1. Thermochemistry
Unit theme: Energy

2. Thermochemistry Potential energy diagrams Heat capacity
Endothermic and exothermic reactions
enthalpy
Specific heat
Heats of changes of state
Units of heat
calorimetry
Thermochemical equations
Hess’ Law
calorie
joule
Gibb’s Free energy
entropy

3. What evidence can you see that energy is involved?

4. Energy Defined as the capacity to do work or supply heat
Chemical potential energy:
aka chemical energy
Energy stored in chemicals because of their compositions
Different substances store different amount of energies

5. Heat
A form of energy that flows from a warmer object to a cooler object Q
Can’t be measured directly
Can only measure its effect on temperature
When heat is added to a system, its temperature rises
Temperature is not heat!
Temperature is a measure of average kinetic energy

6. Thermochemistry
The study of heat changes that occur during chemical reactions and physical changes of state

7. Units of Heat calorie joule
The quantity of heat needed to raise the temperature of 1 g of pure water by 1 degree Celsius
1 kcal = 1000 calories = 1 Calorie (nutrition)
joule
SI unit, named after British physicist
newtonmeter
1 cal = joules
1000 J = 1 kJ
Commonly used because joules are so small

8. Heat Capacity
The amount of heat it takes to change an object’s temperature by 1 Celsius degree
Depends partly on mass
More mass  greater heat capacity

9. Problem
How many calories are required to heat 32.0 g of water from 25.0oC to 80.0oC? How many joules is this?
Answer:
1760 calories
7360 joules = 7.36 kJ

10. Specific Heat
Not all substances respond the same way to the input of heat
Some get hot much more quickly than others!

11. Specific Heat
The amount of heat required to raise the temperature of 1 g of a substance by 1oC
A measure of how well a substance stores heat energy
Substances with low specific heat (ex. metals) heat quickly, cool quickly
Substances with high specific heat (ex. water) take a long time to heat and cool

12. Using Specific Heat C or Cp Units: J/goC or cal/goC Q = m C T where
Q = total heat change
m = mass
T = temperature change

13. Phase Changes Melting, freezing Boiling, condensing
Sublimation, deposition
Phase changes occur without temperature changes

14. Calculating Energy Changes in Phase Changes
Can’t use specific heat, because no temperature change is involved
Instead, use this formula
Q = m  Heat of fusion (or heat of vaporization)
use heat of fusion for freezing, melting
use heat of vaporization for boiling, condensation

15. Heating/Cooling Curves
Regions A, C, E have temperature change
Q = mCDT
Choose specific heat to match state of matter
Regions B, D are phase changes
B: Q = mHfus
D: Q = mHvap

16. Enthalpy of the reaction
The total energy change associated with a chemical or physical change
Given the symbol DHrxn
DHrxn = (energy of products) – (energy of reactants)

17. Classifying Heat Changes
Thermite reaction
Produces molten iron
Formerly used in welding railroad tracks, shipbuilding
Gives off light and heat!
Highly exothermic

18. Exothermic reactions Energy is released to the surroundings
The temperature of the surroundings increases
The products have less energy than the reactants

19. Endothermic Reactions
Energy is absorbed from the surroundings
The temperature of the surroundings decreases
The products have MORE energy than the reactants

20. Thermochemical Equations
Exothermic reactions
Energy released by system
Can treat energy as a product
DH is negative
2 equivalent ways to write equation:
CaO + H2O  Ca(OH) kJ
CaO + H2O  Ca(OH) DH = kJ

21. Thermochemical Equations
Endothermic reactions
Energy absorbed by system
Can treat energy as a reactant
DH is positive
2 equivalent ways to write equation:
2 NaHCO kJ  Na2CO3 + H2O + CO2
2 NaHCO3  Na2CO3 + H2O + CO2 DH = kJ

22. 2 ways to manipulate thermochemical equations
1) Write equation backwards
Sign of DH must change
A + B  C DH = kJ
C  A + B DH = kJ
2) Multiply everything by a coefficient
Must multiply DH by coefficient, too!
3 A + 3B  3C DH = 3 (+123 kJ) = kJ

23. Problems with thermochemical equations
Consider the equation:
2 NaHCO3  Na2CO3 + H2O + CO2 DH = kJ
How many kJ would be released if 4.5 moles NaHCO3 reacted?

24. Hess’ Law
If you add two or more thermochemical equations to give a final equation, then you can also add the heat changes to give the final enthalpy of reaction.

25. Example
What is the enthalpy change, DHrxn, for the decomposition of hydrogen peroxide?
Target: 2 H2O2(l)  2 H2O(l) + O2(g)
Given:
H2(g) + O2(g)  H2O2(l) DH = kJ
H2(g) + ½ O2(g)  H2O(l) DH = kJ

26. Standard Heats of Formation, DHof
The standard heat of formation of a compound is the change in enthalpy that accompanies the formation of one mole of a substance from its elements in their standard states.
The heat of formation of elements in their standard states is arbitrarily set to zero.

27. Using heats of formation
Thermodynamic stability: a measure of the energy required to decompose the compound
Compounds with large, negative enthalpies of formation are thermodynamically stable
Many heats of formation have been measured. (see Appendix A-6 in textbook)
Another way to do Hess’ Law!
DHrxn = SnDHf(products) – SmDHf(reactants)
where
n represents the coefficients for the products
m represents the coefficients for the reactants

28. Example Problems
Compute DHrxn for the following reaction. (Refer to Appendix A-6)
2NO(g) + O2(g)  2 NO2(g)
4 FeO(cr) + O2(g)  2 Fe2O3(cr)
Ans: kJ, kJ

29. Entropy
Symbol: S
A quantitative measure of the degree of disorder in a system
The greater the disorder, the larger the value of S
Solids have a high degree of order (low entropy)
Gases have a low degree of order (high entropy)
More particles (moles) results in higher entropy

30. Entropy, cont. Systems tend to proceed to higher disorder Examples
Stirring sugar into your coffee
The neatness of your locker
The order of cards in a pack of playing cards after shuffling

31. J. Willard Gibbs
First American to earn a Ph.D. in science from a US university
Yale, 1863
One of nation’s best scientists

32. Will a reaction occur spontaneously?
The answer depends on the balance between enthalpy (heat changes) and entropy
Gibb’s Free Energy
The energy available from the system to do useful work

33. Gibb’s Free Energy DG = DH – TDS
If DG is negative, the reaction will occur spontaneously and can proceed on its own.
If DG is positive, the reaction is nonspontaneous and needs a sustained energy input to proceed.
If DG is zero, the reaction is at equilibrium (both the forward and reverse reactions take place!)