a specific example in a natural system • Buffers – solutions that resist changes in pH when small quantities of an acid or base are added to them o Usually contain a weak B-L acid & its conjugate base or a weak B-L base & its conjugate acid o By choosing the correct amounts of the weak acid & weak base in the solution pH of solution can be fixed within narrow limits  E.g. weak acid HA & its conjugate base A-; if HA added to water HA (aq) + H2O (l) A- (aq) + H3O+ (aq) (1) + - H3O ion stronger acid than HA & A stronger base than water thus the equilibrium lies well to the left  If A- added to water (e.g. as a sodium salt) a A- (aq) + H2O (l) HA (aq) + OH- (aq) (2) - OH stronger base than A & HA stronger acid than H2O thus this equilibrium also lies well to the left  By preparing a solution that contains both HA and A- in approximately equal amounts create a buffer - System reaches equilibrium quickly - If add acid (e.g. HCl) as a source of hydronium ions equilibrium (1) shifts to the left - Most of excess acid converted into an unionised molecular acid (HA) resulting in very little pH change - If add base (e.g. NaOH) as a source of hydroxide ions equilibrium (2) shifts to the left - Very little change in pH as the added strong base is converted into a much weaker base (A-) • Natural Buffers – our body fluids & secretions must be maintained in a narrow pH range in order for biochemical processes to occur at an optimal rate – buffers keep our bodies in pH balance o E.g. maintenance of pH of blood in the range 7.35-7.45  Consists of a carbonic acid – hydrogen carbonate buffer linked to the haemoglobin – oxyhaemoglobin equilibrium H2CO3 (aq) + H2O (l) ↔ HCO3- (aq) + H3O+ (aq) (1)  Haemoglobin – bluish-red iron-protein molecule that absorbs & transports oxygen in our blood - Complex molecule that contains 4 haem (Hb) groups bound to an iron (II) ion - Each haemoglobin molecule can bind to 4 oxygen molecules to form oxyhaemoglobin - Haemoglobin is a weak acid – weak proton donor HHb4 (aq) + H2O (l) + 4O2 (aq) Hb4O8- (aq) + H3O+ (aq) (2) (2) o Inhalation – O2 drawn into lungs & diffuses into the bloodstream  Increase in O2 in blood causes equilibrium (2) to shift to the right  Acidity increases as H3O+ formed, blood contains high levels of oxyhaemoglobin as it leaves the lungs  Buffer system prevents oxygenated blood from becoming too acidic – equilibrium (1) shifted to the left in order to remove excess H3O+ 24
-

Carbonic acid that forms breaks down to release CO2 When oxygenated blood reaches oxygen-requiring cells in the body tissues oxygen diffuses into them Equilibrium (2) shifts to the left & hydronium ions are removed from the blood Cells release CO2 (a product of cellular respiration) back into plasma – leads to formation of H2CO3 CO2 (aq) + H2O (l) H2CO3 (aq) (3)  Increase in carbonic acid concentration shift equilibrium (1) to right & blood acidity increases o Exhalation – deoxygenated blood returns to lungs & CO2 diffuses from blood  Partial pressure of CO2 gas rises inside the lungs  As we exhale CO2 (aq) CO2 (g) shifted to right – equilibria (3) & (1) shift to left and pH rises

   

4.0.1 – Gather and process information from secondary sources to trace developments in understanding and describing acid/base reactions Understanding limited to known properties of acids & bases 1776: Antoine Lavoiser – the oxygen theory of acids 1810: Humphry Davy – the hydrogen theory of acids 1884: Svante Arrhenius – the hydrogen ion theory of acids 1923: The Bronsted-Lowry theory of acids o Bronsted & Lowry independently proposed this theory • Lewis theory of acids – acid base reactions treated in terms of electron pairs instead of specific substances • From initial theories new theories created & refined in order to explain and incorporate certain phenomenon that were outside these definitions of acid/base reactions • Using these models, chemists are able to hypothesise and make predictions about acid/base reactions 4.0.2 – Choose equipment and perform a first-hand investigation to identify the pH of a range of salt solutions AIM: To identify the pH of a range of salt solutions METHOD: 1. Examples of salt solutions to test: ammonium chloride, potassium carbonate, sodium chloride, hydrogen sulfate etc. 2. Spot plate used to test colour of universal indicator in each salt 3. pH meter calibrated & used to measure pH DISCUSSION:  pH meter more precise than indicator solution or universal indicator. 25 • • • • •

 

Experiment repeated & average taken Validity – equal concentrations of salt solutions as well as same amount, pH of distilled water measured as a control, pH meter recalibrated after regularly

4.0.3 – Perform a first-hand investigation and solve problems using titrations and including the preparation of standard solutions, and use available evidence to quantitatively and qualitatively describe the reaction between selected acids and bases AIM: To standardise a solution of sodium hydroxide using a primary standard METHOD 1. Rinsing procedure: rinse with tap water, rinse with distilled water twice, dry with paper towel if necessary 2. Required mass of oxalic acid needed to create a 250mL 0.0500 mol/L solution was calculated & weighed 3. Oxalic acid dihydrate crystals (H2C2O4.2H2O) added to volumetric flask using the funnel and wash bottle, 4. Small amount of water was added to the flask, flask swirled until all of the crystals dissolve 5. Water added until the base of the meniscus was on the line, flask stoppered and mixed thoroughly 6. Burette & pipette rinsed (also rinsed with their solutions) and set up 7. Filled (pipette with oxalic acid, burette with NaOH), position of base of meniscus in burette recorded 8. Acid added to the conical flask with the tip of the pipette against the glass wall for 10 seconds 9. 3 drops of phenolphthalein indicator were added to conical flask, rough titration conducted – base from burette added whilst swirling the contents of the flask, wash bottle used to rinse the internal walls of flask 10. When a permanent colour change appeared the volume of base added recorded 11. Conical flask rinsed & titration conducted further three times. DISCUSSION:  Validity – selection of primary standard & indicator, not conducted at standard laboratory conditions so some of measurements taken not precise (e.g. with pipette)  Calculations: nAcid : nBase = 2:1, outlier  Reliability – repeated & averaged, small range VARIABLES: concentration of NaOH (independent), amount of base added to flask (dependent), concentration of standard solution, amount of indicator added, environment (controlled) 4.0.4 – Perform a first-hand investigation to determine the concentration of a domestic acidic substance using computer-based technologies AIM: To determine the concentration of a domestic substance (dilute white vinegar) using phenolphthalein and a pH probe by neutralisation with a standardised sodium hydroxide solution. METHOD: 1. pH probe calibrated, 25.0 mL of acid added to clean beaker with pipette 2. 100 mL of distilled water added to increase volume, 3 drops of phenolphthalein added 3. Placed on magnetic stirrer plate & stirrer inserted, pH probe placed in beaker & clamped 4. Titrated with sodium hydroxide solution 5. Continuous pH measurement – tap of burette opened and allowed to run, reduced as approached endpoint, repeated 6. Single-point pH – 5.0 mL NaOH solution added in intervals & pH measured with each, drops added when approached end point DISCUSSION:  Inflection point of graph is equivalence point, calculations – no. moles, concentration of diluted vinegar, concentration of undiluted vinegar  Validity – correct calibration, inflection occurred within end-point, more precise than indicator alone  pH at equivalence point not equal to 7 – CH3COO- weak base 26

4.0.5 – Analyse information from secondary sources to assess the use of neutralisation reactions as a safety measure or to minimise damage in accidents or chemical spills Neutralisation reactions most effective measure to minimise damage in chemical spills o Advantages – cost effective, quick and easy, amphiprotic substances can neutralise acids and bases  Resulting products of neutralisation are non-toxic and essentially harmless o Disadvantages – exothermic reaction – can be dangerous or further damaging, can’t be carried out on skin • In the laboratory – neutralisation reactions can be used to clean up spilled acids or bases o Because neutralisation is exothermic never use concentrated acids or bases  Potential to cause boiling & evolution of noxious fumes, strong base or acid will cause further damage o If spilled on person’s clothes or skin immediate flushing with copious amounts of water o Spill isolated with sand or vermiculite to prevent acid flowing to other areas, acid-soaked sand or vermiculite then cleaned up & neutralised in a safe location o NaHCO3 also commonly used – amphiprotic so can neutralise both acidic & basic spills  Weak base & non-toxic so can be used in excess & handled safely  Solid at room temperature so can control the spread of chemicals by soaking up solution, cheap Neutralising a base: HCO3- (aq) + OH- (aq)  H2O (l) + CO32- (aq) Neutralising an acid: HCO3- (aq) + H3O+ (aq)  H2O (l) + H2CO3 (aq) • Larger-scale neutralisations – if very large spills occur best procedure is initially to prevent acid escaping into the drains or into soil – large amounts of inert sand or vermiculite used o Once absorbed sand or vermiculite placed in chemical waste container, removed for neutralisation off site o Na2CO3 powder followed by copious amounts of water used to neutralise & dilute any remaining acid o If alkaline material has been spilled can use NaHCO3 as it is amphoteric NaHCO3 (s) + NaOH (aq)  Na2CO3 (aq) + H2O (l) o Sodium hydrogen phosphate solid also used in base spills – amphoteric o In many factory situations alkaline wastes can be carefully neutralised with dilute HCl followed by large volumes of water to prevent high concentrations of ions, or solutions with too high or too low a pH being discharged in wastewater effluents 5. Esterification is a naturally occurring process which can be performed in the laboratory 5.1 – Describe the differences between the alkanol and alkanoic acid functional groups in carbon compounds • Alkanols (CnH2n + 1OH) are compounds that contain the alcohol functional group ( –OH) o Polar molecules due to presence of alcohol functional group o M.p. & b.p. are higher than corresponding alkanes because of dipole-dipole attractions o Methanol & ethanol have significantly high m.p. & b.p. for such small molecules – H-bonding between δ+ H atoms & δ- O atoms of alcohol groups on neighbouring molecules o H-bonding also explains high water solubility of first 4 members of alkanol homologous series o As chain length increases, however, they become increasingly insoluble o Essentially neutral molecules – alcohol functional group strongly bonded to C chain & no tendency for loss of hydroxide ions when alkanols dissolve in water •

27

•

Alkanoic acids (CnH2n+1COOH)contain the carboxylic acid functional group ( –COOH) o Weak organic acids – can be neutralised by strong bases to form salts & water o Naming – identify no. of C present in straight chain, select correct stem to name parent alkane, remove ‘e’ & replace with suffix ‘oic acid’ – COOH functional group always at end of chain so this C is locant 1  Methanoic acid = formic acid, ethanoic acid = acetic acid

•

Comparing the functional groups of alkanols & alkanoic acids o Both polar molecules but carboxylic functional group in alkanoic acids is more polar than hydroxide group in alkanols due to C-O, C=O and O-H bonds o Alkanols & alkanoic acids have higher m.p. & b.p. than their parent alkanes

5.2 – Identify the IUPAC nomenclature for describing the esters produced by reactions of straight-chained alkanoic acids from C1 to C8 and straight-chained primary alkanols from C1 to C8 • Nomenclature of esters – alkyl alkanoate esters named after their parent alkanols & alkanoic acids 1. Count number of C atoms in alkyl group derived from original alcohol – replace ‘anol’ suffix with ‘yl’ 2. Identify no. of C atoms in alkanoate chain derived from alkanoic acid – replace ‘oic acid’ suffix with ‘oate’ 3. Ester is 2 separate words – first name comes from alkanol & second from alkanoic acid, e.g.
Alkanol Methanol Ethanol Propanol Butanol Pentanol Alkanoic acid Ethanoic acid Propanoic acid Butanoic acid Pentanoic acid Hexanoic acid Ester Methyl ethanoate Ethyl propanoate Propyl butanoate Butyl pentanoate Pentyl hexanoate

propyl pentanoate (or 1-propyl pentanoate)

5.3 – Explain the difference in melting point and boiling point caused by straight-chained alkanoic acid and straight-chained primary alkanol structures • Alkanoic acids have higher m.p. & b.p. than their corresponding alkanols o Have greater molar weights than equivalent alkanols thus dispersion forces between molecules greater o Usually slightly more polar than alkanols due to C-O, C=O & O-H bonds thus dipole-dipole forces greater

28

• •

Same molecular weights so dispersion forces between molecules very similar Alkanoic acid & alkanol have significantly higher b.p. than that of ester o All 3 molecules are polar, alkanol & alkanoic acid, however, exhibit H-bonding between molecules & this leads to an elevation in b.p. o Ester cannot form H-bonds as there are no OH functional groups in molecule thus b.p. much lower • When alkanol & alkanoic acid of same molar weight compared alkanoic acid has higher m.p. or b.p. o Mainly due to more extensive H-bonding between alkanoic acid molecules  2 oxygen atoms which can form hydrogen bonds with other hydrogen atoms of another molecule  More energy needed to break forces between molecules

o As carbon chain lengthens increasing dispersion forces dominate & there is a general trend of increasing boiling point with increasing molar weight 5.4 – Identify esterification as the reaction between an acid and an alkanol and describe, using equations, examples of esterification • Esters – produced by acid-catalysed reaction between alcohol & carboxylic acid, endothermic reaction o If alcohol is an alkanol & carboxylic acid is an alkanoic acid then ester formed is an alkyl alkanoate Alkanol + Alkanoic Acid Alkyl Alkanoate + Water o Condensation reaction – a reaction in which 2 molecules combine together with the elimination of a smaller molecule, in this case water  Comes from OH group of alkanoic acid & H of the alkanol functional group o All esters contain the ester linkage or ester functional group ( –COO– )



E.g. methanol + ethanoic acid (in presence of a catalyst)  methyl ethanoate ester + water H2SO4 

+

+

H2O

5.5 – Describe the purpose of using acid in esterification for catalysis • Esterification – the process of making an ester o Usually quite slow & does not proceed to completion o Rate of reaction can be increased by adding a suitable catalyst (e.g. 1-3 mL of conc. H2SO4) & heating the reaction mixture to increase the KE of the molecules  Catalyst decrease the time to reach equilibrium – provides an alternate reaction pathway with lower EA

