1

 Introductory Objectives
1. Explain the relationship between energy and heat.
2. Distinguish between heat capacity and specific heat.
TIP : Do not confuse standard conditions with
STP used in gas law calculations.
2



Energy - the capacity for doing work or supplying heat
Chemical potential energy – the energy stored within the structural units of chemical substances
 Different substances store different amounts of energy.
The kinds of atoms and their arrangement in the substance determine the amount of energy stored in the substance.


All energy in a process can be accounted for as work, stored energy, or heat
Law of Conservation of Energy – In any physical or chemical process, energy is neither created not destroyed
3

 Heat – energy that transfers from one object to another because of a temperature difference between them
 Cannot be detected by the senses or instruments
 Only changes caused by heat can be detected
 Always flows from a warmer object to a cooler object
 If two objects remain in contact, eventually the temperature of both objects will be the same
4

 System- part of the universe on which you focus your attention
 Surroundings – include everything else in the universe
 Universe – the system and the surroundings
 Example: Chemicals and water are in a beaker. (Universe) Your system includes the chemicals and water. The beaker is the surrounding.
5

 Endothermic reaction - heat, q, flows into a system (heat absorbed), >0 (positive number)
Examples: melting of ice, evaporation of a puddle, sublimation of a mothball, heat used to cook food
During endothermic phase changes, energy absorbed does not increase the temperature because the energy is being used to overcome attractions between particles.
Bond-breaking
6

 Exothermic reaction – heat, q, flows out of the system (heat is given off), <0 (negative number)
Examples: combustion of fossil fuels, cooling of skin as perspiration evaporates, freezing of water
Bond-formation
7

 Temperature – a measurement of the average kinetic energy of the particles
Can be detected with a thermometer
 Heat cannot be measured with a thermometer
 Heat can also increase the potential (rather than kinetic) energy. This occurs during phase changes: solid to liquid AND liquid to gas
8

 Suppose two identical candles are used to heat two samples of water. One sample is a cup of water; the other is 10 gallons of water in a drum.
 1. How will the change in temperature of the samples compare?
 Practically no change in the drum; a large increase in the cup
 2. How will the amount of heat received by each container compare?
 Both containers receive the same amount of heat
9

 Joule – the SI unit of heat and energy
A joule of heat raises the temperature of 1 g of pure water 0.2390 °C
 1 J = 0.2390 cal 4.184 J = 1 cal
 Heat capacity – amount of heat needed to increase the temperature of an object exactly
1 °C
 Besides varying with mass, the heat capacity of an object also depends on its chemical composition
10

 calorie- the quantity of heat needed to raise the temperature of 1g of pure water 1 °C
 Calorie = 1000 calories (refers to energy in food)
 1 Calorie = 1kilocalorie = 1000 calories
 “10g of sugar has 41 Calories” means that
10g of sugar releases 41 kilocalories of heat when completely burned to produce carbon dioxide and water
11

 What is the relationship between a joule and a calorie? Calorie and a dietary calorie?
 What is the difference between specific heat capacity and heat capacity? Give examples.
12

 Note that the temperature of water changes less than the temperature of iron because the specific heat capacity of water is larger.
 Specific heat capacity (specific heat) – the amount of heat it takes to raise the temperature of 1 g of the substance 1 °C
 Specific heat (C) is a measure of a substance to store heat. The specific heats of substances can be compared because the quantity (1 g) of matter involved is specified.
13

 C = q ÷ (m x T)
C or c p
= specific heat at a given pressure q = energy lost or gained (heat) m = mass of the sample
T = difference between initial temperature and final temperature
 q = c p x m x T
14

 1. Construct equations that show the heat changes for chemical and physical processes
 2. Calculate heat changes in chemical and physical processes
 Think!
A match won’t ignite unless you strike it and add the heat produced from friction. Is the burning of a match an endothermic reaction?
 Is there a way to measure how much heat is released from a burning match?
15

 No; the reaction releases more energy in the form of heat and light than the amount of energy it absorbs to start.
 Yes, but only indirectly. If the reaction were confined, then any temperature changes in the surroundings could be attributed to heat transfer from the reaction.
16

 The accurate and precise measurement of heat change for chemical and physical processes
 Need insulated container
1. Constant pressure calorimeter
2. Bomb calorimeter – constant volume
Measures the heat released from burning a compound; closed system: the mass of the system is constant
 The heat released by the system is equal to the heat absorbed by its surroundings
17








Heat changes for reactions carried out at constant pressure
Because the reactions presented in the textbook occur at constant pressure, the text uses heat and enthalpy interchangeably
Heat change for a chemical reaction carried out in aqueous solution: q = H = m x C x T
Reacting chemicals = system
Known volumes of water = surroundings
Exothermic – negative number
Endothermic – positive number
18

