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Enzyme

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Submitted By pandacjune
Words 2144
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Centre for Foundation Studies, UTAR

Topic Scopes

TOPIC 2
Chemical Bonding







Electrovalent
Covalent bonding
Co-ordinate/dative covalent bonding
Metallic bonding
Intermolecular bonding (including hydrogen bonding, Van der Waals)
• Electronegativity
2

Chemical Bond

Topic Scopes



Bond energies, bond lengths & bond polarities Drawing Lewis structure and calculate the formal charge

• Chemical bond  the force of attraction that binds atoms together in a chemical compound • When atoms react to form chemical bonds, only the electrons in the outermost
/ valence shell are involved

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Forms of Chemical Bonds
1) ________molecular bond – forces hold the atoms within a molecule
• Ionic / Electrovalent Bond
• Covalent Bond
• Metallic Bond
2) ________molecular bond – forces between the molecules
• Hydrogen bonding
• Van der Waals
3) Co-ordinate / Dative Bond
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Electrons In An Atom
2 groups of electrons:
a) Valence electrons
• ____________ shell electrons
• similar chemical properties
• similar group number in periodic table
b) Core electrons
• ________ shell electrons
• not involved in chemical behaviour
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Valence & Core Electrons

Bond & Lone Pairs Electrons
• Valence electrons are distributed as shared (bond pairs) & unshared (lone pairs) B (1s22s22p1)
Core e- = [He]
Valence e- = 2s22p1



Cl




H
Br ([Ar]3d104s24p5)

lone pair e-

Shared or bond pair e-

Core e- = [Ar]3d10
Valence e- = 4s24p5
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Lewis Symbol &
Lewis Structures

This is called a LEWIS ELECTRON DOT structure (Lewis Structure)

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Building A Lewis Electron
Dot Structure

• Lewis symbol — a chemical symbol to represent the nucleus & core electrons of an atom, together with dots placed around the symbol to represent the valence electron • For example: Si

Electron configuration:
 Si 
[Ne]3s2 3p2

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3) Form a single bond between the central atom and each surrounding atom

Example 1: Ammonia, NH3
1) Decide on the central atom;
 Never H
 Central atom is atom of lowest affinity for electrons.
Therefore, N is central.
2) Count valence electrons
H = 1, N = 5
Total = (3 x 1) + 5 = 8 electrons / 4 pairs 10

Example 2: Carbon Dioxide, CO2
1) Central atom = _______
2) Valence electrons = __ e- or __ pairs e3) Form bonds
6 pairs of electrons are now left

4) Remaining electrons form LONE PAIRS to complete octet (8 e-) as needed


4) Place lone pairs on outer atoms

3 BOND PAIRS and 1 LONE PAIR
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Example 3: Sulfur Dioxide, SO2
1) Central atom = S
2) Valence electrons = ___ e- or __ pairs e3) Form bonds

5) So that C has an octet, we shall form
DOUBLE BONDS between C & O

Each atom is surrounded by an octet of electrons



O


S



Leave 14 e4) Remaining pairs become lone pairs, first on outside atoms and then on central atom 



O


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5) Form double bond so that has an octet — but note that there are 2 ways of doing this bring in left pair




O


S







or bring in right pair

O


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Resonance Structures
• 2 or more Lewis structures can be drawn for a molecule or polyatomic ion
• For example:
a) Sulfur dioxide, SO2

This leads to the following structures


O




S





O






S





O


O


b) Nitrogen dioxide, NO2

These equivalent structures are called
RESONANCE STRUCTURES
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Octet Rule

Exceptions To Octet Rule

• Octet rule — the tendency of molecules & polyatomic ions to gain, lose or share electrons until they are surrounded by 8 valence electrons
• Each atom has a share in 4 pairs of electrons, so each has achieved a stable noble gas configuration

Exceptions when molecules & ions have:
a) fewer than 4 pairs of electrons (8 valence electrons) on a central atom (BF3)
b) more than 4 pairs of electrons (8 valence electrons) on a central atom (PCl5, SF6)
c) have an odd number of electrons (NO,
NO2)

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Formal Atom Charges











nonmetal e- transfer from atom reducing agent to oxidizing agent



O

Formal Charge (C) = 4 – 1/2(8) – 0 = 0
 Sum of formal charge = 0 + 0 = 0


Ionic
compound

b) Cl ([Ne]3s23p5)
Cl- ([Ne]3s23p6)
 [Ar] noble gas configuration

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• The mutual attraction of 2 _________ for the same electrons
• Sharing of valence electrons between atoms +

HB

HA

1) When 2 atoms approach each other, the electrons repel one another due to like charges Electron-electron repulsion

HB

• A region of high electron density resulting from the overlapping atomic orbitals between bonded atoms

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Formation of Covalent Bond
Example: H• •H

