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Honors Chemistry
Chapter 6 – Thermochemistry
Exam Study Guide
Big Idea: Energy is exchanged or transformed in all chemical reactions and physical changes
of matter.
Learning Objectives: Students should be able to:
1.
Calculate the energy change (∆E) of a system given the total heat and/or work transferred
to for from it surroundings.
2.
Describe temperature and heat flow in terms of the motion of molecules (or atoms).
3.
Identify chemical processes that either release (exothermic) or absorb (endothermic)
thermal energy.
4.
Write and interpret thermochemical equations.
5.
Solve problems involving heat flow and temperature changes, using known values of
specific heat.
6.
Apply Hess's law to calculate enthalpy change in a reaction.
7.
o
Use standard heats of formation ( H of ) to find the enthalpy of a reaction ( H rxn
).
Key Equations:
Expressing the change in internal energy in terms of heat and work (p. 188): ∆E = q + w
First law of thermodynamics (law of conservation of energy) (p. 190):
∆Euniverse = ∆Esystem + ∆Esurroundings = 0
Calculating the heat absorbed or released when a substance undergoes a temperature change
(p. 196): q = m x c x ∆T
Calculating the standard heat of reaction (p. 205):
o
H rxn
= m[ H of (products)] – n[ H of (reactants)]
Practice Problems:
1.
A piece of metal is heated and then submerged in cool water. Which statement below
describes what happens?
A) The temperature of the metal will increase.
B) The temperature of the water will increase.
C) The temperature of the water will decrease.
D) The temperature of the water will increase and the temperature of the metal will
decrease.
E) The temperature of the water will decrease and the temperature of the metal will
increase.
2.
If heat is produced by a chemical system, ___________.
A) an equal amount of heat will be absorbed by the surroundings.
B) an equal amount of heat will be absorbed by the universe.
C) an equal amount of heat will be produced by the surroundings.
D) an equal amount of heat will be produced by the universe.
3.
Which of the following equations correctly represents an endothermic reaction?
A) A + B  C + D + heat
B) A + B + heat  C + D
4.
Copper has a specific heat of 0.382 J/g oC. The temperature of an unknown mass of
copper increases by 4.50oC when it absorbs 3.97 J of heat. What is the mass of copper?
5.
Consider the following balanced thermochemical equation for the decomposition of the
mineral magnesite: MgCO3(s)  MgO(s) + CO2(g) Hrxn = 117.3 kJ
(a) Is heat absorbed or released in the reaction? (b) What is Hrxn for the reverse
reaction? (c) What is H when 5.35 mol of CO2 reacts with excess MgO? (d) What is H
when 35.5 g of CO2 reacts with excess MgO?
6.
Hydrazine, N2H4, is a liquid used as a rocket fuel. It reacts with oxygen to yield nitrogen
gas and water.
N2H4(g) + O2(g)  N2(g) + 2H2O(l)
The reaction of 3.60 g of N2H4 evolves 69.8 kJ of heat. Calculate the enthalpy change
per mole of hydrazine for this reaction.
7.
Which of the following chemical reactions or physical changes are exothermic?
A) Ice melting
B) Fog condensing on a windshield
C) Oil burning on water
D) Digestion of food
8.
Calculate q when 0.10 g of ice is cooled from 10.oC to -75oC (cice = 2.087 J/g K).
9.
o
Calculate H rxn
for the following reaction, using standard heats of formation in Appendix B,
p. 820: 2 C2H6(g) + 7 O2(g)  4 CO2(g) + 6 H2O(g)
10.
A system conducts 255 cal of heat to the surroundings while delivering 428 cal of work.
What is the change in internal energy of the system (in cal)?
11.
All of the following statement are true EXCEPT
A) the enthalpy change of an endothermic reaction is positive
B) at constant pressure the heat flow for a reaction equals the change in enthalpy
C) ∆H for a reaction is equal in magnitude but opposite in sign to ∆H for the reverse
reaction
D) enthalpy change is dependent upon the number of steps in a reaction
12.
Determine the heat of vaporization of tin (IV) chloride from the thermochemical equations
below. SnCl4(l)  SnCl4(g)
Sn(s) + 2Cl2(g)  SnCl4(l)
Sn(s) + 2Cl2(g)  SnCl4(g)
∆H = -511.3 kJ
∆H = -471.5 kJ
Solutions:
1) D
2) A
3) B
4) q = m x c x ∆T; m = q / (c x ∆T) = 3.97 J / (0.382 J/g oc)(4.50oC) = 2.31 g
5a) Absorbed
5b) Hrxn (reverse) = –117.3 kJ
 117.3 kJ 
 = –627.555 = –628 kJ
 1 mol CO2 
5c) Hrxn =  5.35 mol CO2  
 1 mol CO2  117.3 kJ 

 = –94.618 = –94.6 kJ
 44.01 g CO2  1 mol CO 2 
5d) Hrxn =  35.5 g CO2  
6) Hrxn = (-69.8 kJ / 3.60 g N2H4) x 32.06 g N2H4/mol = - 622 kJ/mol
7) B, C, & D

8) q(J) = c x mass x T =  2.087

J 
 0.10 g    75  10.  C  = –17.7 = –18 J
g C 
9) H = m[ H (products)] – n[ H of (reactants)]
o
f
o
rxn
o
H rxn
= {4 mol ( H of , CO2(g)) + 6 mol ( H of , H2O(g))} – {2 mol ( H of , C2H6(g)) + 7 mol ( H of ,
O2(g))}
= 4 mol (–393.5
kJ
kJ
) + 6 mol (–241.826 kJ ) – [2 mol (–84.667
) + 7 mol (0)]
mol
mol
mol
= –2855.6 kJ (or –1427.8 kJ for reaction of 1 mol of C2H6)
10) E = q + w = –255 cal + (–428 cal) = –683 cal
11) D
12) 39.8 kJ