...Effect of Table Salt on Water Reaching 99⁰ C Abstract I had always been told that, when cooking, adding salt to water would make the water boil faster. Recently I had heard that this wasn’t true. Beginning with the hypothesis that salt did increase the boiling time, I set out to discover if this was the case. Using the home stovetop, I boiled 2 liters of water each time and the amount of salt added was the control. The experiment was to find out how long it took the water to reach 99⁰ Celsius. After 8 boils at four different salt concentrations, the effect on the time to reach the desired temperature was negligible. The water was already boiling each time the temperature achieved the target. The experiment seemed to indicate the salt has no effect on the time it takes water to reach 99⁰ C. Background As mentioned in the abstract, I had been led to believe that adding even a “pinch” of salt to water would accelerate the time that water would reach boiling. However, this was challenged recently and I was curious if I had been told a “wives tale” all my life. I have no chemistry background, but since the experiment, I have a better understanding of what should have occurred because of the added solvent (NaCl in this case), which will be addressed in the conclusion. Design Eight experiments were run using four measurements of table salt in 2 liters of tap water. Care was taken to ensure each of the variables, other than the control variable of salt, remained constant...
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...Chemistry Revision Hazard Symbols States of Matter As heat is added to a solid the particles start to vibrate more and more vigorously. Eventually when it reaches its melting point the particles have enough energy to break their bonds and melt into a liquid. As it is cooled energy is taken away so the particles vibrate less and if a liquid or gas the bonds become stronger and so it freezes or condenses. In the case of a solid it becomes less flexible. Particles in a solid vibrate around their equilibrium but don’t move and keep a rigid shape with their bonds intact. Liquid particles are similar but have more energy so vibrate faster and have more fluidity. Gas particles have no bonds and move around very quickly Structure of an atom |Particle |Where? |Mass |Charge | |Proton |Nucleus |1 |1+ | |Neutron |Nucleus |1 |0 | |Electron |Energy Levels |1/1840 |1- | Atomic/Proton Number – Number of protons (small number) Mass Number – Sum of protons and neutrons Mass Num – Atomic Num = Number of neutrons Number of protons = number of...
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...properties reflect the solute’s properties. The physical properties that the solution and solute do not share are known as colligative properties and they depend solely on the solute concentration. Some of these properties include vapor pressure lowering, boiling-point elevation, freezing point lowering, and osmotic pressure. The solvent boils when the vapor pressure, or tendency of solvent molecules to escape, is equivalent to the atmospheric pressure. At this moment, the gaseous and liquid states of the solvent are in dynamic equilibrium and the molecules change from the liquid to the gaseous states and from the gaseous to liquid states at equal rates. The dissolution of a solute with very low vapor pressure, or a nonvolatile solute, raises the boiling point and lowers the freezing point. Similarly, anti-freeze lowers the freezing point and lowers the boiling point. The colligative-property law describes these effects, stating that the "freezing point and boiling point of a solution differ from those of the pure solvent by amounts that are directly proportional to the molar concentration of the solute" (Brown, 203-204). The colligative-property law can be expressed using the equation: D T = Km, where D T is the change in freezing or boiling point, K is a solvent-specific constant, and m is the solution’s molality. Adding a solute to a...
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...Hydrostatic vs Osmotic Pressure Hydrostatic pressure is the pressure going outward from the capillary. Hydrostatic Pressure forces the fluid from capillary to move outward into the interstitial space. Osmotic pressure is the water trying to move from interstitial space into the capillary Hydrostatic pressure will be greater in the arterial side as opposed to the venule side. Osmotic Pressure is constant throughout the capillary. Because the hydrostatic pressure drops across the capillary, at the artery side fluid is pushed into the interstitial space whereas in the venule side, fluid is pushed into the vessel. If you have high blood pressure, it can cause build up of fluid in interstitial space (due to high hydrostatic pressure) and cause edema. Also, if fluid is not taken in by lymphatic vessels, it can also cause edema. Water that is lost from fluid due to hydrostatic pressure eventually goes into lymph vessels and is put back into the vessels. Note: Capillary walls are made up of endothelial cells Sodium Potassium ATPase: Pumps 3 Na ions out and 2 K ions in. Ketone Bodies: Ketone bodies are produced when AcetylcoA exceeds krebs cycles capacity. So, when you are starving, ketone bodies are used primarily rather than glucose. Glucose is preserved for brain. Brain, heart, muscle can use ketone bodies. Liver cannot use ketone bodies. Insulin: Insulin helps glucose intake by cells normally. Unsaturated fat is easy to burn off because they produce...