Covalent Bonding

HA

20

• For the formation of NaCl:
a) Na (1s22s22p63s1)
Na+ (1s22s22p6)
 [Ne] noble gas configuration

Na+ Cl


Na   Cl





metal atom 

+  Cl

C

Valence Electron Configurations
& Ionic Compound Formation

• Transferring of electrons from 1 atom to another, creating positive & negative ions
• Electrostatic attraction between 2
________________ charged ions
• For example:

Na 



O



Ionic Bonding

Formal Charge (O) = 6 – 1/2(4) – 4 = 0


• The formal charge for an atom in a molecule / ion is the charge calculated for that atom based on the Lewis structure of the molecule / ion
• Formal charge
= Group number – 1/2 (no. of bonding electrons) – (no. of lone pair electrons)
= No. valence electrons in free atom –
1/2 (no. of bonding electrons) – (no. of lone pair electrons)
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Example:
Carbon Dioxide, CO2

2) Similarly, their nuclei also repel each other Nucleus-nucleus repulsion
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Polar Covalent Bond in Polar
Molecule

4) The nuclei & the electrons of neighboring atoms, however, attract each other
Electron-nucleus attraction

H

F

 Non-uniform distribution of electron

5) When the attraction forces > powerful than repulsive forces, a bond is formed

density between the two atoms

 Due to F is more electronegative than H, the F atom tends to exert a stronger attraction on the bonding e– compared to the H atom

Attraction is greater than repulsion
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Coordinate Covalent Bonding
 Electron density of the bonding electrons to be higher around the F atom than the H atom  ‘F end’ of molecule acquires a partial negative charge while the ‘H end’ acquires a partial positive charge

 The covalent bond in the H–F molecule is polarised • Dative bond
• In some molecules or ions, a single atom contributes both of the electrons (lone pair e-) to be a shared pair of electrons
• Electron-pair donor ― the atom which donates the lone pair electrons
• Electron-pair acceptor ― the atom accepting the pair of electrons

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Metallic Bonding

Dative bond
A dative bond is represented by an arrow pointing from the donor of the electron pair
(N atom) to the acceptor of the electron pair
(H atom)
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FHSC1114 Physical Chemistry

• Metallic bonding
– bonding in solid metals
– Electrostatic attraction between the positively charged metal ions & the ‘sea’
/‘cloud’ of delocalised (mobile) electrons
• For example: Cu, Fe, Al….
• Electron Sea Model – describe the metallic bonding
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Centre for Foundation Studies, UTAR

Metallic Bonding
Electron Sea Model
+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

+

 e-

+

Positive metal ion

Delocalised valence electrons

• In metals, the valence electrons are associated with a particular metal atom but are free to move (mobile) throughout the solid piece of metal
• Electrons are free to move away from a – ve electrode to a +ve electrode
(delocalised) when an electrical potential is applied
• With valence electrons now delocalised, the metal atoms are effectively ionised

31

Effect of Bonding on Physical Properties

32

2) Ionic bonding
• poor conductor in solid state (electrical conductivity only in aqueous solution)
• no thermal conductivity
• no ductility and malleability
• most are colourless
• soluble in polar solvents but insoluble in nonpolar solvents
• high melting points
• solid at room temperature

1) Metallic Bonding

 high electrical conductivity
 high thermal conductivity
 ductility and malleability
 grey, black, brown/yellow in appearance
 not soluble
 melting point (depend on bond strength)
 solid at room temperature & non volatile
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Intermolecular
Forces of Attraction

3) Covalent bond
• no electrical and thermal conductivity
• no ductility and malleability
• water – colourless
• halogen – colour
• nonpolar molecules – insoluble in water but soluble in organic solvents
• polar molecule – soluble in water
• giant structure – insoluble in all solvents

• In addition to the covalent bonds within a molecule, there are also other forces of attraction between molecules
• Intermolecular forces of attraction weaker than covalent / ionic bonds
• 2 types of intermolecular forces:
a) Van der Waals forces
b) Hydrogen bond
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Van der Waals Forces

3 Types of Dipoles

• Van der Waals forces are weak intermolecular forces that contributed by permanent dipole – permanent dipole & temporary dipole – induced dipole attraction • The larger the molecule size, or the larger the number of electrons in a molecule, the larger the Van der Waals forces, & the higher the melting point/boiling point of the molecule

1) Permanent dipole-permanent dipole attractions  The force of attraction between the
_________ end of a polar molecule & the
_________ end of another polar molecule dipole – dipole attraction

37

2) Dipole-induced dipole attractions

• Interaction between a polar molecule & nonpolar molecule
+

neutral

-



The nonpolar is polarized by the polar molecule polarized
+

+

-

-

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3) Induced dipole-induced dipole attractions
(London Dispersion Forces)
• Interaction between non–polar molecules
(O2, N2, CO2 & noble gases)
• This is caused by the random movement of the _________ in an atom or molecule
• For example:
Argon has 18 e- in its atom. Since it is non-polar, the arrangement of e- is symmetrical on the average

