1.

Name________________________________________period________ AP Chemistry
Unit 7 worksheet
1. What is energy? The measure of the ability to do work
2. What is heat? A form of energy that is transferred because of temperature differences
3. What is the SI unit for heat? Joule
4. What is an exothermic reaction? What is the sign of ΔH for an exothermic reaction?
Heat is released; negative
5.
What is an endothermic reaction? What is the sign of ΔH for an endothermic reaction?
Heat is absorbed; positive
6.
Consider the following reaction, which occurs at room temperature and pressure:
2Cl (g) Cl2 (g) ΔH = -243.4 kJ
Which has higher enthalpy under these conditions, 2Cl (g) or Cl2(g)
7.
Consider the following reaction:
2Mg(s) + O2(g) 2MgO(s) ΔH = -1204 kJ
a. Is this reaction exothermic or endothermic? Exothermic
b. Calculate the amount of heat transferred when 2.4 g of Mg react.
-59 kJ
c. Will the surroundings get warmer or colder when the reaction proceeds?
Warmer
d. How many kilojoules of heat are absorbed when 7.50 g of MgO decomposes?
112 kJ
8.
How much heat is released when 15.0 g of copper with a specific heat capacity of 0.385 J g-1K-1 is cooled from 80
⁰C to 35 ⁰C?
260 J
9. How many kilojoules of heat are needed to raise the temperature of 10.00 kg of liquid water from 24.6 °C to
46.2 °C?
904 kJ
10. If 500 J of heat are added to 100.g sample of each of the substances listed below, which will have the largest
temperature increase?
Gold specific heat = 0.129 J g-1K-1
Silver specific heat = 0.237 J g-1K-1
Copper specific heat = 0.385 J g-1K-1
Water specific heat = 4.18 J g-1K-1
11. If 400.0 J of heat are added to a 150.0 g sample of water at 25.0 ⁰C, what is the final temperature of the water?
25.6 °C
12. What are more thermodynamically favored, exothermic or endothermic reactions?
13. The heat released from the combustion of 0.0500 g of white phosphorus increases the temperature of 150.0 g
of water from 25.0 ⁰C to 31.5 ⁰C. Calculate the value of the enthalpy change in kJ mol-1 of the combustion of
phosphorus.
-2,540 kJ/mol
14. When a 6.50 g sample of solid sodium hydroxide dissolves in 100.0 g of water in a Styrofoam cup, the
temperature rises from 21.6 ⁰C to 37.8 ⁰C. Calculate ΔH (in kJ mol-1) for the solution process. Assume the
specific heat of the solution is the same as pure water.
-44.4 kJ/mol
15. Consider the combustion of liquid methanol:
CH3OH (l) + 3/2 O2(g) CO2 (g) + 2H2O (l)
ΔH = -726.5 kJ
a. What is the enthalpy change for the reverse reaction? 726.5 kJ
16. b. Balance the forward reaction with whole-number coefficients. What is ΔH for the reaction represented by this
equation? 2CH3OH (l) + 3 O2(g) 2CO2 (g) + 4H2O (l) ΔH = -1453kJ
c. Which is more likely to be thermodynamically favored the forward or reverse reaction? Forward
d. If the reaction were written to produce H2O (g) instead of H2O (l), would you expect the magnitude of ΔH to
increase, decrease, or stay the same? Explain.
It takes energy to go from liquid water to gaseous water, so ΔH would increase (less negative)
17. Given the following enthalpies of reaction:
P4(s) + 3O2(g) P4O6 (s) ΔH = -1640.1
P4(s) + 5O2(g) P4O10(s) ΔH = -2940.1
Calculate the enthalpy change for P4O6(s) + 2O2(g) P4O10(s)
-1300 kJ
18. Given the following enthalpies of reaction:
H2(g) + F2(g) 2HF(g)
ΔH = -537kJ
C(s) + 2F2(g) CF4(g)
ΔH= -680 kJ
2C(s) + 2H2(g) C2H4(g)
ΔH=52.3 kJ
Calculate the enthalpy change for C2H4 (g) + 6F2(g) 2CF4(g) + 4HF(g)
-2490 kJ
19. Which of the following does not have a standard heat of formation value of zero at 25 ⁰C and 1.00 atm?
a. Cl2(g)
b. I2 (s)
c. Br2(g)
d. Na(s)
20. Calculate the enthalpy change for the reactions
a.Fe3O4 (s) + 2C(s) 3Fe(s) + 2CO2 (g) -670 kJ
b. SiCl4(l) + 2H2O(l)  SiO2(s) + 4HCl(g) -68.3 kJ
c. 4FeO(s) + O2 (g) 2Fe2O3(s)
-556.7 kJ
21. Is breaking bonds endothermic or exothermic? Is forming bonds endothermic or exothermic?
Endothermic; exothermic
22. Using the given bond energies, find the enthalpy change for the following reactions.
a. C2H4 + H2 C2H6
bond
bond energy
bond
-124 kJ
b. 2H2 + O2 2H2O
C-H
413 kJ/mol
H-H
-484 kJ
C-C
348 kJ/mol
O=O
C=C
614 kJ/mol
O-H
bond energy
436 kJ/mol
495 kJ/mol
463 kJ/mol
23. a. What is entropy? Disorder or randomness of a system
b. During a chemical process the system becomes more disordered. What is the sign of ∆S for the process?