29

5.6 – Explain the need for refluxing during esterification • Esterification reaction mixture heated to increase KE of molecules – endothermic reaction • Reactants & products of esterification are volatile & readily vaporise on heating – to avoid loss of material from reaction flask mixture is heated using a reflux apparatus o Refluxing: the process of heating a mixture of liquids in a flask with an attached condenser in order to prevent the loss of volatile reactants or products o Vapours condense back to liquid state & drip back into reaction vessel o Because system is open to atmosphere there is no build up of pressure due to the production of vapours o Organic liquids & vapours are flammable & care must be taken to avoid fires and explosions – naked flames from Bunsen Burner avoided o Small boiling chips (normally pieces of crushed ceramic) used to prevent ‘bumping’ as they provide a large surface area on which vaporisation can occur without the risks of sudden superheating & the explosive ejection of vapours o Refluxing may occur for hours or days until system reaches equilibrium o % yields obtained vary as ester is lost during the separation & purification process 5.7 – Outline some examples of the occurrence, production and uses of esters • Occurrence – esters are sweet-smelling compounds which occur naturally in all natural systems o Can be formed not only from alkanoic acids but other organic & inorganic acids o Energy transfer reactions in cells – phosphate esters formed from phosphoric acid (e.g. ATP) o Fats & oils – triesters of glycerol & various long-chain carboxylic acids – known as triglycerides o Natural waxes – consist of mixed esters of long-chain alkanols & long-chain alkanoic acids (fatty acids) Production – large scale version of laboratory method of production used to make esters o Natural and synthetic esters mixed to produce cosmetics and flavours in foods  Relatively inexpensive production cost compared to natural flavours o Certain esters are monomers for the production of polyesters by condensation polymerisation Uses: Flavours in foods e.g. ethyl butanoate is pineapple flavour, methyl butanoate is apple flavour Used as a fragrance in cosmetics e.g. benzyl ethanoate used for jasmine fragrance. Solvent for paint/lacquers e.g. ethyl ethanoate is used in nail-polish remover. Medications e.g. methyl salicylate used for muscular heat treatment. Lubricants in jet engines and refrigeration systems Plasticisers – added to hard plastics (e.g. PVC) to soften them & increase their flexibility Biofuels – esters such as methyl soyate formed in production of biodiesel – can be used in diesel engines Emulsifying agents (produce stable homogeneous dispersions of one insoluble liquid in another) in cosmetics & foods

•

• o o o o o o o o

30

5.0.1 – Identify data, plan, select equipment and perform a first-hand investigation to prepare an ester using reflux AIM: To prepare an ester using the reflux technique RISK ASSESSMENT:  Glacial acetic acid (CH3COOH) – solutions >13M are corrosive to skin and eyes, toxic if ingested, highly irritant vapour  1-butanol (C4H9OH) - slightly toxic by all routes of exposure; flammable; irritant vapour  Sulfuric acid (H2SO4) >10M – highly toxic by all routes of exposure, highly corrosive to skin or eyes, considerable heat evolved when mixing with water  1-butyl ethanoate (CH3COOC4H9) – breathing vapours may cause drowsiness & dizziness, irritant, flammable METHOD: 1. 24 mL of glacial acetic acid and 18 mL of 1-butanol was measured & poured into dry round-bottom flask 2. 10 drops of concentrated sulfuric acid added to flask in fume hood, 3 boiling chips added to the flask 3. Flask and condenser for reflux were set up, tap turned on so that cold water flowed from the tap to the base of the condensing jacket then out the top and back into the sink 4. Electric heating mantle used to heat flask till the mixture was boiling 5. Heat reduced until condensate dripped back into reaction flask at about 2 drops/second so that no vapours were lost, mixture refluxed for 20 minutes, heating mantle turned off and apparatus was allowed to cool 6. Mixture added to 30 mL of saturated salt water in a beaker, poured into a separating funnel 7. Upper layer of crude ester isolated by running out bottom layer of mixture into waste beaker 8. Saturated NaHCO3 added to separating funnel to remove any acid in small amounts until mixture stopped fizzing 9. Ester isolated again using the separating funnel & run out into a beaker containing 20 mL of water 10. The odour of the ester was compared to those of the original alkanol and alkanoic acid RESULTS:
Chemical Acetic acid 1-butanol 1-butyl ethanoate Odour Vinegar-like, pungent, sour Slightly vinous, sweetish Sweet, fruity

DISCUSSION:     

conc. H2SO4 CH3COOH (l) + C4H9OH (l) CH3COOC4H9 (l) + H2O (l) Validity – catalyst, heating, refluxing, using excess acid – decreased time to reach equilibrium, increased yield Separation – NaCl reduces solubility of ester, NaHCO3 neutralises excess acid, Improvements – longer refluxing, ester removed & dried after separation then fractional distillation Accuracy questionable – reaction did not proceed to completion, still impurities after separation Reliability not gauged – experiment not repeated; could be repeated and obtaining same results

5.0.2 – Process information from secondary sources to identify and describe the uses of esters as flavours and perfumes in processed foods and cosmetics • Far cheaper to manufacture on a large scale than to use natural flavours • Ethyl butanoate – pineapple-like smell o Used in the manufacture of artificial rum as it adds a unique pineapple aroma to the rum • Pentyl butanoate – apricot/pear-like odour, strong, sweet smelling fragrance o Widely used in the cosmetics industry as a major ingredient in many perfumes • Ethyl ethanoate o Used in the cosmetics industry as a nail polish remover o Used in perfumes where it evaporates quickly leaving a scent of the perfume on the skin o Used in the food industry in confectionary and some fruits 31

Chemistry Notes 2010
Core Module 3: Chemical Monitoring and Management
Contextual Outline The state of our environment is an important issue for society. Pollution of air, land and water in urban, rural and wilderness areas is a phenomenon that affects the health and survival of all organisms, including humans. An understanding of the chemical processes involved in interactions in the full range of global environments, including atmosphere and hydrosphere, is indispensable to an understanding of how environments behave and change. It is also vital in understanding how technologies, which in part are the result of chemical research, have affected environments. This module encourages discussion of how chemists can assist in reversing or minimising the environmental problems caused by technology and the human demand for products and services. Some modern technologies can facilitate the gathering of information about the occurrence of chemicals – both those occurring in natural environments and those that are released as a result of human technological activity. Such technologies include systems that have been developed to quantify and compare amounts of substances. This module increases students’ understanding of the nature, practice, applications and uses of chemistry and the implications of chemistry for society and the environment. • Monitoring combustion reactions o Complete combustion of octane 2C8H18 (l) + 25O2 (g)  16CO2 (g) + 18H2O (l) o Incomplete combustion of octane examples 2C8H18 (l) + 17O2 (g)  16CO (g) + 18H2O (l) 2C8H18 (l) + 13O2 (g)  8CO (g) + 8C (s) 18H2O (l)
(low concentrations of O2)

•

Formation of NOx

N2 (g) + O2 (g)  2NO (g) 2NO (g) + O2 (g)  2NO2 (g)
(high temperatures in engines produced by near complete combustion)

•

Haber Process N2 (g) + 3H2 (g) ↔ 2NH3 (g)
Fe3O4

ΔH = –92 kJ/mol
(400-500°C, 25-35 MPa)

•

Ozone Chemistry o Formation O2 (g) + hv  2O. (g) O. (g) + O2 (g)  O3 (g) o Natural Decomposition – many processes O3 (g) + hv  O. (g) + O2 (g) . O (g) + O . (g)  O2 (g) o Decomposition by CFCs CF3Cl (g) + hv  CF3. (g) + Cl. (g) Cl. (g) + O3 (g)  ClO. (g) + O2 (g) ClO. (g) + O. (g)  O2 (g) + Cl. (g) ΔH-ve ΔH -ve

1

1. Much of the work of chemists involves monitoring the reactants and products of reactions and managing reaction conditions 1.1 Outline the role of a chemist employed in a named industry or enterprise, identifying the branch of chemistry undertaken by the chemist and explaining a chemical principle that the chemist uses • • • Dr David Proctor, Primary Researcher CSIRO Thermal & Fluids Engineering Branch of chemistry: Analytical chemistry, developmental chemist Chemical principle: the pulsing effect produced in the combustion process can be used to recirculate unburnt gases in a combustion chamber o New technologies addressing the issue of incomplete combustion will significantly reduce energy costs, boost productivity and help to clean up greenhouse gas emissions. o NGPC (new generation pulse combustion) – developing technology  Combines the pulsing effect of sound waves produced in the combustion process with a new geometry of burner and combustion chamber  Sound waves are driven into resonance by geometry of burner, locking combustion instability into a very stable repetitive pattern  Burner becomes self aspirating – no need for a fan to supply air for combustion  New air drawn into the combustion chamber, unburnt exhaust gases recirculated  Flame consists of a series of discrete ‘flamelets’ – ignited by the hot products of the previous flamelets producing much hotter and cleaner flames  Equipment self-cleaning as a result of the micro-vibrations generated  By operating each pulse combustor out of phase with its partner sound levels are much lower – destructive interference  Able to show that zero levels of unburnt hydrocarbons and carbon monoxide had been achieved  Able to reduce fuel consumption by 50%, greenhouse gas emissions by 30%, NOx emissions to levels far below the strictest environmental standards, thermal efficiencies of 97%.  Proctor & team able to overcome the problem of turning down the burning rate without shutting down the flame – now suitable for industrial use  Proctor’s team first to find a way to scale up laboratory pulse combustor into commercial sized processing facilities  Contracts to develop NGPC for range of applications e.g. power generation, petrochemical liquid heating, food ovens, in steel & smelting industry, commercial water heating

1.0.1 Gather, process and present information from secondary sources about the work of practising scientists identifying: - the variety of chemical occupations The Royal Australian Chemical Institute (RACI) has twelve divisions of chemical occupations for membership within Australia: • Analytical Chemistry • Industrial Chemistry • Biomolecular Chemistry • Inorganic Chemistry • Cereal Chemistry • Materials • Chemical Education • Organic Chemistry • Colloid and Surface Science Division • Physical Chemistry • Electrochemistry • Polymer - a specific chemical occupation for a more detailed study • Polymer chemist: investigates the properties of large polymeric molecules o Manipulates their structure, affecting their properties, to produce new & useful plastic products & materials. 2

1.2 Identify the need for collaboration between chemists as they collect and analyse data • Teamwork, collaboration & communication skills are important for chemists o A company may employ many different types of chemists with different skills o Many chemical problems require expertise & in depth knowledge from range of chemical branches o Collaboration between chemists is essential for solving chemical issues, or when dealing with large amounts of data being collected - provide expertise from own particular field for common goal o An industrial process would require collaboration between physical chemists (for equilibrium considerations), organic chemists (for how the reaction occurs) & analytical chemists (for monitoring products) 1.3 Describe an example of a chemical reaction such as combustion, where reactants form different products under different conditions and thus would need monitoring • Monitoring combustion reactions o Chemists & technicians must monitor emissions from combustion reactions to ensure minimum amounts of toxic materials are released into the environment o Toxic emissions result from incomplete combustion reactions & other reactions that occur in association with the main combustion process  I.e. SO2 from combustion of sulfur impurities in fossil fuels, NOx from high temp combustion engines o Complete combustion: 2C8H18 (l) + 25O2 (g)  16CO2 (g) + 18H2O (l) o In most engines normally insufficient oxygen to achieve complete combustion o Incomplete combustion: normally occurs with release of CO &/or soot, unburnt fuel vapours emitted in exhaust gases, e.g.: 2C8H18 (l) + 17O2 (g)  16CO (g) + 18H2O (l) 2C8H18 (l) + 13O2 (g)  8CO (g) + 8C (s) 18H2O (l) o At high temperatures in the engine, N & O combine to form nitrogen oxides (NOx), e.g.: N2 (g) + O2 (g)  2NO (g) 2NO (g) + O2 (g)  2NO2 (g) o CO concentration is low when engine operates at high speed with maximum O intake, high when operating at low speed o Reverse for NOx concentration – high temperatures in engine produced by near complete combustion promote formation of NOx whereas when engine operating at low speed the concentration of NOx is low • Hence under different conditions chemical reactions can proceed in different ways o Reaction conditions monitored to ensure that maximum amount of desired product is being formed o Must limit pollution – CO poisonous gas, C is carcinogenic & can be irritating to the lungs, SO 2 & NO x contribute to acid rain o Must produce maximum amount of energy • Gas-liquid chromatography (GLC): separation & analysis of mixtures of gases based on their differential absorption while passing through tubes (in a carrier gas stream) containing granulated particles with a thin liquid film • Catalytic converters – reduce CO, NOx & unburnt hydrocarbon emissions from vehicle exhausts o Made from alloys of rhodium & platinum that operate at temperatures as low as 150°C o Speed up reactions that convert pollutant gases to materials which are present in air naturally o For catalyst system to work, oxygen sensors must be fitted to monitor exhaust gases  Fuel injection system controlled so correct amounts of fuel & oxygen are mixed prior to combustion  Ensures sufficient oxygen for oxidation of both CO & unburnt hydrocarbons on the catalyst & provides conditions for the reduction of NOx to nitrogen  Reactions involved: − Step 1: Rhodium catalyst in first chamber 2NO (g) + 2CO (g)  N2 (g) + 2CO2 (g) − Step 2: Platinum catalyst in second chamber (additional air added to exhaust gases in this step) 2CO (g) + O2 (g)  2CO2 (g) 2C8H18 (g) + 25O2 (g)  16CO2 (g) + 18H2O (g) 3

o In most modern engines a lean fuel mixture is used – more than required stoichiometric amount of air (oxygen) to burn the fuel  Typically 1:18 ratio of fuel to air (by mass) compared with 1:14 ratio in a normal engine  Less CO & unburnt fuel vapours produced & additional NOx formed due to the higher combustion temperatures are removed by the catalytic converters  Work efficiently only when fuel-air mixture is computer controlled – electronic fuel injection • Monitoring combustion in various industries – industrial chemist in steel industry or coal-power station would monitor & manage level of combustion products released to environment o Exhausts would be continuously monitored to determine concentrations of CO, unburnt hydrocarbons, nitrogen oxides & SO2 using gas probes & gas chromatographs linked to computer control systems o Levels of particulate matter (e.g. soot & aerosols of metal oxides) as well as H2SO4 aerosols would be monitored by collecting & analysing samples removed from the waste gas stream o Smaller industries where flue gases must be monitored by EPA regulations contract mobile laboratories (trucks outfitted with modern instrumentation) to monitor emission levels on-site o Mass emission rates of pollutant gases measured under a variety of load conditions & adjustments have to be made to combustion chamber whenever levels of emissions fail to meet EPA standards o Coal usually contains variable amounts of sulfur compounds, SO2 & SO3 form on combustion  SO3 dissolves in water vapour to form H2SO4 aerosols – concentration must be monitored 2. Chemical processes in industry require monitoring and management to maximise production • Le Chatelier’s Principle: If a system is in equilibrium and a change is made that upsets the equilibrium, then the system alters in such a way as to counteract the change and a new equilibrium is established o Temperature  If forward reaction endothermic, increasing temperature favours formation of product – yield increased  If forward reaction exothermic, reducing temperature favours formation of product – yield increased o Pressure – reactions with one or more gaseous reactants or products  Increasing system pressure (reducing volume) causes equilibrium to shift to side of equation with least number of gaseous molecules  Decreasing system pressure (increasing volume) causes equilibrium to shift to side of equation with most number of gaseous molecules  When equal numbers of molecules on reactant & product sides, pressure has no effect on equilibrium o Concentration  Increasing concentration of reactant shifts equilibrium right to reduce concentration of added reactant  Decreasing concentration of reactant shifts equilibrium left to make more of that reactant o Catalysts – decrease time needed to reach equilibrium, don’t affect position of equilibrium  Provide alternate reaction pathway of lower activation energy

