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 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

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

ChargeType of

Compound

Very strong ionic full charges ionic

Moderately strong

H-bonding, dipole-dipole

partial charges

polar covalent

Weak dispersiontemporary

partial charges

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)

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

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)

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

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

Heating and Cooling Curves