1. Thermochemistry
AP Chem
Ch. 6

2. Thermochemistry
Thermochemistry – the study of heat changes that accompany chemical reactions and phase changes
Universe = System + Surroundings
Endothermic Reaction – one in which energy (heat) is absorbed into the system
Exothermic Reaction – one in which energy (heat) is released from the system

3. Energy the SI unit for heat is the Joule (J)
Energy – the ability to do work or produce heat
the SI unit for heat is the Joule (J)
Heat CANNOT be measured directly; it is calculated
“q” = quantity of heat

4. In the lab…
Calorimeter – an insulated device used for measuring the temperature change during a chemical process

5. Useful Conversion Factors
1 cal = J
1 Cal = 1000 cal
101.3 J = 1 L•atm
1 J = 1 kg•m2 / s2

6. Specific Heat
Specific Heat – energy required to raise the temperature of one gram of any substance one degree Celsius.
Each substance has its own specific heat
The higher the specific heat, the longer it will take to raise the temp.
water has a very high specific heat J/g°C

7. q = m c ΔT q = heat absorbed or released c = specific heat value
m = mass or moles of substance
ΔT = change in temperature
Use this when there is NO change in state
**Make sure units will cancel (usually “c” will determine what units to use for q, m, DT)**

8. Heat Lost = Heat Gained
When 2 substances at different temperatures come into contact with each other, heat is transferred from the warmer to the cooler substance until both substances are at the same temperature

9. Enthalpy Exothermic (lose heat) = -ΔHrxn
Enthalpy (H) – measure of the amount of energy aborbed/released by a reaction
Changes in enthalpy (ΔH rxn )can be calculated for specific chemical reactions
Exothermic (lose heat) = -ΔHrxn
Endothermic (gains heat) = +ΔHrxn

10. Enthalpy (cont’d) ΔH rxn = H products – H reactants
DH = q only at constant pressure
ΔH° = standard enthalpy change
° = occurs under standard conditions
standard conditions = 1atm, 298K (25°C)

11. Thermochemical Equations
A balanced chemical equation that includes all reactants, products, and the change in energy (ΔH)
Use energy in stoichiometric calculations

12. Thermochemical Equations
Can be written 2 ways:
a. 4Fe + 3O2 2Fe2O kJ or
4Fe + 3O2 2Fe2O ΔH = -1625kJ
b. NH4NO3 +27kJ  NH4+ + NO3-
NH4NO3  NH4+ + NO ΔH = 27kJ

13. DHo = bond energies broken – bond energies formed
Bond Energy
Bond energy = energy required to BREAK a bond
Forming bonds = energy is released
Breaking bonds = energy is absorbed
DHo = bond energies broken – bond energies formed

14. Heat in Changes of State
Molar Heat Values (table 11.5 on handout)
ΔH vap = - ΔH cond
ΔH fus = - ΔH solid

15. Heat in Changes of State
q = mDH
m = mole or mass
DH = molar heat value
**units for m depend on units of DH**

16. Heating/Cooling Curve

17. Changes in Temp & State Use Heating/Cooling Curve
Use q=mcDt and q=mDH to calculate total heat

18. Heat of Formation
Heat of formation-change in enthalpy that takes place when a compound is formed from it’s elements
ΔH°f = standard heat of formation = the change in enthalpy that accompanies the formation of 1 mole of compound with all substances in their standard states
Table 11.6 in handout
all in standard states: 1atm, 298K (25°C)

19. ΔH°rxn = Σ ΔH°f (products) - Σ ΔH°f (reactants)
Heat of Formation
ΔH°rxn = Σ ΔH°f (products) - Σ ΔH°f (reactants)
**ΔH°f of a pure element in it’s standard state is 0.0 kJ**

20. Hess’s Law Add two or more thermochemical equations
together to obtain a desired thermochemical equation
Multiply DH by same coefficient
Divide DH by same coefficient
Change sign of DH
Multiply rxn by a coefficient
Divide rxn by a coefficient
Reverse rxn