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...C2 Revision list Topic 1 Atomic structure and the periodic table ● Explain how Mendeleev arranged the elements, known at that time, in a periodic table by using properties of these elements and their compounds and used his table to predict the existence and properties of some elements not then discovered ● Classify elements as metals or non-metals according to their position in the periodic table ● Describe the structure of an atom as a nucleus containing protons and neutrons, surrounded by electrons in shells (energy levels) ● Demonstrate an understanding that the nucleus of an atom is very small compared to the overall size of the atom ● Describe atoms of a given element as having the same number of protons in the nucleus and that this number is unique to that element ● Recall the relative charge and relative mass of a proton, a neutron and an electron ● Demonstrate an understanding that atoms contain equal numbers of protons and electrons ● Explain the meaning of the terms, atomic number, mass number and relative atomic mass ● Describe the arrangement of elements in the periodic ● Demonstrate an understanding that the existence of isotopes results in some relative atomic masses not being whole numbers ● Calculate the relative atomic mass of an element from therelative masses and abundances of its isotopes ● Draw the electronic configurations of the first 20 elements in the periodic table as diagrams and in the form 2.8.1 ● Describe the connection between the number...
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...with dilute hydrochloric acid is Na2CO3(aq) + 2HCl(aq) → 2NaCl(aq) + CO2(g) + H2O(l) The end-point is marked by using methyl orange as indicator. Introduction : Chemicals : Apparatus : Procedure : solid sodium carbonate, 0.1 M hydrochloric acid 1. 2. 3. 4. 5. 6. 7. 8. Weight out about 1.3 g of anhydrous sodium carbonate accurately using the method of “weighing by difference”. Transfer the weighed carbonate to a beaker and add about 100 cm3 of distilled water to dissolve it completely. After dissolving, transfer the solution to a 250.00 cm3 volumetric flask. Rinse the beaker thoroughly and transfer all the washes into the volumetric flask. Remember not to overshoot the graduation mark of the flask. Make up the solution to the mark on the neck by adding water. Pipette 25.00 cm3 of sodium carbonate solution to a clean conical flask. Add 2 drops of methyl orange indicator to the carbonate solution. Titrate the carbonate solution with the given dilute hydrochloric acid until the colour of solution just changes from yellow to orange. Repeat the titration two times. Calculation : Results : Questions : Calculate the molarity of the sodium carbonate solution prepared and the molarity of the hydrochloric acid. 1 2. What is the meaning of “weighing by difference”? Suggest one method other than using acid-base indicator to...
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...non-metals. Equipment: Hot plates, glass plates, beakers, glass rods, scale to weigh out the salt, graduated flask for water, furnace, crucibles, sand, tongs, safety glasses and gloves. Starting materials: Epsom salts, water, 50% Aluminum - 50% Copper alloy (previously alloyed). Safety: Whenever you are dealing with hot liquids, there is the potential for burns and spills. Protect yourself from the possible risks, especially around the hot metal. Make sure you know where any potential spillage will go and place something in the way to protect yourself. In part ‘B’ be especially careful of the hot metal. Remember, it will still be very hot, even when it has changed back into a solid. Procedure: Part A: A supersaturated solution of salt will separate out into crystals on cooling. The size of the crystals is a function of the rate of cooling, the amount of impurities present and the degree of supersaturation (concentration of salt present in the solution). 1. Dissolve 25 grams of Epsom salts in 25 ml of water. Heat the water until all the salt dissolves, but keep the water below the boiling point. If all the salt will not dissolve, add water in small quantities until it does. Pour some of the solution onto a clean glass plate so as to form a thin film and watch it solidify. If it solidifies too rapidly, return it to your beaker and add more salt. Again, keep track of how much additional salt is added. When you achieve the proper mixture you should be able to observe the crystals...
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...CARIBBEAN EXAMINATIONS COUNCIL Caribbean Secondary Education Certificate CSEC® CHEMISTRY SYLLABUS Effective for examinations from May–June 2015 CXC 21/G/SYLL 13 Published by the Caribbean Examinations Council. All rights reserved. No part of this publication may be reproduced, stored in a retrieval system, or transmitted in any form, or by any means electronic, photocopying, recording or otherwise without prior permission of the author or publisher. Correspondence related to the syllabus should be addressed to: The Pro-Registrar Caribbean Examinations Council Caenwood Centre 37 Arnold Road, Kingston 5, Jamaica Telephone Number: + 1 (876) 630-5200 Facsimile Number: + 1 (876) 967-4972 E-mail Address: cxcwzo@cxc.org Website: www.cxc.org Copyright © 2013 by Caribbean Examinations Council The Garrison, St Michael BB14038, Barbados CXC 21/G/SYLL 13 Contents RATIONALE ................................................................................................................................... AIMS ............................................................................................................................................. CANDIDATE POPULATION ............................................................................................................. SUGGESTED TIME-TABLE ALLOCATION ........................................................................................ ORGANISATION OF THE SYLLABUS .................................................
