Unit 3 Study Guide
Energy and Thermochemistry
10.1 - Energy, Temperature, and Heat
1. Explain the difference between kinetic and potential energy in a chemical reaction.
Potential energy: stored in chemical bonds
Kinetic energy: energy absorbed/released in a reaction
2
a. What do potential, kinetic and thermal energy have to do with the law of
conservation of energy?
Law of conservation is that energy cannot be created or destroyed but can be converted to
one form from another. When you have a ball high up, it has potential energy, but when
you drop it, the potential energy converts to kinetic and thermal energy.
b. What do they have to do with the internal energy of a system?
delta E = q + w
Potential, kinetic and thermal energy are the total internal energy of a system.
3. Explain the difference between temperature and heat.
Temperature is the average kinetic energy of the particles while heat is the flow of energy
due to a change in temperature.
4. Give an example of a state function and explain your answer.
Enthalpy is a state function because no matter how many reactions you perform, if you
start with the same reactants and end with the same products, the change in enthalpy will
always be the same.
5. Give an example of a system and explain your answer.
A cup of coffee is a system because it is a point in the universe that we can focus on. Also, all
of the parts are dependent on each other.
10.2 - Endothermic vs. Exothermic
6. Draw the energy diagram for an endothermic and exothermic reaction:
7.a. Which has higher energy in an endothermic reaction, reactants or products?
Products
b. Which has higher energy in an exothermic reaction, the reactants or products?
Reactants
8. Label the process as either exothermic or endothermic and identify your perspective:
a. When solid KBr (potassium bromide) is dissolved in water, the solution gets colder
_____Endothermic (if KBr is the system)_______
b. Natural gas CH4 (methane) is burned in a furnace __Exothermic (if methane is the
system)
c. Concentrated H2SO4 (sulfuric acid) is added to water,
and the solution gets very hot
_____Exothermic (if sulfuric acid is the system)________________
d. Water boils in a teakettle _____Endothermic (if water is the system)________
10.3 - Thermodynamics
9. Explain what each positive or negative sign means:
State Function?
Work (w)
Heat (q)
∆E = energy
Definition
+ positive
- negative
∆H = enthalpy
∆T = temp
10. If a gas absorbs 45 kilojoules of heat, and does 29 kilojoules of work, calculate the
change in energy.
delta E = q + w
delta E = 45 KJ – 29 KJ = 16 KJ
If work is done BY the system, then it is negative.
11. Look at the specific heats listed in Table 10.1 in the textbook. Which substance retains
the most heat? Which substance will have the largest change in temperature when heat
is applied?
Most heat – water (highest specific heat)
Least change in temperature – gold (smallest specific heat)
12. How much energy is required to heat 7.40 mL of water from 25°C to 46°C?
q = mc delta T
7.40 mL X 1g/1mL = 7.40 g water
q = (7.40 g) (4.184 J/gC) (21 C)
q = 650. J
13. How much energy is required to heat a piece of iron (Fe) that weighs 1.3 grams from
25°C to 46°C? (Use table 10.1)
q = (1.3 g) (0.45 J/gC) (21 C) = 12 J
14. If it takes 5.8 joules of energy to heat a piece of metal that weighs
1.6 grams from 23°C
to 41°C, what is the specific heat of this metal? Is this metal pure gold? Why or why
not? (Consult table 10.1)
q=mc delta T
5.8 J = (1.6 g) (c) (18 C)
5.8 J = 28.8 gC (c)
specific heat (c) = 0.20 J/g C
15. How much energy is required to melt 23 g of ice?
delta H fus H2O = 6.01 KJ/mol
q = delta H fus x mass
23 g x 1 mol / 18.02 g = 1.27 mol
q = 6.01 KJ/mol x 1.3 mol
q = 7.8 KJ
16. Explain how a calorimeter could be used to solve question 15 experimentally. Would
you expect the answer to be exactly the same as what you calculated in 15? Why or why
not?
Put 23 g of ice into water at room temperature, measure the change in temperature (delta
T), and from that use q = mc delta T to find the energy absorbed by the water (which is the
same as the energy released from the ice). The energy would be different due to energy
escaped from the calorimeter.
Thermochemistry (Enthalpy) and Hess’s Law
17. The equation for the fermentation of glucose to alcohol and carbon dioxide is:
C6H12O6  2 C2H5OH + 2 CO2
∆H = - 67 kilojoules
a. Is this reaction endothermic or exothermic? Why?
b. Is energy released or absorbed?
c. How much heat is released when 25 moles of glucose is fermented?
18.
S + 3/2 O2  SO3
∆H = - 395.2 kilojoules
2SO2 + O2  2 SO3
∆H = - 198.2 kilojoules
Calculate the Enthalpy (∆H) for the overall reaction:
S + O2  SO2
19.
2O3  3O2
∆H = - 427 kilojoules
O2  2 O
∆H = + 459 kilojoules
NO + O3  NO2 + O2
∆H = - 199 kilojoules
Calculate the Enthalpy (∆H) for the overall reaction:
NO + O  NO2
Reverse reaction 1
Reverse reaction 2
Multiply reaction 3 by 2
Add the reactions, then divide by 2
delta H = +427 KJ – 459 KJ – 2 (199 KJ) = -430 KJ / 2 = -215 KJ