2.1 Identify and describe the industrial uses of ammonia • Ammonia is a feedstock for a large variety of industrial chemicals – increasing demand o Feedstock: materials required for an industrial process; may be natural raw materials or processed raw materials o NH3 is a colourless, choking gas which produces an alkaline solution when dissolved in water o Used as a refrigerant gas 4

Industrial Product Derived from Ammonia Urea, ammonium sulfate, ammonium nitrate, ammonium hydrogen phosphate Nitric acid

Acrylonitrile Diaminoalkanes Cyanides Hydrazine Sulfonamides Aniline derivatives Alkylammonium hydrocarbons Ammonia solution (ammonia hydroxide)

Use of Product • Fertilisers – account for 80% of worldwide use of NH3 o Nitrogen essential for plant growth Production of explosives (e.g. TNT, nitrocellulose & nitroglycerine) Production of nitrate salts Strong laboratory acid Acrylic plastics Nylon plastics Extraction of gold from gold veins Rocket propellant Antibiotic drugs Dyes Cationic detergents Cleaning agents to dissolve & remove grease & dirt from floors and windows Weak base – used to safely neutralise acid spills

2.2 Identify that ammonia can be synthesised from its component gases, nitrogen and hydrogen Industrial manufacture of ammonia Must balance rate of production and maximisation of yield Yield: quantity of product formed in a process; usually expressed as % of expected amount of product Manufactured by process first developed by Fritz Haber in early 1900s, iron catalyst (Fe3O4)
Fe3O4



N2 (g) + 3H2 (g) 2NH3 (g) Nitrogen feedstocks – filtered air (~78% N2) either fractionally distilled (expensive) or extracted using chemical reactions involving natural gas or methane (more common, also yields the H2 required) Hydrogen feedstocks – source varies from one plant to another  Obtained by electrolysis of salt water: 2H2O (l) + 2NaCl (aq)  H2 (g) + Cl2 (g) + 2NaOH (aq)  Derived from steam reforming of natural gas, mainly composed of methane (more common) − Natural gas purified to remove sulfur compounds that will ‘poison’ the Haber catalyst − H2 extracted from pressurised gas by reacting natural gas with steam over Ni catalyst at 750°C CH4 (g) + H2O (g) CO (g) + 3H2 (g) ΔH+ − Air introduced to ensure almost all of remaining methane combusts to form CO2 or reacts with steam to produce CO & H2, ΔH– − CO removed by passing cooled gas mixture over 2 catalysts, combines with steam & forms CO2 & H2 − Acidic CO2 removed by neutralisation with hot solution of K2CO3 − Gaseous mixture methanated – all traces of CO & CO2 removed − Final gaseous mixture contains N2 & H2 in approximate mole ratio of 1:3 as required for reaction − Very small amounts of methane and argon present as impurities − Mixture compressed to 20-25 MPa before sent to Haber chamber

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2.3 Describe that synthesis of ammonia occurs as a reversible reaction that will reach equilibrium Reversible reaction – can proceed in both directions Equilibrium – reversible reaction proceeding in a closed system with rate of forward reaction = rate of reverse reaction At standard temp & pressure reaches equilibrium very slowly, lies well to the left – very little NH3 in mixture High activation energy – a lot of energy is required to break the covalent bonds of H2 & N2 2.4 Identify the reaction of hydrogen with nitrogen as exothermic N2 (g) + 3H2 (g) 2NH3 (g) ΔH = –92 kJ/mol • Exothermic equilibrium: –92 kJ/mol of N2, –46 kJ/mol of NH3 produced o Energy required to break bonds of H2 & N2 is less than the energy given out when bonds of NH3 form 2.5 Explain why the rate of reaction is increased by higher temperatures • Kinetic theory of gases predicts rate of reaction increases when more successful collisions occur in the shortest time • High temperatures = increase in reaction rate o More molecules have sufficient energy to overcome activation energy barrier o Molecules have higher KE therefore more collisions – increased frequency of successful collisions o Increase in temperature increases rate of reaction in both directions – time to reach equilibrium decreases
Fe3O4

2.6 Explain why the yield of product in the Haber process is reduced at higher temperatures using Le Chatelier’s principle • Le Chatelier’s Principle: If a system is in equilibrium and a change is made that upsets the equilibrium, then the system alters in such a way as to counteract the change and a new equilibrium is established • Equilibrium is exothermic – increase in temperature causes equilibrium to shift left to use up added heat 6

o High temperature reduces yield of ammonia, low temperatures favour higher yields 2.7 Explain why the Haber process is based on a delicate balancing act involving reaction energy, reaction rate and equilibrium • Compromise Conditions – conditions are a compromise between equilibrium, kinetic & economic factors o Temperature: 400-500°C – equilibrium & kinetic factors conflict  Compromise temperature, together with a catalyst, allows economical rate of reaction & achieves reasonable yield per cycle through reaction vessel o Pressure: 25-35 MPa  Although equilibrium & kinetic considerations suggest pressure should be as high as possible economic & safety considerations require lower pressure o Yield of ammonia:  ~15-20% per cycle, after 5-6 cycles about 98% of reactants  NH3  Constant removal of ammonia by liquefaction (under pressure) shifts equilibrium right o Reactants: N2 & H2 although equilibrium can be shifted by increasing reactant concentration it is important that 1:3 mole stoichiometry is maintained o Catalyst: provides alternate pathway of lower EA, decreases time to reach equilibrium o Economic factors:  Production of NH3 must have high rate of production due to economic considerations  Strong pipes & maintaining high-pressure reactor vessel very expensive – pressure mustn’t be too high  Plants should be located near supplies of component gases & haulage centres for cheap transport 2.8 Explain that the use of a catalyst will lower the reaction temperature required and identify the catalyst(s) used in the Haber process • Presence of catalyst provides alternate reaction pathway of lower EA o Increased rate of reaction – decreases time to reach equilibrium o Iron oxide catalyst (magnetite Fe3O4 fused with K2O, Al2O3 & CaO then reduced to porous iron)  Ground to fine powder to expose high surface area for gases to adsorb onto its surface  Allows lower temperatures & pressures to be used to achieve acceptable reaction rates without too great a loss in yield of product per cycle

2.9 Analyse the impact of increased pressure on the system involved in the Haber process • 4 moles of gaseous reactants are converted to 2 moles of gaseous products o Increase in pressure in system favours forward reaction – higher system pressure = higher yield o High gas pressure also favours frequency of successful collisions – higher reaction rate

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2.10 Explain why monitoring of the reaction vessel used in the Haber process is crucial and discuss the monitoring required

•

Chemists are employed to ensure production efficiencies & quality control o Continuous monitoring of high-pressure reaction vessel (25–35 MPa)  Ensures production of NH3 under safe conditions  If pressure too low yield of NH3 drops, pressure data from sensors monitored centrally o Temperature monitoring – critical to ensure temp remains in optimum range (400-500°C)  Temp too high, yield of NH3 reduced, temperature data from sensors monitored centrally o Monitoring furnaces that produce H2 & N2 feedstocks – must be produced in correct mole ratios (1:3) & kept free from contamination with sulfur compounds or CO, oxygen must be absent to prevent explosions  CO & CO2 sensors used to analyse information about gas stream o Monitoring activity of catalyst – lifetime of 5-10 years, particle size monitored to ensure high surface area o Monitoring ammonia liquefaction process during each cycle to ensure optimal yield of ammonia

2.0.1 Gather and process information from secondary sources to describe the conditions under which Haber developed the industrial synthesis of ammonia and evaluate its significance at that time in world history • Fritz Haber – patriotic German chemist, his experiments in 1905 produced small yields of ammonia when N2 & H2 gases combined at 1000°C over iron catalyst • Other catalysts tested – osmium best but very expensive • Further investigations showed pressure needed to be raised & temp needed to be lower to increase yields • 1909 – synthesised ~100g of ammonia • Carl Bosch scaled process up to industrial levels – yields of about 15% at 15-20 MPa & 500°C using ironbased catalyst • By continuously recycling unreacted gases yield could be raised to 98% • By end of 1914 production raised to 200 tonnes per day • Prior to WWI, Haber realised manufacture of nitric acid from oxidation of ammonia would be vital for production of nitrogenous fertilisers needed for agriculture • Synthetic fertiliser important in producing greater crop yields for increasing human population • Before the Haber process the global source of nitrates (for fertilisers and explosives) came from saltpetre from Chile • At start of WWI Britain & allies blockaded Atlantic sea routes from Chile – prevented Germany from importing nitrates – threatened to cause widespread starvation & Germany’s rapid defeat • Haber’s synthetic NH3 process allowed Germany to supply explosives for its war effort & to produce fertilisers to grow crops needed to feed large German population & its troops 8

• • • • •

The Haber process had a significant impact on the course of history during the early 20th century as it allowed Germany to continue the war for much longer than otherwise would have been possible Awarded Noble Prize – NH3 process improved nutrition of humanity, Bosch awarded Nobel Prize for enabling commercial production Haber, along with Einstein, was forced out of Germany by the Nazis – fled to USA & forced into retirement Haber process important industrial process, today ~85% of ammonia produced in the world used to manufacture fertilisers for farming Without synthetic fertilisers could not sustain growing population

3. Manufactured products, including food, drugs and household chemicals, are analysed to determine or ensure their chemical composition 3.1 Deduce the ions present in a sample from the results of tests • Qualitative tests – formation of gases or precipitates, depends on whether only 1 anion or a mixture o Reagent: a substance or solution that causes a reaction to occur o Confirmation test: additional tests used to confirm the original analysis o Solutions to be tested should have concentrations >0.1 molar, 1-2 mL samples of unknown solution placed in separate clean test tubes for testing o If unknown solution contains more than 1 ion, the 1st ion identified has to be removed (e.g. by precipitation of precipitate formed in excess reagent) before the next tests are conducted Soluble Salts* Group 1 and NH4+ compounds Nitrates Acetates (Ethanoates) Chlorides Sulfates (Lead Chloride) (Calcium / Silver Sulfate) (Silver chloride) (Barium / Lead Sulfate) Carbonates, Sulfites, Phosphates (except Group 1 and NH4+ compounds) Hydroxides & Oxides (except Group 1, NH4+, Ba2+, Sr2+) Sulfides (except Groups, 1, 2 & NH4+ compounds) * Solutions > 0.1 M can be prepared • Anion analysis – CO32-, Cl-, PO43-, SO42o All anions are colourless in solution – colour of solution due to cation(s) present o Anion elimination tests: Procedure A. Add 2M nitric acid until effervescence ceases A. Observation/conclusion Effervescence of colourless gas (CO2) indicates carbonate – use limewater to confirm CO32- (aq) + 2H+ (aq)  CO2 (g) + H2O (l) Should have alkaline pH (8-11) CO3 2- (aq) + H2O(l) HCO3- (aq) + OH- (aq) Thick white precipitate indicates SO42- present SO42- (aq) + Ba2+ (aq)  BaSO4 (s) White lead (II) sulfate precipitate forms Slightly Soluble Insoluble Salts

(Calcium Hydroxide)

Anion CO32-

SO42-

B. Confirmation test: test original solution with universal pH paper A. Add Ba(NO3)2 (aq) B. Confirmation test: to the acidified solution add drops of Pb(NO3)2 (aq) 9

B. A. B.

Cl-

PO43-

A. Filter solution from previous step, add AgNO3 A. White precipitate indicates Cl- ions Cl- (aq) + Ag+ (aq)  AgCl (s) (aq) to filtrate B. Confirmation test: add 1-2 mL of 2M NH3 (aq) B. White precipitate dissolves to form colourless to the suspension & heat in water bath silver complex ion AgCl (s) + 2NH3 (aq)  Ag(NH3)2+ (aq) + Cl- (aq) A. Add 6M NH3 (aq) to make filtered solution A. White precipitate indicates PO43- ions 2PO43- (aq) + 3Ba2+ (aq)  Ba3(PO4)2 (s) from previous step slightly alkaline (pH ~ 10), add Ba(NO3)2 (aq) B. Confirmation test: acidify solution with HNO3 B. Yellow precipitate forms & add ammonium molybdite, warm gently (NH4)3PO4.12MoO3.3H2O

•

Cation analysis – Pb2+, Ba2+, Ca2+, Cu2+, Fe2+/Fe3+ o Colour of solution – some cations have distinctive colours in solution  Fe3+ (aq) – yellow-orange to pale yellow solution  Fe2+ (aq) – pale green to colourless  Cu2+ (aq) – blue to green-blue o Flame test – many metal ions produce characteristic colours when their salts are volatised in a blue (nonluminous) Bunsen flame  When a metal salt is vapourised in a flame the outer shell electrons of the metal ion may absorb energy & move to higher ‘excited’ energy levels  Excited electrons are unstable & emit light of characteristic frequencies as they fall back to lower levels  Some of these frequencies are in visible EM spectrum

