2. 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)

3. 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

4. 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 10 3 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

5. Attractive Forces between Molecules Molecules are held together in liquids and solids by intermolecular forces Forces are due to attraction of opposite charges Strength of Force Type(s) of Force Charge Type of Compound Very strongionicfull chargesionic Moderately strong H-bonding, dipole- dipole partial charges polar covalent Weakdispersion temporary partial charges nonpolar covalent

7. 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

9. 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)

10. Calculations Using Heat of Fusion Use heat of fusion to calculate heat required to melt or heat removed to freeze (80. cal/g for H 2 O) 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

11. 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 (H 2 O higher than F 2 )

12. 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

13. 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 10 2 cal Heat = mass x heat of fusion = 10.0 g x 80. cal/g = 8.0 x 10 2 cal Heat = mass x  T x specific heat = 10.0 g x 10.0 ºC x 1.00 cal/g ºC = 1.00 x 10 2 cal Total heat = 1.00 x 10 2 cal + 8.0 x 10 2 cal + 1.00 x 10 2 cal = 1.0 x 10 3 cal 1.0 x 10 3 cal x 1 kcal/1000 cal = 1.0 kcal

14. Heating and Cooling Curves