 An equation that includes the heat change
 Heat of reaction – the heat change for the reaction exactly as it is written (Usually heat change at constant pressure)
 The physical state of the reactants and products must be given
 Standard conditions = 101.3 kPa (1atm) and
25 °C
 Amount of heat absorbed or released depends on the number of moles
19

 Heat of reaction for the complete burning of one mole of a substance
 Like other heats of reaction, heats of combustion are reported as the enthalpy changes when the reactions are carried out at
101.3 kPa of pressure and the reactants and products are in their physical states at 25 °C
20

 Classify, by type, the heat changes that occur during melting, freezing, boiling, and condensing
 Calculate heat changes that occur during melting, freezing, boiling, and condensing
21







Specific heat capacity (specific heat) – the amount of heat it takes to raise the temperature of 1 g of the substance 1 °C
C = q ÷ (m x T)
Enthalpy (H) - Heat changes for reactions carried out at constant pressure
Heat change for a chemical reaction carried out in aqueous solution: q = H = m x C x T
Like other heats of reaction, heats of combustion are reported as the enthalpy changes when the reactions are carried out at
101.3 kPa of pressure and the reactants and products are in their physical states at 25 °C
H is enthalpy or heat content
H represents a change in the heat content
22

 Solid ----------------Liquid -------------------Vapor
+ Fusion +Vaporization
Low enthalpy-----------------------High enthalpy
 Molar heat of fusion ( H fus
) – heat absorbed by one mole of a substance in melting from a solid to a liquid at a constant temperature
 Molar heat of vaporization ( H vap
) – the amount of heat necessary to vaporize one mole of a given liquid
 Endothermic reactions
23

 Vapor------------------Liquid------------------Solid
-Condensation -Solidification
High enthalpy-----------------------Low enthalpy
 Molar heat of condensation ( H cond
) – amount of heat released when one mole of vapor condenses
 Molar heat of solidification ( H solid
) – the heat lost when one mole of a liquid solidifies at a constant temperature
 Exothermic reactions
24

 Solid ----------------Liquid -------------------Vapor
+Fusion +Vaporization
-Solidification -Condensation <
 The molar heat of fusion is the heat absorbed by one mole of a substance in melting from a solid to a liquid at a constant temperature.
The heat lost when one mole of a liquid solidifies at a constant temperature is the molar heat of solidification. Because energy is conserved in all chemical and physical changes, the quantity of heat absorbed by the melting solid must equal the quantity of heat lost when the liquid solidifies.
25



H fus
= - H solid
H vap
= - H cond
Values are numerically the same, but the values have different signs
 Fusion —endothermic—Vaporization (+)
 Solidification – exothermic—Condensation (-)
 The melting of one mole of ice at 0 °C to one mole of water at 0 °C requires the absorption of 6.01 kJ of heat. What is the heat of fusion?
26

 The heat of fusion is 6.01 kJ/mol.
 The heat of solidification is -6.01 kJ/mol.
 Ice is commonly used to refrigerate perishable foods. What happens to the temperature of the ice as it begins to melt?
 The ice and the water are both at 0 °C. The temperature will not rise above 0 °C until all of the ice has melted.
27

 How many grams of ice at 0 °C and 101.3 kPa could be melted by the addition of 2.25 kJ of heat?
 Standard conditions for ice exist. Use heat of fusion for water.
 Grams for one mole (18g)/6.01 kJ = x/2.25 kJ
 Answer = 6.74 g ice
28

 H soln
– heat change caused by the dissolution of one mole of a substance
 Examples:
1. Exothermic molar heat of solution: sodium hydroxide dissolved in water, hot pack that mixes calcium chloride and water
2. Endothermic molar heat of solution: cold pack that allows water and ammonium nitrate to mix (Heat is released from the water and the temperature of the solution decreases.)
29

 How much heat (in kJ) is absorbed when
24.8 g of H
2
O( l ) at 100 °C is converted to steam at 100 °C ?
 Use heat of vaporization for water.
 24.8g/x = 18g/40.7 kJ
 Answer = 56.1 kJ
30

 Apply Hess’s law of heat summation to find heat changes for chemical and physical processes
 Calculate heat changes using standard heats of formation
31

 Hess’s law of heat summation – If you add two or more thermochemical equations to give a final equation, then you can also add the heats of reaction to give the final heat of reaction
32

 Standard heat of formation ( H f
°) – the change in enthalpy that accompanies the formation of one mole of a compound from its elements with all substances at their standard states at 25 °C
 The standard heat of formation of a free element in its standard state is arbitrarily set at 0. (Includes diatomic molecules and graphite form of carbon)
33

 The standard heat of reaction ( H °) is the difference between the standard heats of formation of all the reactants and products.
 H ° = H f
° (products) - H f
° (reactants)
34