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...PRASAD UNIT1 WATER CYCLE Evaporation is the process by which water changes from a liquid to a gas or vapor. Evaporation is the primary pathway that water moves from the liquid state back into the water cycle as atmospheric water vapor. Studies have shown that the oceans, seas, lakes, and rivers provide nearly 90 percent of the moisture in the atmosphere via evaporation, with the remaining 10 percent being contributed by plant transpiration. A very small amount of water vapor enters the atmosphere through sublimation, the process by which water changes from a solid (ice or snow) to a gas, bypassing the liquid phase. This often happens in the Rocky Mountains as dry and warm Chinook winds blow in from the Pacific in late winter and early spring. When a Chinook takes effect local temperatures rise dramatically in a matter of hours. When the dry air hits the snow, it changes the snow directly into water vapor, bypassing the liquid phase. Sublimation is a common way for snow to disappear quickly in arid climates. (Source: Mount Washington Observatory) Why evaporation occurs Heat (energy) is necessary for evaporation to occur. Energy is used to break the bonds that hold water molecules together, which is why water easily evaporates at the boiling point (212° F, 100° C) but evaporates much more slowly at the freezing point. Net evaporation occurs when the rate of evaporation exceeds the rate of condensation. A state of saturation exists when these two process rates are equal, at...
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...Written Assignment 1: Intermolecular Forces and Liquids and Solids Answer all assigned questions and problems, and show all work. 1. Explain and give an example for each type of intermolecular force. A: a. Dipole-dipole interaction: a dipole-dipole interaction is the electrostatic attraction between the positive end of one polar molecule and the negative end of the other. Dipole-dipole attraction occurs between molecules which are permanent dipoles (polar covalent molecules). An example of a dipole-dipole interaction is HCl and HCl. b. Dipole-induced dipole interaction: a dipole-induced dipole interaction is produced in neutral molecules when they are introduced into a magnetic field (i.e induced by an electric current or by a permanent dipole). Subjecting a neutral molecule to such magnetic fields has effects on the charge of the molecule. The negative charges concentrate in a specific point totally opposite from the positive charges. An example of dipole-induced dipole interaction is HCl and H2 c. Ion-dipole interaction: an ion-dipole interaction is the force between an ion and a neutral polar molecule which possess a dipole moment. Polar molecules are dipoles; they have a positive end and a negative end. The positive ions are attracted to the negative end of a dipole, while negative ions are attracted to the positive end. An example of ion-dipole interaction is K+ ---H2O d. Dispersion forces (London forces): London forces are weak intermolecular...
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...ions, or between molecules and ions. Table 13.1 Summary of Intermolecular Forces Ions Dipoles Induced Dipoles (Overhead & book p 585) Covalent bond energies 100-400 kJ/mol Attractive forces between ions 700-100 kJ/mol Intermolecular attractions less than 15% of bond energies Intermolecular Forces Ion-Ion Forces Na+ — Cl- in salt. These are the strongest forces. Lead to solids with high melting temperatures. NaCl, mp = 800 oC MgO, mp = 2800 oC Intermolecular Attractions Coulomb’s Law Force ~ (n+)(n-)/d2 Distance - twice the distance = 1/4 the force Charge on the Ion Magnitude of the dipole Composition - Solids and Liquids are closer so composition has greater role in attractive forces Attraction Between Ions and Permanent Dipoles Water is highly polar and can interact with positive ions to give hydrated ions in water. Attraction Between Ions and Permanent Dipoles Water is highly polar and can interact with positive ions to give hydrated ions in water. Dissolving Ionic Solids Attraction Between Ions and Permanent Dipoles Many metal ions are hydrated. It is the reason metal salts dissolve in water. Attraction Between Ions and Permanent Dipoles Attraction between ions and dipole depends on ion charge and ion-dipole distance. Measured by DHhydration for Mn+ + H2O --> [M(H2O)x]n+ Solvation (aka hydration) Attraction Between Ions and Permanent Dipoles Attraction between ions and dipole depends on ion charge and ion-dipole distance...