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Procedure – dip platinum wire into conc. HCl to clean, heat wire to remove impurities, dip wire into acid then into powdered salt, volatile chlorides of the metal ions are produced when wire is heated; or - Dissolve chloride salts in water & spray solution into flame using atomiser  Many chemicals are contaminated with Na+ ions – strong yellow flame colour so may mask colour of unknown metal  Useful for distinguishing between Ca2+ & Ba2+  Normally used as confirmatory tests rather than primary tests, same anion should be used for validity  Some metals cannot be tested using flame tests – poisonous & produce toxic vapours e.g. Pb2+ & iron Cation Ca2+ Ba2+ Cu2+ Sr2+ Na+ Li+ K+ o Cation elimination tests: Procedure A. Add 5 drops of HCl (aq) to unknown solution Observation/conclusion A. A (faint) white precipitate indicates Pb2+ ions (PbCl2 soluble in hot water but slightly soluble in cold water) Pb2+ (aq) + 2Cl- (aq)  PbCl2 (s) B. Yellow precipitate forms Pb2+ (aq) + 2I- (aq)  PbI2 (s) A. White precipitate indicates either Ca2+ or Ba2+ present (CaSO4 may not precipitate if too dilute) Ca2+ (aq) + SO42- (aq)  CaSO4 (s) Ba2+ (aq) + SO42- (aq)  BaSO4 (s) B. White precipitate confirms Ca2+, none suggest Ba2+ Ca2+ (aq) + 2F- (aq)  CaF2 (s) Brick-red flame indicates Ca2+, yellow green flame indicates Ba2+ A. Blue precipitate forms from an original blue or green solution, dissolves in excess NH3 to form deep blue solution containing Cu complex ion Cu2+ (aq) + 2OH- (aq)  Cu(OH)2 (s) Cu(OH)2 (s) + 4NH3 (aq)  Cu(NH3)42+ (aq) + 2OH- (aq) B. Green flame indicates Cu2+ ions A. Brown precipitate indicates Fe3+, green precipitate indicates Fe2+ (may turn brown rapidly) Fe3+ (aq) + 3OH- (aq)  Fe(OH)3 (s) Fe2+ (aq) + 2OH- (aq)  Fe(OH)2 (s) B. a. Dark blue precipitate indicates Fe2+ 3Fe2+ (aq) + 2Fe(CN)63- (aq)  Fe3(Fe(CN)6)2 (s) b. Deep blood-red colour indicates Fe3+ Fe3+ (aq) + SCN- (aq)  FeSCN2+ (aq) Cation Pb2+ Flame Colour Brick red Yellow-green (apple green) Green Scarlet Yellow Carmine (dull red) Violet (lilac)



Ba2+, Ca2+

B. Confirmation test: add sodium iodide solution to original solution A. Filter solution & add 5-10 drops of 1M H2SO4

B. Confirmation test: to fresh sample add NaF or conduct a flame test

Cu

2+

A. Add 1M NaOH to original solution, add NH3 (aq) when precipitate forms B. Confirmation test: use original solid or solution to conduct a flame test A. Ad 1M NaOH to original solution B. Confirmation test: add 10 drops of 4M HCl a. Remove 2 drops of solution to a white tile & add a drop of potassium hexacyanoferrate (III) b. To two other drops on a white tile add a drop of potassium thiocyanate

Fe2+, Fe3+

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Quantitative analysis – variety of techniques used to determine concentration of an element, ion or compound in a sample of material: o Gravimetric analysis – weighing materials & determining % composition of elements compounds or components of a mixture o Volumetric analysis – measuring the volumes of solutions that react with other solutions of known concentrations o Instrumental analysis – using instruments that determine the concentration or amount of a material by measuring a property of the material (e.g. pH, colour, absorbance of light)

3.0.1 Perform first-hand investigations to carry out a range of tests, including flame tests, to identify the following ions: - phosphate - carbonate - barium - lead - iron - sulfate - chloride - calcium - copper RISK ASSESSMENT:  AgNO3 is toxic if ingested, stains skin black  Pb(NO3) highly toxic cumulative poison  Concentrated HCl toxic, highly corrosive, lung/skin irritant  4 M NH3 used in fume hood  Quantities minimised, wash hands  Cu, Pb & Ag compounds collected for treatment & subsequent recycling or disposal METHOD: 1. Anion tests – each precipitation reaction was conducted in separate test tubes 2. Flame tests for cations – atomiser bottles, platinum wire cleaned with HCl, dipped in powdered salts 3. Cation tests – each precipitation reaction was conducted in separate test tubes, confirmation tests DISCUSSION:  Gas bubbles in the test for carbonate ions: acid + metal carbonate  salt + water + carbon dioxide gas  Silver nitrate test for chloride ions must be done in acidified solution to make sure that carbonate ions are not present − H3PO4 is a weak triprotic acid – each of its ionisation equilibriums will be shifted to the left and very minimal amounts of phosphate ions will remain in solution – Ag3PO4 will not form precipitate  Addition of Ba(NO3) in acidic conditions will precipitate sulfate ions − H2SO4 strong diprotic, 1st ionisation relatively unaffected ∴SO42- will still remain in solution − H3PO4 weak tripotic so won’t precipitate  Flame test for Fe3+ ions produced a yellow flame – not as definitive as precipitation reactions are for it 12

3.2 Describe the use of atomic absorption spectroscopy (AAS) in detecting concentrations of metal ions in solutions and assess its impact on scientific understanding of the effects of trace elements • Atomic vapours selectively absorb & emit various frequencies of light • If a sample of an element is vaporised in a hot flame, electrons are promoted from the ground state into more unstable or excited energy levels • As the electrons fall back to more stable levels they emit light of characteristic frequencies • If white light is passed through an atomic vapour at a suitably low temperature to prevent electron excitation, some wavelengths are selectively absorbed & dark lines appear in the spectrum produced o Correspond to the exact bright line wavelengths in atomic emissions spectra • Spectrometer: device that detects & measures a substance by its absorbance of certain wavelengths of light from the EM spectrum; also called spectrophotometer • AAS developed using principle of selective light absorption of metal ions – very sensitive technique o E.g. if nickel to be analysed, special nickel light passed through flame in which sample has been vaporised, Ni atoms if present in vaporised sample absorb some of the nickel light, this loss measured o Hollow-cathode lamp selection – light source of AAS, corresponds to element being measured  Generates specific wavelengths of light characteristic of element being analysed o Standard solution preparation of metal to be analysed – standard volumetric techniques used, solution then systematically diluted o Aspirating the solutions – dilution standards & unknown solution sprayed in turn (using a nebuliser) into the flame, alternatively sample can be heated in a graphite furnace  Flame AAS uses a slot type burner (~1000°C) to increase total absorbance of light  Graphite furnace (~3000°C) more efficient than flame method – can be used for smaller quantities of material & provides a reducing environment for samples that are readily oxidised o Measuring light absorption – as light beam passes through vaporised sample some of light absorbed by the hot atoms, a 2nd reference beam bypasses the sample  Emerging light beam passes through a monochromator which contains diffraction grating & focussing mirrors, light passes through a narrow slit to select only 1 of the wavelength bands to be measured  Light now monochromatic – light of a single (or very narrow) wavelength  Photomultiplier tubes measure light intensity & convert it into an electrical signal  Amount of light absorbed relative to reference beam (measured by the absorbance, A) related to the concentration, c, of element in the vaporised sample – greater the concentration = greater the absorbance A = kc  The constant k depends on the characteristic of the apparatus & the metal being analysed o Calibration – concentration determined from calibration curve with the standards of known concentration  Control blank (contains only solvent) is run – should register 0 absorbance

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Impact on scientific understanding of the effects of trace elements – large impact o Essential trace elements – elements required by living organisms in very small (trace) quantities (typically 1-100 ppm), in high doses can be toxic

o Plants absorb minerals from the soil When animals & humans eat these plants they absorb the essential micronutrients If soil lacks certain elements, humans or animals may exhibit deficiency diseases Significant health problems as a result o Existence of trace elements not known until sensitive analytical methods such as AAS developed o Old ‘wet’ (gravimetric or volumetric analysis) methods too insensitive & very time consuming, often failed when other ions present o Chemical analyses now took minutes instead of days with great accuracy, used on metal solids & liquids o AAS very specific & can determine concentration of a metal ion in the presence of other metals o AAS led to an understanding of importance of trace elements in crop growth & livestock health o Using AAS can quickly & reliably establish which trace metals are required for specific biochemical pathways – large impact on our understanding of the functioning of the body o Using AAS, discovered that trace elements have a variety of essential roles, e.g.  Fe – required for production & functioning of haemoglobin in the blood  Co – component of vitamin B12; vital for production of haemoglobin o Blood & urine samples in humans or sap samples in plants can be analysed by AAS o Used in farming to test soil to ensure soil is fertile    3.0.2 Gather, process and present information to describe and explain evidence for the need to monitor levels of one of the above ions in substances used in society • Lead – toxic heavy metal, dangerous even in small amounts, passed up food chain – bioconcentration o Over last 10 a significant increase in awareness about effects of lead on human health and the environment o Absorbed through ingestion, inhalation & other exposures o Only 2mg can be excreted per day, excess accumulates in the bones & teeth replacing natural calcium o Harms virtually every human system – especially brain, kidney and reproductive system o Long term exposure can cause anaemia, nervous system disorders, mental retardation, kidney disease & decreased fertility o No beneficial role to organisms, different species react differently to lead exposure o Does not decay, biodegrade or dissipate unless disturbed – remains a long-term pollution source o Major cause of illness in children and babies living in old homes painted using lead-based paints which contained up to 50% lead – young children who crawl on floor inhale dust and flakes o Uses in society:  Lead-paints banned in 1972 but many homes built before ban are likely to contain lead paint  Emissions, discharges, waste stockpiles and disposal practices from lead industries have resulted in wide-scale dispersion of lead into their local environment  Was used to reduce corrosion, increase the longevity of structures and reduce maintenance and replacement costs in structures such as bridges, towers and water tanks  Present in contaminated drinking water from lead pipes or solder, car batteries, roof flashing & lead shot  Over the last 10 000 years the amount of lead in the air has increased 20 times  Lead compounds (e.g. tetraethyl lead) used in leaded petrol to improve engine performance in older cars  Much of this lead was expelled into the air as exhaust gases – since 1986 cars designed to use unleaded o Need to monitor blood levels, water contamination, soil concentration & air pollution of lead o National goal for blood levels of lead can only be achieved by ensuring minimal contamination of the environment with lead and that existing contamination is cleaned up o Vital that levels of lead ions in substances used in society are monitored so that further contamination of the environment does not occur and so that the adverse effects of this contamination are not experienced 14

3.0.3 Identify data, plan, select equipment and perform firsthand investigations to measure the sulfate content of lawn fertiliser and explain the chemistry involved EQUIPMENT:  Lawn Fertiliser  Hotplate  Weighing tray (ammonium sulfate)  Watch glass  Magnetic stirrer  Concentrated HCl  Side armed conical flask  Acetone  Ba(NO3)2 solution (5%)  Sintered glass funnel  Electronic balance (porosity 4)  Beakers  Measuring cylinder METHOD: 1. 0.6g of fertiliser was weighed on clean, dry weighing tray, transferred to a 250mL beaker 2. 25mL of deionised water was added to dissolve the fertiliser 3. 5 drops of conc HCl was added to the solution to aid the dissolution 4. The beaker was placed on a hot plate and brought to the boil 5. 25mL of 5% barium nitrate solution was heated on a hot plate until it boiled 6. The hot Ba(NO3)2 was added to the hot fertiliser solution very slowly, stirred using a magnetic stirrer 7. Solution was allowed to cool for 5 minutes 8. 5mL of acetone was added to the beaker 9. Sintered glass funnel was weighed, the warm solution was then filtered using the funnel 10. The precipitate was washed with cold deionised water followed by a small amount of acetone to dry it 11. The funnel was reweighed, [SO42-] in the lawn fertiliser calculated using the mass of BaSO4 MBaSO4 = 233.37g nBaSO4 = 1.29/233.37 = 0.005527702 moles mSO4 = n x (32.07 + 4 x 16.00) = 0.531046407g % SO4 = 0.53/0.56 x 100 = 95% Theoretical % = 96.07/132.154 x 100 = 72.7% DISCUSSION          
2+

Mass (g) Fertiliser Filter Filter + sample BaSO4 (s) 0.56 98.93 100.22 1.29

Ba2+ (aq) + SO42- (aq)  BaSO4 (s) Ba & SO4 react to form insoluble BaSO4 – collected & measured to find the [SO42-] in the fertiliser [SO42-] found experimentally to be much higher than the theoretical value Conc. HCl added to aid the dissolution of the fertiliser Acidic conditions also prevented formation of barium salts of other anions such as chromate, carbonate & phosphate which are insoluble in neutral solutions Other salt ions adhering to the precipitate after filtration were washed from the residue with distilled water to prevent contamination (NH4)2SO4 heated & constantly stirred whilst adding the hot (BaNO3)2 slowly – helped to prevent sudden supersaturation & subsequent immediate precipitation Due to this the precipitate consisted of larger crystals due to coagulation – easier to filter Also helped to avoid coprecipitation of barium chloride Excess Ba(NO3)2 used to ensure complete precipitation Sintered glass funnel that was used had pores much smaller than that of ordinary filter paper – able to trap much more precipitate than would be obtained using ordinary filter paper.
2-

15

 

Acetone (CH3COCH3) is a polar aprotic solvent – a highly polar substance that dissolves ions but lacks an acidic hydrogen Helped increase the coagulation of BaSO4, then used as a drying agent due to the readiness with which it mixes with water and its high volatility