39

• The positive end of the temporary dipole will distort the e- cloud of the neighboring atoms giving rise to induced dipoles

• When these e- revolve around the nucleus, the e- density might be higher at one end than the other & cause a temporary dipole

temporary dipole then attract one another

Electron cloud
+

-

-

+

-

+

-

+

Temporary dipole-induced dipole attractions

temporary dipole
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Hydrogen Bonding in
Hydrogen Fluoride (HF)

Hydrogen Bond
• A special type of permanent dipole – permanent dipole attraction between a H atom (which is bonded to a small & highly electronegative atom - O, N or F) and a lone pair of electrons of another very electronegative atom.

–
+

+

–

–

+

 If 2 molecules of HF are close to another, the H atom of 1 molecule will be attracted to the F atom of another molecule

 Electrostatic attraction between the partial
43

+ve charge on the H atom & the partial – ve charge on F atom
44

Hydrogen Bonding in
Ammonia (NH3)
+

+






–


+

+
–


+

Bond Properties

+

• In NH3 molecule, the N atom has 1 lone pair of electrons
• Each NH3 molecule can form 1 hydrogen bond Bond Polarity
Bond Order
Bond Length
Bond Energy

45

46

• Polar bond in polar molecules – the bond between the two atoms has a partial +ve end & a partial -ve end (dipole moment)
• The unequal sharing of electrons leads to:
a) a partial _________charge on the more electronegative element ( )
b) a partial _________charge on the less electronegative element ( +)
• A covalent molecule is polar if the covalent bond is polarised

• Electronegativity, 
a) a measure of the ability of an atom in a molecule to attract bonding electrons in a covalent bond to itself
b) decide whether a bond is polar, which atom of the bond is -ve & which is +ve
c) increase from left to right across a period & decrease down a group
 Electronegativity trend: F > O > N > C > H
F > Cl > Br > I

47

48

Bond Polarity

FHSC1114 Physical Chemistry

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Bond Order

Bond Order

• Bond order
= number of bonding electron pairs shared by 2 atoms in a molecule

Bond

Single

1

CC

Double

2

CC

Number of shared pairs linking X – Y
Number of X-Y links

Bond order

CC

• Bond order (BO) =

Type of bonding Triple

3

49

50

Bond Length

Bond Order
• Bond order is depends on to 2 important bond properties:
a) Bond length
– the distance between the centers of 2 atoms joined by a bond
b) Bond energy
– the energy required to break a bond

• Bond length depends on size of bonded atoms • Size of bonded atoms , bond length ___


H—F


H—Cl

H—I

51

Bond Length
• Bond length depends on bond order





• The bond lengths become shorter as the bond orders increase
• Bond length:
Single bond  double bond  triple bond 53

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Bond Energy
Type of bond

Typical bond energy
(kJ mol-1)
Covalent bonds
200 – 400
Hydrogen Bonds
20 – 80
Van der Waals bonds
Less than 20
The relative strengths of chemical bonds: covalent bonds > hydrogen bonds > Van der
Waals bonds
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Relation of Bond Length & Bond
Strength with Bond Order
• The GREATER the number of bonds
(bond order), the HIGHER the bond strength (energy required)

Learning Outcomes
• Explain principles that govern properties of chemical systems.
• Discuss properties of gases, chemical bonds and chemical kinetics.

• The GREATER the number of bonds
(bond order), the SHORTER the length of bond 55

FHSC1114 Physical Chemistry

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...Enzymes An enzyme is a protein molecule that helps to increase the rate in which chemical reactions occur. Recognizing an enzyme is very simple as they are named after a substrate with the suffix “ase” attached to the end. In terms of activation energy, an enzyme’s function is to decrease the amount of energy necessary for the reaction to take place. An enzyme inhibitor prevents that function from happening by binding to the enzyme. There are two different ways that enzyme inhibition can occur. The substrate imposter can either bind directly to the active site so the real substrate cannot bind or it can bind to a different site other than the active site. When it binds to another site it causes the enzyme to change its shape, therefore not allowing the real substrate to be accepted into the active site. Two different theories exist that explain how substrates fit into the active site of an enzyme. First is the Lock and Key theory which states that a substrate (key) fits into the active site of an enzyme (lock) perfectly without any change of shape of the active site at all, according to Elmhurst College’s virtual chemistry book. This theory differs from the Induced Fit theory, which states that when a substrate binds to the active site, the active site’s shape slightly changes to perfectly fit the substrate. Just like a protein, the class of macromolecule that an enzyme belongs to, the function of an enzyme depends greatly on its tertiary structure. If the structure of an...