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...2 tsp (10 mL) baking soda * 1/2 tsp (2 mL) salt * 2 cups (500 mL) cold coffee or water * 1 cup (250 mL) vegetable oil * 2 tsp (10 mL) vanilla * 3 tbsp (45 mL) cider vinegar Preparation In large bowl, whisk together flour, sugar, cocoa powder, baking soda and salt. Whisk in coffee, oil and vanilla. Stir in vinegar. Prepare pans as per below. Pour in batter, smoothing top. Bake in 350°F (180°C) oven until cake tester inserted into centre comes out clean (see below for times). Let cool for 10 minutes. Invert onto rack and remove pan. Remove paper; let cool completely. Metal Pan Size: One 13- x 9-inch (3.5 L) cake pan; Grease and line with parchment paper; Bake 30 to 35 minutes. Metal Pan Size: Two 8-inch (1.2 L) or 9-inch (1.5 L) round cake pans; Grease and line with parchment paper; Bake 20 to 25 minutes. Metal Pan Size: Two 8-inch (2 L) or 9-inch (2.5 L) square cake pans; Grease and line with parchment paper; Bake 20 to 25 minutes. Metal Pan Size: 24 muffin cups for cupcakes; Grease or line with paper cups; Bake 18 to 20 minutes. 2. FABULOUS FUDGE CHOCOLATE CAKE * 2 1/4 cups all-purpose flour * 2 teaspoons baking soda * ½ teaspoon salt * 1/2 cup butter * 2 1/2 cups packed brown sugar * 3 eggs * 1 1/2 teaspoons vanilla extract * 3 (1 ounce) squares unsweetened chocolate, melted * 1 cup sour cream * 1 cup boiling water * 1/2 cup butter * 1 cup packed...
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...O2) or different (e.g. H2O). The chemical formula shows the number and type of atoms present. Non-metal compounds are made of molecules: Carbon dioxide contains CO2 molecules Methane (natural gas) contains CH4 molecules AN ION is an atom or group of atoms with an electrical charge (+ or -). Metal compounds such as sodium chloride or copper sulphate contain ions. Sodium chloride is made of Na+ and Cl- ions Copper Sulphate is made of Cu2+ and SO42- ions Note that metals form positive ions while non-metals form negative ions. A solid is represented by (s). e.g. H2O(s) is ice. A liquid is represented by (l) e.g. Fe(l) is molten iron. A gas is represented by (g) e.g. H2O(g) is steam. A solution in water is represented by (aq). Salt dissolved in water is NaCl(aq). You should remember that the common gases are diatomic (have 2 atoms in each molecule). These are Oxygen O2; Hydrogen H2; Nitrogen N2; and Chlorine Cl2. Elementary Particles Atoms are made up of smaller particles called protons, neutrons and electrons. The protons and neutrons cluster together in a small nucleus at the centre of the atom while the electrons orbit the nucleus. The main properties of the particles are: |Particle |Mass |Charge | |PROTON |1 |+1 | ...
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...evidence) 5.1 Unit 2: Phases, phase changes, and the effect of heat on matter Activity 2a: The properties of solids, liquids and gases 2.3, 2.4, 2.7 Activity 2b: Understanding atmospheric pressure 16.5 Activity 2c: Heat and the motion of sub-microscopic particles 2.5, 2.6 Activity 2d: Absolute zero and the Kelvin temperature scale 2.6 Activity 2e: Exploring the phase changes 2.7, 8.2-8.6 Activity 2f: The difference between boiling and evaporation 8.4 Unit 3: An overview of the periodic table Activity 3: Properties of the elements and the periodic table 3.2, 3.3, 2.2 Unit 4: The structure of the atom Activity 4a: The nucleus, isotopes, and atomic mass 4.2, 4.3, 4.4 Activity 4b: The electrons and the shell model 4.2, 4.5-4.8 Unit 5: An introduction to ionic, covalent, and metallic bonding Activity 5: Conductivity and models of chemical bonding 3.4, 6.2-6.5 Unit 6: Exploring covalent compounds (molecules) Activity 6a: Molecules and lewis dot structures 6.5 Activity 6b: VSEPR theory and molecular shape 6.6 Activity 6c: Polarity, intermolecular forces and boiling point 6.7-6.8, 7.1, 7.5 Activity 6d: The amazing properties of water 8.1-8.4 Activity 6e: Symbolic representations of molecules 3.4, 12.1-12.4 Activity 6f: An introduction to polymers 12.4 Unit 7: Exploring ionic compounds (salts)...
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... | |Chlorine |Green Gas |Green | |Bromine |Red/Brown Liquid |Orange | |Iodine |Dark Grey/Violet Solid |Purple/Violet | |Astatine |Black Solid |Dark Purple | The halogens become darker as you go down the group. Fluorine is very pale yellow, chlorine is yellow-green and bromine is red-brown. Iodine crystals are shiny purple-black but easily turn into a dark purple vapour when they are warmed up. Common properties The halogens have the following properties in common: • they are non-metals • they have low melting and boiling points • they are brittle when solid • they are poor conductors of heat and electricity • they have coloured vapours • their molecules each contain two atoms (they are diatomic) Similarities |Halogen |Atomic Number |Electronic Configuration |Formula | |Fluorine |9 |2, 7...
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