3.0.4 Analyse information to evaluate the reliability of the results of the above investigation and to propose solutions to problems encountered in the procedure RELIABILITY  As the experiment was not repeated reliability could not be gauged  By repeating this experiment and averaging the results reliability could be improved  By using more precise measuring devices the accuracy of the results could be increased  Validity questionable − Some of the precipitate was lost in the filtrate despite efforts to maximise the residue − On the other hand the mass of the precipitate was increased by not ensuring that the residue was completely dried before weighing  Impurities such as Ca2+ ions could have been present in the fertiliser sample – CaSO4 is slightly insoluble – could have formed a precipitate and contributed to the increased mass of the precipitate SOLUTIONS TO PROBLEMS  Numerous problems encountered in this experiment may explain the high % error  The major problem encountered in this procedure was the small particle size of barium sulfate  Drying the precipitate − Placing funnel into a low temp. oven or placing it into a dessicator & leaving overnight (more effective) − BaSO4 should be weighed several times and dried in between until a constant weight is obtained to ensure that all of the water is removed 3.0.5 Gather, process and present information to interpret secondary data from AAS measurements and evaluate the effectiveness of this in pollution control • E.g. AAS was used in Bangladesh to monitor levels of arsenic (extremely toxic) in drinking water o The following AAS absorbance data was obtained by using standard solutions of arsenic of known concentration: Total Arsenic Concentration Absorbance 50 μg/L 0.12 100 μg/L 0.23 150 μg/L 0.35

o This data was then graphed on a calibration curve:
• Samples of drinking water were taken and their absorbencies were measured: Sample Absorbance 1 0.28 2 0.13 3 0.31 These absorbencies were then plotted on the graph, and from there the arsenic concentrations were estimated: o Sample 1 (124 ppb) o Sample 2 (61 ppb) o Sample 3 (163 ppb) Arsenic levels higher than 100 ppb were considered polluted 16

•

•

•

Effectiveness in pollution control – very important tool – used to measure concentrations of pollutants in the atmosphere, soil & waterways o Advantages:  Identification measuring of low concentration pollutants (e.g. heavy metals) from many environments  Extremely accurate, quick & reliable o Disadvantages:  Cannot detect non-metal pollutants e.g. (CO 2 ), can only test for one ion at a time  Equipment expensive to purchase, however little cost after initial purchase  Destructive testing. o Based on ability for fast, efficient analysis of pollutants AAS is an extremely effective tool in pollution monitoring and management

4. Human activity has caused changes in the composition and the structure of the atmosphere. Chemists monitor these changes so that further damage can be limited • Free radical: atoms or molecules with one or more unpaired valence electrons 4.1 Describe the composition and layered structure of the atmosphere • Atmosphere is a thin gaseous layer that extends 600km above Earth’s surface, changes in temperature gradient mark boundaries between atmospheric layers • Composition essentially constant at all altitudes, although concentration of total gas particles drops with increasing altitude, amount of water vapour in atmosphere varies from 1-5% Layer Troposphere Altitude (km) 0 – 15  Air pressure decreases from ~100 kPa at surface to 10 kPa Composition Contains 75% (by mass) of atmosphere  78% N2, 21% O2, 0.9% Ar  25 NTU o Aquatic plants cannot photosynthesise if too little light penetrates the water – impacts food chain o Freshwater streams should not have total suspended solids > 20mg/L - acidity • Potable water should have pH between 6.5 & 8.5 • Acid rain can cause lakes to be below pH of 4 at which most aquatic organisms die • Acidic water leads to corrosion of metal pipes, alkaline water tastes bitter and produces scaling in pipes • pH > 9.0 aquatic organisms suffer from toxic ammonia poisoning as NH4+  NH3 • Acidic or alkaline waste discharges from industries can cause local changes in pH - dissolved oxygen and biochemical oxygen demand • Dissolved Oxygen (DO) – the amount of oxygen (mg) dissolved in 1 L of water at a fixed temp o High level of DO is vital for water quality – aquatic animals rely on DO for respiration o 1 at even 550°C − Heat exchanger  Gas pressure: pressure between 100 & 200 kPa increases collision frequency between reacting gases − Higher pressures (achieved by large fan blowers) increase yield as equilibrium driven to the right − Despite possibility of higher yield, low to avoid expense & safety risks of high-pressure apparatus  Oxygen concentration: excess of oxygen drives equilibrium right & increases yield − 5:1 mixture of air (21% O2) to SO2 is used – equivalent to mole ratio of O2:SO2 ~1:1 – double required 11
V2O5

1. Combustion of sulfur to form SO2

Catalyst: vanadium oxide catalyst used to increase reaction rate by providing alternate reaction pathway with lower EA − Catalyst impregnated onto surface of porous silica pellets – high surface area ensures rapid reaction  Removal of sulfur trioxide: before gas mixture enters 4th catalyst bed SO3 is removed by absorption into 98% H2SO4 in an ‘interpass absorption tower’ producing oleum which on dilution forms 98% H2SO4 − Removal helps shift equilibrium to the right, remaining gases pass through 4th catalyst bed − All but a trace of SO2 has been converted after passing over the 3rd or 4th catalyst bed – yield ~99.7% − Gases that are vented to the atmosphere are largely N2 & O2, no more than 0.3% SO2 3. Absorption of SO3 in H2SO4 to form oleum (H2S2O7) SO3 (g) + H2SO4 (l)  H2S2O7 (l) o Cooled SO3 from final catalyst bed dissolved in 98% H2SO4 in 2nd absorption tower  SO2 enters bottom of tower & H2SO4 sprayed at top down vertical plates that increase surface area of contact of the acid & gas  If SO3 dissolved directly into water stable sulfuric acid mists form due to large amount of heat generated − Very dangerous, difficult to handle, cannot be easily made to coalesce 4. Conversion of oleum to H2SO4 using water H2S2O7 (l) + H2O (l)  2H2SO4 (l) o Water mixed with oleum in diluter to produce 98% H2SO4 (18 mol/L)



3.4 Describe the reaction conditions necessary for the production of SO2 and SO3 3.5 Apply the relationship between rates of reaction and equilibrium conditions to the production of SO2 and SO3 3.0.1 Gather, process and present information from secondary sources to describe the steps and chemistry involved in the industrial production of H2SO4 and use available evidence to analyse the process to predict ways in which the output of sulfuric acid can be maximised • S (s) + O2 (g)  SO2 (g) ΔH = –297 kJ/mol o Not an equilibrium reaction – goes to completion, reaction rate increased using high temperatures o Air (source of O2) must be dry to avoid SO2 reacting with the moisture in the air to form sulfurous acid  Air scrubbed of water vapour by passing through 99% H2SO4 (dehydrating agent) o Excess of air (& oxygen) ensures sulfur reacts completely o Considerable heat released – gas stream must be cooled from about 1000°C to 400°C before next stage  Heat exchangers used, heat recycled to remelt more sulfur or to produce steam in waste heat boiler & power turbines to generate electricity for the factory 12 Production of SO2

o If SO2 is to be obtained from the metal industry it is cleaned, dried & heated to the correct temperature– conversion stage: 2ZnS(s) + 3O2(g)  2ZnO(s) + 2SO2(g)  Polluting SO2 is captured & provides a relatively cheap starting material for the sulfuric acid production • Production of SO3 SO2 (g) + ½O2 (g) ↔ SO3 (g) ΔH = –99 kJ/mol o Compromise between rate of reaction & yield of SO3 o Temperature: if too low rate of reaction decreases, compromise temperature (400-550°C) used as too high a temperature will reduce equilibrium yield (exothermic)  At 400°C, K = 100 so equilibrium lies far to the right – yield of SO3 98%, still >1 at even 550°C  Heat exchanger o Gas pressure: pressure between 100 & 200 kPa increases collision frequency between reacting gases  Higher pressures (achieved by large fan blowers) increase yield as equilibrium driven to the right  However, increase in yield that can be obtained by using high pressures is outweighed by the high costs and safety risks of high-pressure apparatus so a moderate pressure is used o Oxygen concentration: excess of oxygen drives equilibrium right & increases yield  5:1 mixture of air (21% O2) to SO2 is used – equivalent to mole ratio of O2:SO2 ~1:1 – double required o Catalyst: vanadium oxide catalyst used to increase reaction rate by providing alternate reaction pathway with lower EA  Catalyst impregnated onto surface of porous silica pellets – high surface area ensures rapid reaction  Gas stream passes over 4 catalyst beds, at each bed temp is lowered to increase yield  Unreacted gases are recycled o Removal of sulfur trioxide: before gas mixture enters 4th catalyst bed SO3 is removed by absorption into 98% H2SO4 in an ‘interpass absorption tower’ producing oleum which on dilution forms 98% H2SO4  Removal helps shift equilibrium to the right, remaining gases pass through 4th catalyst bed  All but a trace of SO2 has been converted after passing over the 3rd or 4th catalyst bed – yield ~99.7% 3.6 Describe, using examples, the reactions of sulfuric acid acting as: – an oxidising agent • Concentrated H2SO4 is a moderately strong oxidising agent (it itself is reduced), especially at high temps o SO42- ion is the oxidant – in hot concentrated H2SO4 it oxidises:  Unreactive metals such as Ag, Cu, Hg and Pb to produce the metal sulfate, SO2 & H2O, e.g. Sn (s) + 2H2SO4 (l)  SnSO4 (aq) + SO2 (g) + 2H2O (l) − Metals below lead in activity series are not oxidised by conc. sulfuric acid − In some cases oxidation ceases rapidly because insoluble product (e.g. PbSO4) coats the metal surface − Under various temps & with certain reactants S or H2S may be produced  Bromide or iodide ions to elemental bromine or iodine while the sulfuric acid is reduced to SO2 2KBr (s) + 3H2SO4 (l)  2KHSO4 (aq) + SO2 (g) + 2H2O (l) + Br2 (l) 2KI (s) + 3H2SO4 (l)  2KHSO4 (aq) + SO2 (g) + 2H2O (l) + I2 (l) − As it is only moderately strong it cannot oxidise chloride  Carbon, sulfur and phosphorous to CO2, SO2 & PO5 respectively • Dilute H2SO4 is a replacement acid – H3O+ ions oxidise reactive metals with the release of H2 gas, SO42ions is a spectator ion: Zn (s) + 2H3O+ (aq)  Zn2+ (aq) + H2 (g) + 2H2O (l) o Some reactive metals (e.g. Al, Ni, Cr) not readily attacked by dilute H2SO4 due to the passivating oxide layer on their surfaces – a dehydrating agent • Conc. H2SO4 (18 M) dries gases that do not react with it (e.g. O2, N2, Cl2, CO2) – used in the contact process • It’s loosely held proton transfers to water in a highly exothermic formation of hydronium ions • Hydrated crystals can be dehydrated by conc. H2SO4 o Blue copper (II) sulfate crystals turn white when conc. H2SO4 added CuSO4.5H2O (s)  CuSO4 (s) + 5H2O (l) 13 conc. H2SO4

V2O5

•

Carbohydrates, e.g. sugars, can be dehydrated by conc. H2SO4 o Sucrose, C12H22O11 readily dehydrated – black, porous solid produced is mainly carbon, the heat from the reaction causes water to be produced as steam which expands the black mass C12H22O11 (s)  12C (s) + 11H2O (g) Removes water from alkanols, e.g. converts ethanol to ethene C2H5OH (l) ↔ C2H4 (g) + H2O (l) conc. H2SO4 conc. H2SO4

•

3.0.2 Perform first-hand investigations to observe the reactions of sulfuric acid acting as: – an oxidising agent – a dehydrating agent METHOD Part A: Sulfuric acid as an oxidant 1. 3 mL of 2 mol/L sulfuric acid was added to a granule of tin in a test tube 2. Observations recorded for 5 minutes 3. Test tubes placed in a beaker of hot water and the effect on rate of reaction was observed The teacher performed the following experiments in a fume cupboard with small amounts of concentrated sulfuric acid – gloves and glasses worn when handling the sulfuric acid 1. A few crystals of potassium bromide were placed in a test tube & a few drops of conc. H2SO4 were added. 2. Tube placed in a test-tube rack, experiment repeated using a few crystals of potassium iodide 3. 2 mL of water and 1–2 mL of cyclohexane were added to the potassium iodide/sulfuric acid mixture 4. Tube stoppered and the mixture was agitated, the layers were allowed to separate Part B: Concentrated sulfuric acid as a dehydrating agent – in fume cupboard 1. A small scoop of fine blue copper sulfate crystals was placed in one Petri dish 2. Drops of concentrated sulfuric acid were added to some of the crystals in the dish, observations recorded 3. Sucrose was placed in a 150 mL beaker to a depth of 1 cm, a small amount of water was added 4. 10 mL of concentrated sulfuric acid was added, mixture stirred with a glass rod, observed for 5 minutes RESULTS Part A – H2SO4 Oxidising Agent Observations Slight effervescence, quicker in hot 1.a) water

Chemical Equation/Half-Equation Sn (s) + 2H2SO4 (aq)  SnSO4 (aq) + SO2 (g) + 2H2O (l) Oxidation: Sn (s)  Sn2+ + 2eReduction: SO42- + 4H+ + 2e-  SO2 (aq) + 2H2O (l) 2KBr (s) + 3H2SO4 (aq)  2KHSO4 (aq) + SO2 (g) + 2H2O (l) + Br2 (g) Oxidation: 2Br-  Br2 (s) + 2eReduction: SO42- + 4H+ + 2e-  SO2 (aq) + 2H2O (l) 2KI (s) + 3H2SO4 (aq)  2KHSO4 (aq) + SO2 (g) + 2H2O (l) + I2 (s) Oxidation: 2I-  I2 (s) + 2eReduction: SO42- + 4H+ + 2e-  SO2 (aq) + 2H2O (l)

2.a)

A brown colour indicated the formation of molecular bromine

b)

Darkening of the mixture indicated molecular iodine formation A violet, organic layer was characteristic of molecular iodine, layers separated

14

Part B – H2SO4 Dehydrating Agent Observations Turned white, dried up 1.

Chemical Equation CuSO4.5H2O (s)  CuSO4 (s) + 5H2O (l) Hydrated copper sulfate  anhydrous copper sulfate + water conc. H2SO4 C12H22O11 (s)  12C (s) + 11H2O (l) conc. H2SO4 conc. H2SO4

2.

