Dr. L. Baxley
Chem 201A
Chapter 6 Notes: Thermochemistry
I. Forms of Energy and Energy Change
A) Heat in Chemical Reactions
examples:
Chemical hand warmers: 4 Fe(s) + 3 O2(g) → 2 Fe2O3 + heat
Chemical ice packs: heat + NH4NO3(s) → NH4+(aq) + NO3–(aq)
B) Kinetic and Potential Energy
Energy, E:
Potential Energy, PE:
Kinetic Energy, KE:
Kinetic energy is related
to thermal energy
Potential energy is related
to chemical energy
Chapter 6 notes / page 1
Dr. L. Baxley
Chem 201A
C) Energy Transfer: Heat and Work
Energy can be transferred as heat or work
Heat, q:
Work, w:
D) System and Surroundings
System:
Surroundings:
E) Sign Conventions
Chapter 6 notes / page 2
Dr. L. Baxley
Chem 201A
Concept check: Identify each of the following energy exchanges as heat or
work and whether the sign of heat or work for the system is positive or
negative.
1. An ice cube melts and cools the water it is in. (ice cube is the system)
2. A bike is pushed up a hill. (bike is the system)
F) Energy units
G) State Functions
State function:
Path Function:
Concept check: Is ΔE a path property or a state property?
II. Enthalpy, H
A) Meaning of Enthalpy
Change in Enthalpy, ΔH:
Chapter 6 notes / page 3
Dr. L. Baxley
Chem 201A
B) Some Standard Enthalpy Changes
Heat of a reaction, ΔHrxn:
2 C2H6(g) + 7 O2(g) → 4 CO2(g) + 6 H2O(g) ∆Hrxn = –2.86 kJ
Heat of fusion, ΔHfus:
NaCl(s) → NaCl(l)
∆H= ∆Hfus
Heat of vaporization, ΔHvap:
H2O (l) → H2O (g) ∆H= ∆Hvap
C) Exothermic and Endothermic
Exothermic:
Endothermic:
Chapter 6 notes / page 4
Dr. L. Baxley
Chem 201A
III. Measuring Heat Transfer: Calorimetry
A) Types of Heat Transfer
Concept check: Which would be most difficult to change its
temperature?
1. A 1.29 g piece of aluminum has a temperature of 26.0 °C.
If it loses 1.90 J of heat to its surroundings, what is its
new temperature.
Chapter 6 notes / page 5
Dr. L. Baxley
Chem 201A
B) Measuring Heat Changes
C) Heat Transfer between two bodies
2. A 60.61 g piece of metal at 85.4 °C is placed into a cup containing 359 g of
water at 20.0 °C. The final temperature is 22.3°C. Cwater = 4.184 J/g°C.
What is the specific heat of the metal?
Chapter 6 notes / page 6
Dr. L. Baxley
Chem 201A
3. A 20.0 g piece of ice at 0.0 ºC is dropped into 78.9 g of ethanol at 60.0 °C,
which melts the ice. The final temperature is 17.3 °C. What is the specific
heat capacity of ethanol?
Cwater = 4.184 J/g °C
∆Hfus of water = 6.01 kJ/mol
IV. Thermochemistry in Chemical Reactions
A. Stoichiometry of Thermochemical Equations
Thermochemical equation
Examples:
H2 (g) + ½ O2 (g) → H2O (l)
2 H2 (g) + O2 (g) → 2 H2O (l)
∆H° = –286 kJ
∆H° = –572 kJ
Chapter 6 notes / page 7
Dr. L. Baxley
Chem 201A
Just In Time Learning questions
3NO2(g) + H2O(l) → 2HNO3(aq) + NO(g) ∆H = –138 kJ
3 moles NO2 or
–138 kJ
–138 kJ
3 moles NO2
1. Write the ratio and its inverse that relates moles of nitric acid formed to kJ of heat.
2. Using one of the mole ratios given above, solve this problem. If 2.50 mole of NO2 react in
water according to the equation above, what is the change in heat, in units of kJ? [–115 kJ]
3. Given the thermochemical equation shown below, calculate the change in heat when 95.4 g
of aluminum reacts with excess oxygen.
4 Al(s) + 3 O2(g) → 2 Al2O3(s)
∆H = –3352 kJ
4. According to the equation above, how many grams of aluminum oxide are produced when
the change in heat is –824 kJ?
In-Class Problem:
4. How many grams of O2 react if 1.280x103 kJ are released in the following
reaction?
CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)
ΔH = –8.90x102 kJ
Chapter 6 notes / page 8
Dr. L. Baxley
Chem 201A
5. In a calorimeter 37.1 mL of 0.350 M H2SO4 is added to 85.7 mL of
0.280 M NaOH. For both solutions, Ti = 25.00 ºC and Tf = 27.86 ºC.
Csolution = 4.18 J/g ºC, dsoln = 1.00 g/mL. What is the ∆H of reaction, in kJ?
B. Hess's Law
Hess's Law:
Chapter 6 notes / page 9
Dr. L. Baxley
Chem 201A
Rules for applying Hess’s Rule:
6. From the known thermochemical equations below, determine the ∆H of the
reaction, Ge (s) + ½ O2 (g) → GeO(s).
GeO2(s) → Ge(s) + O2(g)
2 GeO(s) + O2 (g) → 2 GeO2(s)
∆H = +535 kJ
∆H = −558 kJ
7. Calculate the ∆H for the reaction 2 Ni(s) + ¼ S8(s) + 3 O2(g) → 2 NiSO3(s)
from the following information:
(1) NiSO3(s) → NiO(s) + SO2(g)
(2) 1/8 S8(s) + O2(g) → SO2(g)
(3) Ni(s) + ½ O2(g) → NiO(s)
∆H = 156 kJ
∆H = –297 kJ
∆H = –241 kJ
Chapter 6 notes / page 10
Dr. L. Baxley
Chem 201A
C. Heats of Formation
To write an equation of formation of a compound:
8. Write the equation for the heat of formation of CH3OH(l)
D. Using ∆H°f to determine ∆H° of a reaction
9. Use ∆H°f to calculate the ∆H for the reaction,
2 CH3OH(l) + 3 O2(g)  2 CO2(g) + 4 H2O(g)
Chapter 6 notes / page 11
Dr. L. Baxley
Chem 201A
10. Calculate the standard enthalpy of vaporization for water
Additional Challenge Problem:
Consider the reaction,
4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) ΔH = –1650 kJ
In a calorimeter containing 6.52 kg of water at 12.5 °C a 35.6 g of iron reacts with 95.3 g of
oxygen according to this reaction. What will be the final temperature of the water,
assuming all of the heat from the reaction is transferred to the water?
Chapter 6 notes / page 12
Dr. L. Baxley
Chem 201A
Thermochemistry Concept Map
Chapter 6 notes / page 13