1

Chapter 11: Thermochemistry

The Two-Day Chapter Extravaganza

2

Thermochemistry- heat changes that occur during chemical reactions

Introductory Objectives

1. Explain the relationship between energy

and heat

2. Distinguish between heat capacity and

specific heat

3. Know the key terms on page 321

TIP : Do not confuse standard conditions with

STP used in gas law calculations.

3

Energy

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

4

Heat (q)

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

5

Systems

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.

6

Endothermic and Exothermic Reactions

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

7

What value of q is endothermic?

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

8

Heat vs. Temperature

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

9

Think!

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

10

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

11

Heat Capacity and Specific Heat

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

12

Do Now

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.

Explain how you could manipulate a liquid to bring it to a boil.

13

Refer to Table 11.2 on page 296 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.

14

Specific Heat

Memorize the formula on page 297 C = q ÷ (m x ΔT) or q = m C ΔT where:

C = specific heat capacity

q = heat transferred

ΔT = change in temperature

m = mass Review Sample-Problem 11-1 on page 299 Review homework questions 1-10 on page

299

15

Objectives

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?

16

Answer to Think!

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.

17

Calorimetry

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

18

Enthalpy (H)

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 = ΔH = m x C x ΔT Reacting chemicals = system Known volumes of water = surroundings Exothermic – negative number Endothermic – positive number

19

Thermochemical Equations

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)

Refer to page 303 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

20

Heat of Combustion

Heat of reaction for the complete burning of one mole of a substance

Refer to Table 11.4 on page 305 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

21

Objectives

Review sections 1 and 2 Explain question 16 on page 306 Know key terms and concepts Complete Interpreting Graphics Handout 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

22

Review

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 ΔT)

Enthalpy (H) - Heat changes for reactions carried out at constant pressure

Heat change for a chemical reaction carried out in aqueous solution:

q = ΔH = m x C x ΔT 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 ΔH represents a change in the heat content

23

Refer to Question 16 on page 306

To solve: Calculate the molar mass of ethanol

Answer: One mole of ethanol contains 46 grams.

12.5 grams (given in problem) is .272 mole .272 x -1235 kJ (given in problem) = -335.9 Heat given off = 3.36 x 102 kJ

24

Key Terms and Concepts

Solid ----------------Liquid -------------------Vapor

+ Fusion +Vaporization

Low enthalpy-----------------------High enthalpy Molar heat of fusion (ΔHfus) – heat absorbed

by one mole of a substance in melting from a solid to a liquid at a constant temperature

Molar heat of vaporization (ΔHvap) – the amount of heat necessary to vaporize one mole of a given liquid

Endothermic reactions

25

Vapor------------------Liquid------------------Solid

-Condensation -Solidification

High enthalpy-----------------------Low enthalpy Molar heat of condensation (ΔHcond) –

amount of heat released when one mole of vapor condenses

Molar heat of solidification (ΔHsolid) – the heat lost when one mole of a liquid solidifies at a constant temperature

Exothermic reactions

26

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.

27

ΔHfus = - ΔHsolid

ΔHvap = - ΔHcond

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?

28

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.

Complete the Interpreting Graphics Handout

29

Problem Solving 1

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

30

Heat of Solution

ΔHsoln – heat change caused by the dissolution of one mole of a substance

Examples: (Refer to page 312) 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.)

31

Problem Solving 2

How much heat (in kJ) is absorbed when

24.8 g of H2O(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

32

Objectives

Apply Hess’s law of heat summation to find heat changes for chemical and physical processes

Calculate heat changes using standard heats of formation

33

Hess’s Law of Heat Summation

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

Refer to question 32 on page 318 Reverse the second enthalpy change

(change sign to +) and cancel the oxygen Subtract +824.2 kJ from -1669.8 kJ Answer = -8.456 x 102 kJ

34

Standard Heats of Formation

Standard heat of formation (ΔHf°) – 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)

Refer to Table 11.6 on page 316

35

Standard Heat of Reaction

The standard heat of reaction (ΔH°) is the difference between the standard heats of formation of all the reactants and products.

ΔH° = ΔHf° (products) - ΔHf° (reactants)

Refer to Sample Problem 11-7 on page 317.