Carbon rod formed, rapid effervescence, heat

3.7 Describe and explain the exothermic nature of sulfuric acid ionisation • The dilution & ionisation of concentrated sulfuric acid in water is highly exothermic H2SO4 (l) + H2O (l)  HSO4- (aq) + H3O+ (aq) HSO4- (aq) + H2O (l) ↔ SO42- (aq) + H3O+ (aq) H2SO4 (98%)  SO42- (aq) + 2H3O+ (aq) ΔH = –90 kJ/mol • Both dissociations exothermic – resulting solution becomes quite hot, may boil if amount of water is small • The ionisation of sulfuric acid is exothermic because the energy released when the hydronium ions are formed is much greater than the energy absorbed to break the bonds of sulfuric acid • Unlike concentrated H2SO4 (98%), concentrated HCl (35%) & HNO3 (70%) are prepared in much lower concentrations – most of the molecules already ionised so much less heat liberated upon dilution 3.8 Identify and describe safety precautions that must be taken when using and diluting concentrated sulfuric acid • When diluting sulfuric acid add small volumes of concentrated acid to large volumes of water, running it the down side of the container or stirring rod with constant stirring to disperse the heat, wear safety glasses • When pouring sulfuric acid into a beaker pour slowly down a glass rod to avoid splashing, use rod to stir mixture to distribute heat generated • Plastic trays can be used to ensure that any drips don’t contact the workbench while being used • If H2SO4 vapours are inhaled immediately move patient to fresh air, if breathing stops give artificial resuscitation, call for medical attention immediately • To protect skin & clothing always wear protective gloves & lab coat – skin & clothes will char rapidly • If sulfuric acid contacts the skin it must be washed off rapidly with copious amounts of running tap water o If larger amounts spilled on skin wipe excess away rapidly with paper towel then wash – minimises heat produced on dissolution with water • If sulfuric acid is spilt on bench or floor isolate to prevent from spreading, if fumes are present evacuate o Sand can be spread over acid to absorb it, place acid-soaked sand in plastic buckets for disposal & neutralisation, clean surface using solid sodium bicarbonate followed by water & detergent 3.0.3 Use available evidence to relate the properties of sulfuric acid to safety precautions necessary for its transport and storage o In concentrated form (98%) consists mainly of unionised hydrogen sulfate molecules o The small amount of water present is bound as a hydrate to hydrogen sulfate molecules, H2SO4.H2O o As water is added, protons are donated to the water molecules to form H3O+ o H2SO4 becomes increasingly ionised on dilution H2SO4 (l) + 2H2O (l) ↔ SO42- (aq) + 2H3O+ (aq) o At concentrations less than 0.001 mol/L the acid is essentially 100% ionised • Storing H2SO4 o Conc. H2SO4Can be stored in steel drums because low concentration of hydronium ions  Inside of drums coated with FeSO4 – passivating layer that protects steel against further attack o Dilute solutions normally stored in glass or plastic containers  Use well-sealed containers, smaller bottles more suitable & safer for regular use  Store bottles in secure, cool, ventilated room away from sunlight in plastic trays in case of breakage  Store away from metals (particularly metal powders), bases and water as will react exothermically 15

 Should not store with oxysalts (e.g. chlorates & nitrates) as unstable products are produced on mixing  Should not store in contact with wooden shelves – any spillage leads to exothermic dehydration • Transportation of H2SO4 – by road, rail or sea o Large volumes of conc. H2SO4 are stored in steel drums & tankers  Steel much stronger than glass or plastic – less likely to rupture in an accident  Acid does not attack drums as passivating layer forms  Water mustn’t be allowed to contaminate acid – significant ionisation & heat build up will occur, steel will be attacked by the acid & H2 gas will be released o Dilute acid more costly to transport – more acid is ionised, cannot be transported in steel, less product being transported 4. The industrial production of sodium hydroxide requires the use of electrolysis 4.1 Explain the difference between galvanic cells and electrolytic cells in terms of energy requirements • Catholyte: electrolyte present in the cathode compartment • Anolyte: electrolyte present in the anode compartment Galvanic Cells Use spontaneous reaction involving oxidation & reduction half-reactions to generate electrical energy • Chemical energy  electrical energy • Used is society as portable sources of electrical energy Simple cell constructed from 2 half-cells connected electrically • External circuit – metallic conductor joins electrodes • Internal circuit – ion transfer through a semipermeable membrane As electrons move through external circuit from anode to cathode, ions move through semipermeable membrane • Negative ions move from the cathode compartment (catholyte) into the anode compartment (anolyte) • Positive ions move in opposite direction to maintain charge neutrality Electrolytic Cells Use electrical energy to drive redox reactions in the non-spontaneous direction • Used to decompose compounds electrically into elements or simpler compounds Usually consist of one compartment • External source of DC power connected to the two electrodes

Structure

Electron flow

• •

Applied voltage pushes electrons onto negative cathode Electrons are removed by the power source from the positive anode

16

Cell Feature Oxidation at… Reduction at… Electron flow in external circuit Net cell reaction Electrical energy •

Galvanic Cell Anode (–) Cathode (+) From anode to cathode Spontaneous (Eө > 0) Produced

Electrolytic Cell Anode (+) Cathode (–) From negative battery terminal to cathode & from anode to positive battery terminal Non-spontaneous (Eө < 0) Required

Comparing galvanic & electrolytic cells Zn2+ (aq) + 2e- ↔ Zn (s) Eө = –0.76 V Cu2+ (aq) + 2e- ↔ Cu (s) Eө = +0.34 V o Simple galvanic cell Zn (s) | ZnSO4 (aq)|| Cu (s) | CuSO4 (aq)  Zn is stronger reductant & forms anode, Cu forms cathode  e- flow through external circuit from Zn anode to Cu cathode, K+ & NO3- move through salt bridge Zn (s) + Cu2+ (aq)  Zn2+ (aq) + Cu (s) Eө = +1.10 V  Cell potential is positive – spontaneous reaction o Same redox couples used to form an electrolytic cell  In this case 2 compartments so electrolytes don’t mix readily  External DC power source that provides a reverse voltage greater than 1.10 V will cause a redox reaction that is the reverse of the cell reaction in the galvanic cell  Zn deposited on the zinc cathode & copper anode will oxidise Eө = –1.10 V Zn2+ (aq) + Cu (s)  Zn (s) + Cu2+ (aq)

4.2 Outline the steps in the industrial production of sodium hydroxide from sodium chloride solution and describe the reaction in terms of net ionic and full formulae equations • NaOH, ‘caustic soda’, is industrially produced by the electrolysis of concentrated brine (salt water) o Steps: 1. NaCl solution obtained from sea water or underground salt mines 2. Impurities removed from saturated brine − Impurities would precipitate as [NaOH] rises & block pores in diaphragm and membrane cells − Ca2+ & Fe2+ removed by precipitation with sodium carbonate (Na2CO3) − Mg2+ & Fe3+ removed by precipitation with sodium hydroxide (NaOH) − SO42- removed by precipitation with CaCl2 or BaCl2 17

Ca2+ (aq) + CO32- (aq)  CaCO3 (s) Fe3+ (aq) + 3OH- (aq)  Fe(OH)3 (s) Ca2+ (aq) + SO42- (aq)  CaSO4 (s) 3. NaCl solution is concentrated forming a saturated brine (26% w/w) –formation of Cl2 gas favoured 4. Solution is electrolysed in a mercury, diaphragm or membrane cell − Cathode (-ve): 2H2O (l) + 2e-  H2 (g) + 2OH- (aq) –0.83 V − Anode (+ve): 2Cl- (aq)  Cl2 (g) + 2e–1.36 V − Net ionic: 2Cl- (aq) + 2H2O (l)  H2 (g) + Cl2 (g) + 2OH- (aq) − Overall: 2NaCl (aq) + 2H2O (l)  H2 (g) + Cl2 (g) + 2NaOH (aq) − Na+ is a spectator ion, water is reduced, chloride ions are oxidised − The oxidation of Cl- & H2O have similar Eө values, big difference between reduction of Na+ & H2O 5. Products separated − Chlorine & hydrogen gases are pumped out of cell, compressed & stored − Sodium hydroxide removed & stored 4.3 Distinguish between the three electrolysis methods used to extract sodium hydroxide: – mercury process – diaphragm process – membrane process by describing each process and analysing the technical and environmental difficulties involved in each process • 1890 manufactured using a 2-compartment electrolytic cell separated by a porous cement diaphragm o Not very efficient • 1895 new electrolytic cell established using mercury as a flowing cathode – NaOH made extremely pure • 1897 electrolytic cell developed that used a percolating asbestos diaphragm to separate the 2 compartments o NaOH produced was not as pure Feature Mercury Cell Diaphragm Cell Membrane Cell Cathode (–) Liquid mercury flowing Steel mesh Stainless steel mesh (or Ni) in long, steel trough 2Na+ (aq) + 2Hg (l) + 2e-  Reduction 2H2O (l) + 2e-  H2 (g) + 2OH- (aq) 2Na/Hg (l) half-equation Anode (+) Titanium plates Titanium (or titanium-steel alloys) Oxidation 2Cl (aq)  Cl2 (g) + 2ehalf-equation Overall cell 2Na+ (aq) + 2Cl- (aq)  2Na (s) + Cl2 (g) 2NaCl (aq) + 2H2O (l)  H2 (g) + Cl2 (g) + 2NaOH (aq) equation Decomposer:
2Na/Hg (l) + 2H2O (l)  2NaOH (aq) + H2 (g) + 2Hg (l)

Electrolyte

Anolyte: saturated brine Catholyte: pure water (or 32% NaOH) Partition None PTFE (teflon) polymer – metal Thin, porous polymer membrane oxide diaphragm (replacing coated with anionic groups – porous asbestos diaphragm) impermeable to anions but not Na+ • The Mercury Cell – phased out due to concerns about mercury contamination of environment o Designed to ensure products of electrolysis are kept separate – H2 & Cl2 are explosive mixtures o Na metal rapidly dissolves in the mercury to form a sodium-mercury amalgam (alloy) o Na amalgam flows into ‘decomposer’ – contains pure water o Na amalgam decomposes – reacts with H2O at surface of graphite catalyst balls in decomposer chamber 2Na/Hg (l) + 2H2O (l)  2NaOH (aq) + H2 (g) + 2Hg (l) o Mercury recycled

Saturated brine

Purified saturated brine

18

•

Diaphragm cell – most common cell used in industry o Consists of 2 compartments separated by porous diaphragm – traditionally composed of asbestos but replaced by safer & superior diaphragms, 3mm thick, permeable to both cations & anions o Allows movement of Na+ ions and reduces migration of Cl- & OH- ions o Electrolyte on both sides of cell is concentrated brine (26% w/w)  Fresh brine constantly replaces spent brine at the anolyte, preventing dissolved Cl2 diffusing into catholyte  Level of anolyte above catholyte level to reduce mixing of OH- & Cl diffusion – Cl2 combines with NaOH to form sodium hypochlorite o After electrolysis, catholyte contains 10% NaOH & 15% NaCl o Unwanted NaCl solution removed – concentrated using heat energy from steam then fractionally crystallised by cooling  NaCl less soluble than NaOH, crystallises preferentially o Leaves supernatant of 50% NaOH & 1-2% NaCl o The electrolytic cells are quite small so large numbers are assembled in series

•

Membrane cell – similar to diaphragm cell o Catholyte contains no brine – pure water or dilute NaOH used o Ion selective membranes prevent Cl- & OHdiffusion, Na+ can diffuse from anolyte into catholyte  Membrane is a thin, porous polymer coated with anionic groups that repel anions but not cations o NaOH solution formed in cathode half-cell is concentrated (in a falling-film evaporator) to 50% using heat from steam, salt impurity of only 0.02% 19

•

Technical & environmental issues o Energy & operating costs  Electrolytic cells used to make NaOH & Cl2 use currents up to 80 000 A to ensure rapid production  Cost of electricity a major technical issue  Electricity costs reduced by using high purity salt – impurities reduce efficiency  Also reduced by maintaining anodes & their surfaces, ensuring brine is kept hot o Pollution & health issues  Chlor-alkali industry requires burning vast amounts of fossil fuels for electricity − CO2 contributes to global warming  Cl2 is a toxic gas, H2 will react explosively with Cl2 or O2  Corrosive action of NaCl, Cl2 and NaOH o Product purity & quality control  Quality of feedstocks, brine solutions & final products need to be controlled; monitored by − Acid-base titration to monitor alkalinity of NaOH − AAS to determine level of Ca2+ & Mg2+ impurities in brine feedstocks − Gravimetric analysis to determine moisture content of Cl2 produced − Ion-selective electrodes to monitor concentrations of ions, e.g. Cl-, in solution o Choosing an industrial site  Location of markets – new plant should be sited near other industries that use products  Energy – power must be readily available, energy companies often provide lower power cost incentives to new industries  Raw materials & feedstocks – brine or salt must be easily available, many plants located on coast  Transport networks – must be close to road, rail & shipping  Scientists, technicians & site workers – large workforce required, many need tertiary qualifications Mercury Cell Diaphragm Cell 3.5-4.5 V 2720 (energy consumption for steam producing concentrated NaOH 50% = 610 kWh/Mt Cl2) Membrane Cell 3.0-3.6 V 2650 (energy consumption for steam producing concentrated NaOH = 180 kWh/Mt Cl2) • Lowest energy consumption • Lower establishment & operating costs o Process not affected by variation in individual cell loads or shutdowns • High cost of polymer membranes • Voltages reduced by placing anode close to membrane o Reduces cell’s resistance

Energy & operating cost Working voltage Energy consumption (kWh / Mt Cl2) • •

3.9-4.2 V 3360

Highest energy consumption & working voltage Mercury is a very expensive metal

20

Pollution & health issues • Health risks associated with • Original diaphragm cell use of toxic mercury used toxic asbestos o Although it is recycled diaphragm there is a mechanical loss o Health conditions which to the environment arise from exposure to o Waste water containing airborne asbestos fibres mercury converted to toxic are mesothelioma and methyl-mercury asbestosis o Bioaccumulation • Brine waste from diaphragm cell may be contaminated with hypochlorite o Must be removed before discharge into environment Product purity & quality control • Highest purity NaOH as • Least purity NaOH sodium is reacted with pure • NaOH produced as 50% water solution (after • NaOH produced as 50% evaporation) with 1-2% solution salt impurity Other technical issues • A leak in the diaphragm • Escape of mercury o To minimise mercury loss would cause H2 to react recycled mercury is explosively Cl2 or O2 washed through special o Must carefully monitor mercury trap where it is equipment precipitated as HgS, Hg • Minimising contact vapour trapped and between OH- & Cl- in condensed solution to prevent • Large electric currents formation of hypochlorite create heating and magnetic • Minimising the amount of effects chloride that is present in the final NaOH • Large electric currents create heating and magnetic effects Additional information • Mercury plants being phased • Asbestos diaphragms out banned • More efficient & inert polymer-oxide composites developed

