Energy

1. Energy • Energy is defined as having the ability to do work • Energy allows objects to move and to change • Walking, lifting, chemical reactions, etc. involve work • Two kinds of energy: - Kinetic = energy of motion (e.g. climbing ladder) - Potential = stored energy (e.g. object at top of ladder) • Potential and kinetic energy can be interconverted • Kinetic and potential energy come in many forms (heat, light, electrical, mechanical, chemical, rotational) • Energy produced by chemical reactions can be used to do work in biological systems (ATP produced by oxidation of glucose powers many cellular processes)

2. Measuring Heat • Heat is the amount of thermal energy transferred between two objects at different temperatures (Not the same as temperature, a measure of molecular kinetic energy that predicts direction of heat flow) • Heat is usually measured in units of calories (cal) or joules (J); kcal or kJ are used for larger amounts of heat • Specific heat = amount of heat to raise the temperature of 1 gram of a substance by 1ºC • Water has the highest specific heat of any substance • Water keeps the temperature stable around oceans and large lakes and also in the body • Metals have low specific heats, so they heat up quickly

3. Calculations Using Specific Heat • Specific heat is used for temperature changes • Heat (gained or lost) = mass x T x Sp. Heat • Example 1: How much heat is absorbed (in cal) when 25 g of water is heated from 0.0ºC to 100.0 ºC (given that specific heat of water is 1.00 cal/g ºC )? 25 g x 100.0 ºC x 1.00 cal/g ºC = 2.5 x 103 cal • Example 2: How much heat is released (in kcal) when 100.0 g of water cools from 22ºC to 0.0ºC ? 100.0 g x 22ºC x 1.00 cal/g ºC x 1 kcal/1000 cal = 2.2 kcal

4. Attractive Forces between Molecules • Molecules are held together in liquids and solids by intermolecular forces • Forces are due to attraction of opposite charges

6. States of Matter • Recall: matter = mass + volume (occupies space) • Matter exists in 3 physical states: solid, liquid and gas • Solids: definite shape and volume, strong intermolecular forces (ionic, H-bonding) • Liquids: definite volume, take shape of container, moderate intermolecular forces (H-bond, dipole-dipole, dispersion) • Gases: takes shape and volume of container, no intermolecular forces (particles are too far apart) • Physical state is temperature (and pressure)-dependent • At lower T compounds have lower KE, so even compounds with weak intermolecular forces can form solids at very low temperatures

8. Melting and Freezing • When matter is converted from one physical state to another it’s called a “change of state” • Solid goes to liquid = melting - Heat increases movement of particles in solid - At melting point E is high enough to overcome strong intermolecular attractive forces - This E is called the “heat of fusion” - Solid absorbs heat until all is melted, then can rise in T • Liquid goes to solid = freezing - Freezing point = melting point - At melting/freezing point both states coexist at equilibrium (melting rate = freezing rate)

9. Calculations Using Heat of Fusion • Use heat of fusion to calculate heat required to melt or heat removed to freeze (80. cal/g for H2O) • Heat = mass x heat of fusion • Example: If 12.0 g of water at 0.0ºC is placed in the freezer, how much heat (in kJ) must be removed from the water to form ice cubes? Heat = 12.0 g x (80. cal/g) = 960 cal 960 cal x (4.18 J/cal) x (1 kJ/1000 J) = 4.0 kJ

10. Boiling and Condensation • Liquid goes to gas = evaporation - Happens when enough heat is added to overcome attractive forces (heat increases KE of liquid particles) - This E is called “heat of vaporization” • Gas goes to liquid = condensation - Condensation point = boiling point • At boiling point bubbles of gas form throughout liquid and rise to top • In open container, liquid can all evaporate • In closed container, liquid reaches equilibrium with gas (evaporation rate = condensation rate) • Compounds with stronger intermolecular forces have higher boiling points (H2O higher than F2)

11. Calculations Using Heat of Vaporization • Use heat of vaporization to calculate heat required to vaporize or heat removed to condense (540 cal/g for water) • Heat = mass x heat of vaporization • Example: How much heat is released (in kcal) when 25.0 g of steam condenses at 100.0ºC Heat = 25.0 g x (540 cal/g) = 13500 cal 13500 cal x 1 kcal/1000 cal = 14 kcal

12. Combined Energy Calculations • Calculate each step separately, then total them • Example: How much heat (in kcal) is required to warm 10.0 g of ice from -10.0 ºC to 0.0 ºC, melt it, then warm it to 10.0 ºC ? Heat = mass x T x specific heat = 10.0 g x 10.0 ºC x 1.00 cal/g ºC = 1.00 x 102 cal Heat = mass x heat of fusion = 10.0 g x 80. cal/g = 8.0 x 102 cal Heat = mass x T x specific heat = 10.0 g x 10.0 ºC x 1.00 cal/g ºC = 1.00 x 102 cal Total heat = 1.00 x 102 cal + 8.0 x 102 cal + 1.00 x 102 cal = 1.0 x 103 cal 1.0 x 103 cal x 1 kcal/1000 cal = 1.0 kcal

13. Heating and Cooling Curves