1. Chapter 17
Thermochemistry

5. Law of conservation of energy
The law of conservation of energy states that energy can be neither created or destroyed but can be converted from one form to another.

6. Heat
Heat, which is represented by the symbol q, is energy that is in the process of flowing from a warmer object to a cooler object.
When the warmer object loses heat, its temperature decreases.
When the cooler object absorbs heat, its temperature rises.

7. Measuring heat
In the metric system, the amount of heat required to raise the temperature of one gram of pure water by one degree Celsius (1°C) is defined as a calorie (cal).
The SI unit of heat and energy is the Joule (J). One Joule is the equivalent of calories, or one calorie equals Joules.

8. Temperature
Temperature is a measure of the average kinetic energy of the particles in a sample of matter.
Temperature is measured in °C or K

10. Important Concept
Heat and
temperature
are not the
same thing

11. Specific Heat
Specific heat (c) is the amount of energy that is required to raise the temperature of one gram of pure water by one degree Celsius (1°C).
The specific heat of water is 4.18 J/g∙°C

13. Specific Heat
Because different substances have different compositions, each substance has its own specific heat.
Substance Specific Heat J/g ∙ °C
Water
Ethyl Alcohol 2.44
Aluminum
Copper
Gold

14. Calculating Specific Heat
The heat absorbed or released by a substance during a change in temperature depends not only upon the specific heat of the substance, but also upon the mass of the substance and the amount by which the temperature changes.

15. Calculating Specific Heat
The formula for calculating specific heat is:

16. Calculating Specific Heat
q = the heat absorbed or released
c = the specific heat of the substance
m = the mass of the sample in grams
∆T is the change in temperature in °C
∆T is the difference between the final temperature and the initial temperature or, Tfinal – Tinitial.

17. Try this Problem
How much heat is given off when an 869 g iron bar cools from 940.0°C to 50.0°C?
c Fe = J/g • °C
m = 869 g
ΔT = Tfinal – Tinitial = 50.0°C – 940.0°C = -890°C
q = c m ΔT
q = J/g • °C x 869 g x –890.0°C
q = -343,000 J

18. Enthalpy
Enthalpy is the amount of energy absorbed or lost by a system during a process at constant pressure.
Types of enthalpy
Heat of Reaction
Heat of Formation
Heat of Combustion
Heat of Vaporization
Heat of Fusion

20. Thermochemical Equation
A thermochemical equation is a balanced chemical equation that includes the physical states of all reactants and products and the energy change, usually expressed as the change in enthalpy, ∆H.
2H2(g) + O2(g)  2H2O(g) kJ

22. Enthalpy and Change in Enthalpy
Although you cannot measure the actual energy or enthalpy of a substance, you can measure the change in enthalpy, which is the heat absorbed or released during a chemical reaction.
The change in enthalpy for a reaction is called the Heat of Reaction (∆Hrxn).
You have already learned that ∆ means a change in the property.

23. Heat of Reaction ∆Hrxn is the difference between the
enthalpy of the substances that exist at the end of the reaction and the enthalpy of the substances present at the start.
The formula for Heat of Reaction is
ΔHrxn = H°f products - H°f reactants

24. Heat of Combustion
The Heat of Combustion is the amount of energy released as heat by the complete combustion of one mole of a substance.
The formula for Heat of Combustion is
ΔHc = H°f products - H°f reactants
As you can see, it is the same as Heat of reaction except for the beginning: ΔHc

25. Heat of Formation
Heat of formation (H°f) is the energy needed to form one mole of a compound from its elements.
Elements have 0 H°f
A Heat of Formation table is found in your textbook on page 902

26. Heat of Vaporization
Heat of Fusion
Heat can also be absorbed or released during a change of state.
The heat required to vaporize one mole of a liquid is called its heat of vaporization (∆Hvap).
The heat required to melt one mole of a solid is its heat of fusion (∆Hfus).

27. Solve This Problem ΔHc = H°f products - H°f reactants
ΔH°f CO2(gas) = kJ/mol
ΔH°f H2O(liquid) = kJ/mol
ΔH°f C5H12(gas) = kJ/mol
Calculate Heat of Combustion for this reaction
C5 H12 (g) + 8O2(g)  5CO2(g) + 6 H2O(l)
ΔHc = H°f products - H°f reactants
ΔHc = [(5)( kJ/mol) + (6)( kJ/mol)] - [ kJ/mol + 0 kJ/mol]
= kJ or kJ

28. Endothermic and Exothermic
Endothermic reactions absorb energy from the system. The container gets cold.
Exothermic reactions release energy from the system. The container gets hot.

32. Endothermic Reactions
The products have more energy than the reactants. The products have absorbed energy.

33. Exothermic Reactions
The products have less energy than the reactants. The products have given off energy.

34. Changes in Enthalpy

35. Entropy (S)
Entropy is a measure of the degree of randomness of the particles in a system.
Gases have high entropy and solids have low entropy.
Does your bedroom
have high Entropy?

37. Reaction Rate
Reaction rate is the change in concentration of reactants per unit time as a reaction proceeds.
Factors that influence reaction rate:
Nature of reactants – some substances tend to react
faster than other.
2. Surface area – The more surface area, the faster a
reaction will occur.
3. Temperature – Higher temperatures make a reaction
rate increase.
4. Concentration – Higher concentrations make a reaction rate increase.
5. Presence of a catalyst – Can speed up or slow down a reaction depending on the catalyst.

38. Phase Changes
Phase changes occur when a substance changes from one state to another.
For a phase change to occur, energy must be added to or taken away from the substance.
Phase changes are physical changes.

40. Phase Changes
Heating Curves
and

41. Energy Changes Terminology
At the melting/freezing point
Heat of Fusion
Heat of solidification
At the boiling/condensation point
Heat of Vaporization
Heat of Condensation

42. THE END