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Enzymes

...Title: Enzyme Introduction The main reason for conducting this experiment is to establish the various factors that affect enzymes and reaction rates. Various experiments have been conducted to help gain a wide range of the factors that affect enzyme controlled reactions. Enzymes are affected by very many factors. It was the main aim of this experiment to establish these factors and the manner in which they affect them. This experiment also seeks to establish the manner in which some enzymes like Catalase affect the rates of reactions (Cohnheim 2009). Methods To establish the factors that affect enzymes, the procedures for the experiments to be carried out had to be almost perfect. For this reason the apparatus to be used had to be cleaned thoroughly just before commencing the experiment. To avoid differentiated results, similar kinds of apparatus were used all through the experiment. In this case glass test tubes were used. Also measuring apparatuses used were of the same size and volume. In this case four experiments were carried out. The first experiment is to establish the manner in which the enzyme Catalase affects reaction rates. The procedure of this experiment is as follows; using a pencil, label tree test tubes as test tube 1, 2 & 3. On these test tubes, label two marks using the pencil. These are at the 1cm mark and at the 5 cm mark. For the first test tube, pour in Catalase enzyme up to the first mark and add Hydrogen Peroxide up to the...

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Enzymes

...Enzymes and pH pH is a measure of H+ concentration. The higher the concentration of H+ the lower the pH values (acids) A hydrogen ion has a (+) charge so will be attracted to negatively charged molecules or parts of molecules. As like charges repel, positive molecules or parts of molecules will repel hydrogen ions. Large numbers of hydrogen bonds and ionic bonds are responsible for holding the tertiary structure of an enzyme protein in place. This ensures that the active site is also held in the right place. These bonds are due to the attraction between oppositely charged groups on the amino acids that make up the enzyme protein. Because of their charge, hydrogen ions can interfere with the hydrogen and ionic bonds in the molecule holding the tertiary structure in place. This means increasing or decreasing the concentration of hydrogen ions can alter the shape of the tertiary structure and therefore the shape of the active site. This can also aler the rate of an enzyme-controlled reaction. The induced-fit hypothesis suggests that an important part of catalysis in the active site relies on charged groups on the R-groups of the amino acids that make up the active site. Increasing the concentration of hydrogen bonds will alter the charges around the active site, as more hydrogen ions are attracted towards any negatively charged groups in the active site. Optimum pH At the optimum pH, the concentration of hydrogen ions in the solution gives the tertiary structure...

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Biology- Enzymes

...LAB REPORT: ENZYMES Part I: Graphs and Data TIME COURSE: ABSORBANCE VS. TIME Provided data: Time(minutes) | Experimental ABS @ 405nm | Control ABS @ 405 nm | Exp. ABS minus Control ABS | Micromoles p-Nitrophenol | 0 | 0.057 | 0.051 | 0.006 | 0.0004 | 10 | 0.207 | 0.053 | 0.154 | 0.0064 | 20 | 0.351 | 0.054 | 0.297 | 0.0120 | 30 | 0.501 | 0.055 | 0.446 | 0.0181 | 60 | 0.955 | 0.064 | 0.891 | 0.0362 | Personal data: Time(minutes) | Experimental ABS @ 405nm | Control ABS @ 405 nm | Exp. ABS minus Control ABS | Micromoles p-Nitrophenol | 0 | 0.092 | 0.064 | 0.028 | 0.0010 | 10 | 0.262 | 0.048 | 0.214 | 0.0085 | 20 | 0.429 | 0.054 | 0.375 | 0.0140 | 30 | 0.599 | 0.051 | 0.548 | 0.0208 | 60 | 0.976 | 0.050 | 0.926 | 0.0350 | STANDARD CURVE: Provided data: Micromoles p-Nitrophenol | Absorbance @ 405 nm | 0.0000 | 0.000 | 0.0025 | 0.058 | 0.0050 | 0.118 | 0.0100 | 0.245 | 0.0200 | 0.496 | 0.0400 | 1.000 | Personal data: Micromoles p-Nitrophenol | Absorbance @ 405 nm | 0.0000 | 0.000 | 0.0025 | 0.071 | 0.0050 | 0.167 | 0.0100 | 0.228 | 0.0200 | 0.519 | 0.0400 | 1.050 | TIME COURSE: PRODUCT VS. TIME Provided data: Time | Micromoles product | 0 | 0.0004 | 10 | 0.0064 | 20 | 0.0120 | 30 | 0.0181 | 60 | 0.0362 | Personal data: Time | Micromoles product | 0 | 0.0010 | 10 | 0.0085 | 20 | 0.0140 | 30 | 0.0208 | 60 | 0.0350 | TEMPERATURE: PRODUCT VS. TEMPERATURE Provided data: ...

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