•

High purity NaOH (0.02% salt impurity)

• •

The high oxygen content of chlorine is a disadvantage Most expensive

•

Gradually replacing both mercury & diaphragm cells as produces high purity NaOH & requires less energy

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4.0.1 Identify data, plan and perform a first-hand investigation to identify the products of the electrolysis of sodium chloride RISK ASSESSMENT: Cl2 highly toxic by all routes of exposure, H2 highly flammable, NaOH corrosive  Well ventilated area, part 1: 2.0 M solution conducted in fume cupboard  Part 2: teacher demonstration in fume cupboard METHOD:  Part 1 – Student experiment 1. 50 mL of 0.001 mol/L NaCl solution was placed in a 150 mL beaker 2. 2-3 mL of universal indicator was added to the solution 3. The electrodes were inserted on either side of the beaker, connected to a DC transformer 4. Solution was electrolysed with a voltage of 4 V leaving undisturbed 5. Evolution of gases and colour change of indicator noted over 10 minutes 6. After 10 minutes the electrodes were removed and the solution was stirred with a glass rod 7. The final colour of the solution was recorded 8. The same experiment was conducted concurrently with 0.1 mol/L and 2.0 mol/L NaCl solutions  Part 2 – Teacher demonstration 1. A 2.0 M NaCl solution was coloured with sufficient universal indicator to produce a deep green colour. 2. A clean voltameter set up in a fume cupboard was filled with this solution 3. It was ensured that no air was present in either gas collection tube 4. NaCl solution electrolysed with 8 V to ensure a rapid rate of gas collection in each tube of voltameter 5. The colour change of the indicator in each arm of the voltameter was observed 6. When sufficient gas had been collected the current was turned off 7. A test tube of each gas was collected by downward air displacement 8. The gas from the negative terminal was tested for hydrogen using a lighted taper 9. Moist pieces of red and blue litmus paper were placed in the gas from the positive electrode 10. The tube was stoppered and the change in colour of the litmus was observed RESULTS:  Part 1
Concentration of NaCl (mol L-1) 0.001 0.1 2.0 Rate of effervescence Slow Moderate Rapid Final colour of solution Colourless Purple Colour at anode Red Red Red Pop test -------Positive result Colour at cathode Violet Violet Violet Final colour Green Yellow Colourless



Part 2
Electrode Anode (+) Cathode (–) Litmus paper Turned white --------

DISCUSSION:  Anode equations:

2Cl- (aq)  Cl2 (g) + 2e–1.36 V H2O (l)  ½ O2 (g) + 2H+ (aq) + 2e–1.23 V  Cathode equation: 2H2O (l) + 2e  H2 (g) + 2OH (aq) –0.83 V  Bleaching of solution & litmus  Cl2 produced at anode – higher concentrations favour production of Cl2  Indicator turned red near anode due to production of H+ ions from oxidation of water – pH lowered  Indicator turned violet near cathode due to production of OH- ions from reduction of water – pH raised  Positive pop test  hydrogen gas produced at cathode CONCLUSION:  Different products formed with different concentrations of NaCl  Dilute NaCl – O2 & H+ produced at anode, H2 & OH- produced at cathode  Concentrated NaCl – Cl2 produced at anode, H2 & OH- produced at cathode

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4.0.2 Analyse information from secondary sources to predict and explain the different products of the electrolysis of aqueous and molten sodium chloride • Electrolysis of molten NaCl – The Down’s Cell o Electrolysis of molten salts can be used to produce elements o Sodium metal produced commercially by electrolysis of molten NaCl in electrolytic cell using 30 000 A  The Down’s cell does not use pure NaCl as it has too high a melting point (801°C)  CaCl2 added to produce 2:1 mole mixture that melts ~600°C – acts as a flux  Lower temperature of molten mixture reduces formation of sodium metal mists  Only sodium & chloride ions present therefore only 1 oxidation reaction & 1 reduction reaction  Cl- oxidised at graphite anode to form Cl2 gas – pumped from cell, compressed & stored 2Cl- (l)  Cl2 (g) + 2e+  Na reduced at iron cathode which forms a ring around the graphite anode Na+ (l) + e-  Na (l)  Molten sodium that forms is less dense than the molten salt – as Na collects at surface it overflows into collection vessel 2NaCl (l)  Cl2 (g) + 2Na (s) • Electrolysis of aqueous NaCl is more complicated due to presence of water o Reaction which has most positive Eө value is most likely to proceed Cathode (-ve): 2H2O (l) + 2e-  H2 (g) + 2OH- (aq) Na+ (aq) + e-  Na (s) 2Cl- (aq)  Cl2 (g) + 2e2H2O (l)  O2 (g) + 4H+ (aq) + 4e–0.83 V –2.71 V –1.36 V –1.23 V

Anode (+ve):

o Big difference between Eө values for reduction of Na+ & H2O – water is the only species reduced o The oxidation of Cl- & H2O have similar Eө values – both species are reduced due to overpotential  Oxidation of water favoured in dilute NaCl solutions (2 M) , 5. Saponification is an important organic industrial process • • • Fatty acid: a carboxylic acid molecule with a long hydrocarbon chain (usually 12-20 carbon atoms long) Many common fats & oils are classified chemically as triglycerides or fatty esters Fatty esters have a more complex structure than simple esters o Consist of a condensed alkantriol called glycerol o Glycerol (C3H5(OH)3) has 3 long-chain carboxylic acid molecules (commonly called fatty acids) o In some fats each fatty acid chain is the same, in others the chains have different structures o Fatty acids can be classified as saturated or unsaturated based on presence/absence of C=C  Saturated if no C=C, unsaturated if C=C present  Solid fats tend to contain saturated fatty acids while oils tend to contain unsaturated fatty acids

5.1 Describe saponification as the conversion in basic solution of fats and oils to glycerol and salts of fatty acids • Saponification: the chemical reaction in which fatty esters (fats or oils) are hydrolysed in alkaline solution to produce glycerol (1, 2, 3-propanetriol) and salts of fatty acids – soap o Selected lipid is mixed with an alkaline solution & heated – alkali attacks the fat or oil molecules  Produces salts of the fatty acids present in the lipid, glycerine (or glycerol) is produced as a by-product  Fatty acid salts are called ‘soap’ o Soaps produced from vegetable oils tend to feel less greasy due to their shorter hydrocarbon chains 23

5.2 Describe the conditions under which saponification can be performed in the school laboratory and compare these with industrial preparation of soap • Preparation of soap in a school laboratory o Safety glasses & protecting clothing essential due to hot, concentrated alkaline solution o Procedure:  ~10g of fat weighed into a 500 mL beaker  25 mL of water & 25 mL of methylated spirits mixed in a flask, ~12g NaOH dissolved in mixture  NaOH solution stirred into the fat in the beaker (alcohol helps dissolution of fat into the aqueous phase & allows it to react faster with the alkali)  Beaker covered with a clock glass & heated on hotplate for 30mins, stirring occasionally  More water-methylated spirits added periodically to maintain volume of liquid  No oil droplets visible when reaction complete, beaker removed form heat & allowed to cool  Saturated salt water used to ‘salt out’ the soap (much less soluble in salt water) – forms thick curds on top of mixture  Vacuum filtration, washed with water to remove any alkali, crude soap collected & tested • Industrial preparation of soap o Kettle boiled batch process  Performed in immense steel containers – kettles, fats or oils blended & mixed with conc. NaOH  Some soap & salt from previous batch is left in kettle to assist with mixing as it promotes formation of emulsions, salt also assists in solidification of the soap as it forms  Mixture boiled using high-pressure & high-temperature steam that emerges through injectors deep inside the kettles, injection also assists mixing  Hot brine & steam added at end of saponification reaction to ‘salt out’ soap & wash it free of glycerol  Mixture settles for several days, soap curd gradually collects at surface  Lower aqueous layer containing dissolved glycerol is removed, brine recycled  Glycerol extracted by neutralising remaining alkali then distilling solution to remove water, followed by low-pressure distillation of the glycerol  Soap curd washed with water to remove excess salt & alkali, then spray & vacuum dried  End of process – soap contains up to 12% water  Converted into small pellets that are combined with fragrance & colours before reblending

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• • • Differences • • • • • Similarities

Fats & oils are mixed with concentrated alkali and heated Concentrated brine is used to separate the soap from the aqueous phase The crude soap is washed A blend of fats & oils is used in industry rather than 1 fat or oil High-pressure steam is used to heat & stir the mixture instead of glass rod & hotplate Glycerol is removed and purified in industry, some remains in soap in laboratory Settling of soap occurs over several days, in lab soap is relatively crude & collected rapidly Some old soap is left in the kettle to emulsify the new reactants for the next batch, in the lab methylated spirits is added to help emulsification • No fragrances or colours were added to soap produced in laboratory o Continuous saponification process – faster but must be done on smaller scale  Capacity of plant divided into 3 to 6 separate kettles  Saponification performed continuously in a reactor fitted with a recycling system  Liquid-liquid extractor used to separate the soap, glycerine & alkali solution  Process only takes a few hours, less space required on factory floor, reduced losses of feedstocks  Less steam required per tonne of soap but energy consumption is much higher  Economical for soap production only if a minimum of 50 tonnes is produced per day o Fatty acid neutralisation process  High pressures (~5 MPa) & temperatures (~250°C) used to break down fatty esters into fatty acids & glycerol in long steel tubes with zinc oxide catalyst  Fats or oils react with the steam forming glycerol & fatty acids  Removing glycerol helps increase yield as equilibrium is driven to the right fat + water(steam) ↔ glycerol + fatty acids  Fatty acids purified & fractionally distilled to obtain different boiling point fractions  Various fatty acid fractions are stoichiometrically neutralised with bases to produce soap fatty acid + sodium hydroxide  soap + water  Neutralisation is exothermic so conditions adjusted to ensure mixture does not boil, soap salted out

5.3 Account for the cleaning action of soap by describing its structure • Soap is a surfactant (‘surface acting agent’) – a substance that alters the physical properties of a surface o Soap is an ionic compound  Positive ion usually a sodium or potassium ion  Negative ion is a long hydrocarbon chain called the tail & a carboxylate end group called the head o Unique structure responsible for the special properties  When dissolved in water the cations separate from the fatty carboxylate anions  Long non-polar hydrocarbon chain is hydrophobic (little or no affinity with water) 25

 

Charged head group is hydrophilic (interacts strongly with water) In dilute solution soap anion moves to surface of water where hydrocarbon chains can form oily layer with the negative head groups attracting the water dipoles below  Presence of soap at interface interferes with normal hydrogen bonding between water molecules − Surface tension lowered – less ability to form droplets  Above a specific concentration of soap, clusters (micelles) of 40-100 soap anions form spontaneously in the bulk of the liquid − The long hydrocarbon chains form an oily central core where dispersion forces stabilise the chains − Negative head groups are at surface of the micelle – interact through ion-dipole attractions with positive ends of the water dipoles

•

The cleaning action of soap o A greasy stain can be removed from a piece of material or your skin using the properties of a surfactant  Same process whether using soap or a synthetic detergent o Soap enters water – positive ion dissociates in water o Interaction of soap & grease – when stained clothing added to soap-water mixture the long non-polar hydrocarbon tails of the soap start to dissolve in the greasy stain – interact through dispersion forces  Charged heads remain at the surface of the grease – can interact with the water molecules o Formation & stabilisation of grease-soap micelles  Interaction between long hydrocarbon tails with oil droplet & anion with water molecules are strong enough to lift grease layer from material – helped by agitation  Droplets become stabilised in the water by the formation of micelles  Stabilised mixture of grease, water & soap is an emulsion  Surface of grease droplets is covered with the negatively charged head groups of the soap  Negative droplet surfaces repel one another so droplets don’t tend to coalesce or join together again  Hydrocarbon tails are buried deep in the greasy droplet  Emulsification assisted by agitation & using hot water  If excess soap added some of soap anions will be present as soap micelles  Soap will also form a lather consisting of soap, water & air – helps suspend particles of grease o Rinsing away the emulsion  Lather & soapy emulsion of grease & water rinsed away from material with fresh water

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5.4 Explain that soap, water and oil together form an emulsion with the soap acting as an emulsifier • Emulsion: a dispersion of small droplets of one liquid within another liquid, e.g. soap + water + oil o When 2 insoluble liquids mix they often form 2 layers  If oil mixed in water by vigorous agitation temporary dispersion formed  Eventually oil droplets join together to form larger droplets that rise to the surface, form a separate layer o Surfactants can be used to stabilised the dispersion o If soap, an emulsifier (allows substances which were previously immiscible to mix), is added mixture does not separate due to the formation of an emulsion o Hydrocarbon tails of soap anions dissolve in the oil & promote droplet formation o Surface of droplets is covered in negatively charged heads so droplets repel each other o Micelles stabilised since the soap dissolved in the oil interacts with the water around it

o 2 types of emulsions  Oil-in-water emulsions – consist of oil colloidal particles dispersed in water (the continuous phase) − Greater proportion of water than oil so dissolve better in polar solvents, e.g. water, rather than in organic solvents, e.g. kerosene − Non-greasy feel  Water-in-oil emulsions – consist of colloidal droplets of water dispersed in oil (the continuous phase) − Feel much greasier than oil-in-water emulsions − Mix better with organic solvents  Some emulsions can be converted into the opposite type – e.g. cream churned into butter

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5.5 Distinguish between soaps and synthetic detergents in terms of: - the structure of the molecule - chemical composition - effect in hard water • Detergents are cleaning agents, classed as anionic, cationic or non-ionic • Soaps, a subset of anionic detergents, are non-petrochemical derived • Any surfactant that does not belong to this subset is a synthetic detergent Soaps o Soluble salts of long, naturally occurring non-polar fatty acids  Soaps are surfactants  Consist of positive ion (usually Na+ or K+) and hydrocarbon chain with carboxylate end group (R-COO-)  E.g. sodium stearate CH3-(CH2)16-COO-Na Detergents o Hydrocarbon surfactants  Anionic − Alky sulfates/sulfonates are the most common  Cationic − Usually ammonium compounds − Negative ion often a halide ion, e.g. chloride or bromide  Non-ionic − Most common are low molecular weight polymers consisting of 5-50 ethylene oxide monomers − Alcohol functional groups commonly present − Oxygen atoms form hydrogen bonds with water o Long, non-polar hydrocarbon chain obtained from petroleum  Hydrophobic, oil soluble  Often contain side chains

