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)

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

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

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
Very strong
Moderately strong
Type(s) of
Force ionic
H-bonding, dipoledipole
Charge full charges partial charges
Weak dispersion temporary partial charges
Type of
Compound ionic polar covalent nonpolar covalent

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

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)

• 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

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
)

• 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

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