Chemical composition

Structure Tail

o Long, non-polar hydrocarbon chain  Hydrophobic, oil soluble  May be: − Monounsaturated (1 C=C bond) − Polyunsaturated (more than 1 C=C bond) − Saturated (no C=C bonds) Head o Negatively charged carboxylate group (COO-)  Hydrophilic

Behaviour in hard water

Effect of pH

o Soap will produce a lather of bubbles only in soft water o In hard water soap anions complex with metal ions (Ca2+ & Mg2+), forming an insoluble scum that does not lather o Carboxylate group protonated in acidic solutions (pH < 4.5)  Forms uncharged, insoluble molecule which is not surface-active

o Hydrophilic o Anionic – negatively charged  E.g. sulfate (SO42-) or sulfanoate (SO3-) ion o Cationic – positively charged  E.g. ammonium (NH4+) ion o Non-ionic – neutral, hydrophilic groups  E.g. ethoxy groups (-CH2-CH2-O-) o Most detergents lather in hard water  Some anionic detergents (alkyl sulfates & alkyl phosphates) do complex with Ca2+ and Mg2+, forming a scum o Alkyl sulfates & alkyl phosphates will not precipitate out in low pH solutions o Ammonium salts are deprotonated to amine in high pHs (pH > 10)  Forms uncharged, insoluble molecule which is not surface-active

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Diagrams

Anionic

Sodium stearate

Sodium dodecyl sulfate (SDS)

Cationic

Cetrimonium bromide (CTAB)

Non-ionic

Polyethylene oxide chain

5.6 Distinguish between anionic, cationic and non-ionic synthetic detergents in terms of: - chemical composition - uses • Chemical composition: see above

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•

Uses of synthetic detergents

Anionic Cationic Non-ionic • Anionic detergents are • Cationic head group binds strongly • Used in: o Paints, cosmetic & adhesive where strongly foaming to negatively charged surfaces improved contact between polar & (unless pH is above 10) • Nearly 50% of all non-polar substances is required synthetic detergents are • Used in: o Emulsion polymerisation o Fabric softeners & conditioners anionic o Low-foaming cleaners o Hair conditioners • Used in: o Can be combined with other  Once they bond to the fabric or o Laundry detergents classes of detergent (e.g. anionic hair surface the long  Effective in removing laundry detergents) hydrocarbon tails reduce static grease & stains from  Provides greater stability during friction & tangling of fibres natural fibres the action of the detergent o Used to clean plastic  Bubbles hold dirt  Improves ability of laundry o Used in water treatment works particles away from detergent to clean synthetics as flocculating agents the fabric until they o Blended with cationic detergents o Mild biocidal properties can be rinsed away to provide greater stability over a  Antiseptic properties make o Excellent for cleaning wider pH range them useful in household glass o Used in addition with a caustic disinfectants & sanitisers o In heavy-duty engine cleanser in automatic dishwashers including antiseptic soaps, lubricants  Doesn’t produce too much foam mouthwashes & lozenges for o In personal care o Combined in shampoos with sore throats products, shampoos, anionic detergents to emulsify the o Not used in dishwashing liquids and dishwashing & natural oil produced by hair as they leave the negatively laundry liquids follicles charged glass with a greasy feel o Insecticides, herbicides & due to the adsorbed detergent pharmaceuticals o Useful for creating antistatic o Used in hard water as they don’t surfaces on glass though precipitate 5.0.1 Perform a first-hand investigation to carry out saponification and test the product RISK ASSESSMENT:  NaOH extremely toxic if ingested, highly corrosive to skin & eyes, exothermic dissolution  Methylated spirits toxic if ingested, highly flammable (cannot use Bunsen burner)  Gloves, lab coats & glasses must be worn – risk of spattering – continuously stir, boiling chips METHOD: 1. 10 mL of oil (e.g. olive oil) & 50mL of NaOH/ethanol solution was placed in a 500mL beaker 2. Boiling chips added, beaker covered with a clock glass, gently boiled with heating mantle for 30 mins a. Stirred regularly, additional water added to replace water lost by evaporation 3. Removed from heat, 20 mL of saturated NaCl solution added to salt out the soap 4. Soap forms solid layer on top, glycerol dissolves in the lower ‘brine’ layer 5. Soap filtered out using Buchner funnel & vacuum filtration 6. Crude soap washed with water to remove glycerol & excess NaOH 7. Small amount of soap added to 5 mL of hard water (0.1 M CaCl2) & to tap water in 2 test tubes 8. Tubes stoppered & shaken – lathering ability compared with sample of commercial dishwashing liquid 9. Some soap dissolved in warm water in small beaker 10. 5 mL of soap solution poured into one test tube, 5 mL of tap water poured into 2nd test tube 11. 5 drops of olive oil added to each tube to produce thin, visible layer of oil on top 12. Tubes stoppered & shaken, mixtures allowed to settle RESULTS:  Soap lathered in tap water but formed precipitate in hard water, detergent lathered in both  Oil & water emulsion rapidly separated, emulsion stabilised by presence of soap 30

5.0.2 Gather, process and present information from secondary sources to identify a range of fats and oils used for soap-making • Fats & Oils Used in Soap Making Fat or Oil Properties Tallow • Very common animal fat derived from beef processing • Produces hard, greasy soap unless other oils are blended with it before alkaline hydrolysis Lard • Derived from pigs • Produces hard soap that lathers quickly but does not readily dissolve in water Coconut • Derived from pressed, dried fruit of coconut palm Oil • Produces soaps that lather in salty or hard water, often blended with tallow to produce softer soap that dissolves faster in water Palm Oil • Derived from flesh of the palm fruit • Long hydrocarbon chains give it properties more similar to tallow than vegetable oils Palm • Derived from the kernels of palm tree seeds Kernel Oil • Fatty acid composition similar to that of coconut oil Olive Oil • Derived from crushed fruits of olive trees • Very high percentage of unsaturated fatty acid 5.0.3 Perform a first-hand investigation to gather information and describe the properties of a named emulsion and relate these properties to its uses AIM: To test the properties of milk (an oil-in-water emulsion) METHOD: 1. Drops of milk were added to a small beaker of water to produce a cloudy mixture 2. A light ray (from ray box) was shone through the mixture 3. Observed from above as well as looking towards the light source 4. Compared observations with light shone through transparent solution (NaCl solution) as a control 5. Sample of milk warmed in a test tube in hot water bath, small amount of white vinegar added, observations recorded RESULTS:  Part 1: colloidal particles in milk scattered light ray, no scattering in transparent solution  Part 2: solid & water phases of the emulsion separated DISCUSSION:  Emulsifying agents are added to the milk to prevent the fat from forming a separate layer  Casein (protein in milk) micelles scattered shorter wavelengths of visible light but allow longer wavelengths to be transmitted  Acid caused casein micelles to destabilise & aggregate by decreasing their surface charges  Properties: milk is composed of an emulsion of fat globules (surrounded by a thin membrane which helps to stabilise them in an emulsion) & a suspension of casein micelles in water − Fat droplets suspended throughout the milk rather than in a layer at the bottom − Makes it useful as a drink as it looks & tastes better − Milk fat is the major source of lipid used by mammalian newborn for accumulating fatty tissue − Micelle structure of milk casein important for digestion in the stomach 5.0.4 Perform a first-hand investigation to demonstrate the effect of soap as an emulsifier METHOD: 1. Vegetable oil and water mixed in test tube with and without soap, repeated with hard water RESULTS  When soap was added to the water + oil mixture, a cloudy emulsion was observed – tiny oil droplets were suspended in the soapy water  No emulsion without soap  Soap scum formed in hard water, no emulsion 31

5.0.5 Solve problems and use available evidence to discuss, using examples, the environmental impacts of the use of soaps and detergents • Biodegradability o Anionic detergents, including soap, are usually fairly biodegradable and can be precipitated out with cations due to their negative charge  Soap most biodegradable – readily broken down to CO2 & H2O by organisms such as bacteria in sewage works & natural waterways  Ability to be precipitated can lead to clogged sewage reticulation systems and problems in kettles, boilers and washing when in the presence of hard water o Cationic detergents can kill some of the microbes that biodegrade the detergent – least biodegradable o Earliest synthetic detergents were anionic alkylbenzene sulfanoates with highly branched hydrocarbon tails.  Persisted in environment in as they took around a month to decompose, causing many waterways to build up massive heads of foam  Problem solved by synthesising biodegradable detergents with non-branching tails such as linear alkylbenzene sulfanoates • Presence of phosphates o While synthetic detergents are more successful in hard water than soap, the presence of Ca2+ & Mg2+ ions causes small colloidal particles to flocculate soiling clothes in the wash  Phosphate ‘builders’ were added to detergent powders to prevent – builders created alkaline conditions and complexed these ions, preventing them from interfering in the washing process o The phosphate residues were carried into natural bodies of water, leading to algal blooms and eutrophication in developed countries o ‘Green products’ contain little or no phosphate builders  Use sodium zeolites which exchange sodium ions for calcium and magnesium ions  Non-ionic detergents less affected by hard water as they are not charged – require less environmentally damaging builders • Biocidal properties o Cationic detergents have mild biocidal properties as they are attracted to membrane surfaces of bacterial cells where they disrupt cellular processes o The presence of these detergents in wastewater in sewage treatment works can alter the balance of bacterial decomposers  At high concentrations the bacteria that decompose sewage are killed  Impacts on the biodegradability of these detergents, disrupts the decomposition of sewage  As cationic detergents represent only a small % of detergents used the problems associated with their use are minimal as high concentrations of these detergents do not occur often 6. The Solvay process has been in use since the 1860s 6.1 Identify the raw materials used in the Solvay process and name the products • The Solvay process is an industrial method of producing sodium carbonate from salt & limestone 2NaCl (aq) + CaCO3 (s)  Na2CO3 (s) + CaCl2 (aq) • Raw materials: o Purified brine (30% w/w NaCl solution) o Limestone (calcium carbonate) o Ammonia – used in reaction to increase efficiency, reused during the process • Products: o Sodium carbonate o Calcium chloride (waste product) 6.2 Describe the uses of sodium carbonate • Soda ash – anhydrous sodium carbonate o White, odourless, crystalline, readily soluble in water, moderately weak base 32

Use Description Glass • Main use, used to produce glass for windows & bottles manufacture o Acts as a flux in glass production – lowers m.p. of mixture of silicon dioxide & CaCO3 o Heated with sand (silica) & calcium carbonate to over 1500°C o Sodium carbonate & calcium carbonate decompose to their oxides o Oxides combine with silica to form silicates which are found in glass Soap & • Used as base in soap & detergent manufacture in place of stronger alkalis – cheap detergents alternative • Na2CO3 often reacted with CaOH to produce NaOH which is used to produce soap • Also used to make sodium phosphate & sodium silicate builders Softening • CO32- ions precipitate Ca2+ & Mg2+ ions hard water o Contained in some laundry powders Baking soda • Na2CO3 dissolved to make concentrated solution & stream of CO2 passed through it so that production NaHCO3 crystallises out of solution • NaHCO3 (bicarbonate of soda) is also produced as part of the Solvay process • Commonly called baking soda & used in cooking to produce CO2 as a rising agent • NaHCO3 also used in CO2 fire extinguishers & chemical spills Borax • Na2B4O7, used in manufacture of glazes & glass, used as a cleaning agent in laundry products manufacture • Formed when boric acid, H3BO3, neutralises Na2CO3 Volumetric • Pure, anhydrous Na2CO3 used as basic primary standard in volumetric analysis analysis o Can be used in titrations to determine concentration of an acidic solution 6.3 Identify, given a flow chart, the sequence of steps used in the Solvay process and describe the chemistry involved in: – brine purification – hydrogen carbonate formation – formation of sodium carbonate – ammonia recovery • Steps in the Solvay process 1. Brine Purification o Concentrated brine solution (30% w/w) prepared from salt by evaporation of salt water or from underground rock salt deposits o Salt from sea water must have most of impurities removed before used in Solvay process – would interfere with crystallisation of NaHCO3  Ca2+ ions in sea water are removed by precipitation using previously manufactured Na2CO3 Ca2+ (aq) + CO32- (aq)  CaCO3 (s) 2+  Mg ions & other insoluble hydroxides removed by precipitation using NaOH Mg2+ (aq) + 2OH- (aq)  Mg(OH)2 (s)  Precipitates are flocculated & filtered off leaving purified NaCl solution 2. Formation of CO2 & CaO: o Crushed limestone decomposed in a mixed-feed, vertical-shaft kiln o Produces CaO & CO2 gas, CO2 is compressed & cooled CaCO3 (s)  CaO (s) + CO2 (g) 3. Formation of milk of lime: o The CaO is slaked (added to water) to produce a suspension of calcium hydroxide (milk of lime – slightly soluble) – used to recover ammonia CaO (s) + H2O (l)  Ca(OH)2 (s) Ca(OH)2 (s) ↔ Ca2+ (aq) + 2OH- (aq) 4. Formation of ammoniated brine: o Purified brine is saturated with ammonia from Haber process in ammonia absorber tower 33

o Exothermic – considerable heat released so tower must be cooled with cold water pipes to ensure brine is fully ammoniated o Ammoniated brine (7% w/w) is alkaline o Input of NH3 is a once-only process as the ammonia is recycled o Ammoniated brine is partially carbonated – 40% of CO2 is absorbed into ammoniated brine due to high solubility of acidic CO2 in the alkaline solution 5. Hydrogen Carbonate Formation o The partially carbonated ammoniated brine is pumped into carbonators at the top of Solvay towers o CO2 gas is pumped under pressure (about 300 kPa) into base of the Solvay towers (carbonators) o Partially carbonated ammoniated brine trickles downwards over the surfaces of serrated, mushroomshaped plates – provides high surface area for reaction between ammoniated brine & carbonic acid o Carbonic acid forms when the rising CO2 dissolves in the ammoniated brine o Cooling increases solubility of the gas CO2 (g) + H2O (l) ↔ H2CO3 (aq) o Acid-base reaction – HCO3 ions formed NH3 (aq) + H2CO3 (aq)  NH4+ (aq) + HCO3- (aq) 6. Formation of sodium hydrogen carbonate: o Lower temperature (