positive
24. How does the entropy of the system change when the following occur?
a. A solid melts b. a liquid vaporizes
c. a solid dissolves in water
d. a gas liquefies
Increases
increases
increases
decreases
25. For each of the following pairs, choose the substance with the higher entropy per mol at a given temp.
a. Ar(l) or Ar(g)
b. He(g) at 3 atm or He(g) at 1.5 atm
26. Predict the sign of the entropy change for each of the following
a. 2SO2(g) + O2(g)  2 SO2(g)
negative
b. Ba(OH)2 (s)  BaO(s) + H2O(g) positive
c. CO(g) + 2H2(g)  CH3OH (l)
negative
d. FeCl2(s) + H2(g)  Fe(s) + 2HCl (g) positive
27. Calculate ∆S, given the following S values
2CH3OH(g) + 3O2(g)  2CO2(g) + 4H2O(g)
92.3 J/K
28. a. Write a balanced chemical equation when cyclohexane (C6H12) combusts. C6H12 +9 O2  6CO2 + 6H2O
b. Predict whether ∆G for this reaction is more or less negative than ∆H. because ∆S is negative, ∆G is less negative
than ∆H
29. For a certain chemical reaction, ∆H = -35.4 kJ and ∆S=-85.5 J/K
a. Is the reaction exothermic or endothermic? exothermic
b. Does the reaction lead to an increase or decrease of disorder? decrease
c. Calculate ∆G for the reaction at 298 K. -9.9 kJ/mol
d. Is this reaction spontaneous at 298 K? yes, it is spontaneous
30. Find ∆G at 25 °C for each and tell if it is spontaneous or not.
a. N2 + 3F2  2NF3
∆H = -249 kJ
∆S= -278 J/K
b. N2F4 2NF2
∆H = 85 kJ
∆S= 198 J/K
c. N2 + 3Cl2  2NCl3
∆H = 460 kJ
∆S= -275 J/K
-166 kJ; spontaneous
26.0 kJ; nonspontaneous
542 kJ; nonspontaneous
31. A particular reaction is spontaneous at 450 K. The reaction is endothermic by 34.5 kJ. What can you conclude
about the sign and magnitude of ∆S for the reaction? S > +76.7 J/K
32.
sample of CH3CH2NH2 is placed in an insulated container, where it decomposes into ethene and ammonia
according to the reaction represented above.
Substance
Absolute Entropy, S°,in J/(mol⋅K) at 298 K
CH3CH2NH2(g)
284.9
CH2CH2(g)
219.3
NH3(g)
192.8
a. Using the data in the table above, calculate the value, in J/(molrxn⋅K), of the standard entropy change, ΔS°,
for the reaction at 298 K.
127.2 J/(molrxn K)
b. Using the data in the table below, calculate the value, in kJ/molrxn , of the standard enthalpy change, ΔH°,
for the reaction at 298 K.
49 kJ/molrxn
c. Based on your answer to part (b), predict whether the temperature of the contents of the insulated container
will increase, decrease, or remain the same as the reaction proceeds. Justify your prediction.
The temperature of the contents should decrease because the
reaction is endothermic, as indicated by the positive ΔH°.
d. Is the reaction spontaneous at 298 K?
No, it is nonspontaneous
Review:
33. An unknown element is shiny and a good conductor of electricity. What other properties would you predict for
it?
Malleable, ductile, high melting point/boiling points
34. A certain ion has an atomic number of 16, a mass number of 33, and 18 electrons.
a. What is the charge on the ion? -2
b. What is the identity of this ion? sulfur
c. How many neutrons does the nucleus of this ion have? 17
35. Find the mass for each of the following
a. 6.75 mol zinc nitrate
1280 g
36. Find the number of moles in each of the following
b. 3.01 x 10 23 atoms of F
9.5 g
a. 0.11 kg sodium oxide
1.8 mol
37. Label each as forming ionic, covalent, or metallic bonds
a. AlN
b. CO2
c. Al foil
d. H2O
ionic
covalent
metallic
covalent
e. SnO
ionic
b. 2.25 x 10 25 atoms Zn
37.4 mol
f. CuF2
ionic
38. Which of the following would have the highest boiling point and why? H2O or H2S
Water has hydrogen bonding holding the molecules together. Hydrogen sulfide has dipole-dipole which are not as
strong
39. What kind of attractive forces must be overcome to
a. Boil water
b. Melt KCl
c. Sublime(solid to gas) I2
hydrogen bonding
ionic
London dispersion
d.Boil H2S
dipole-dipole
40. Nitrogen and carbon monoxide have almost equal masses. Explain why the boiling point of carbon monoxide is
slightly higher than that of nitrogen
It is polar so it has dipole-dipole forces where nitrogen only has London dispersion which is weaker.
41. CH4(g) + 2Cl2(g)  CH2Cl2 (g) + 2HCl(g)
Methane gas reacts with chlorine gas to form dichloromethane and hydrogen chloride, as represented by the
equation above.
a. A 25.0 g sample of methane gas is placed in a reaction vessel containing 2.58 mol of Cl2(g).
(i) Identify the limiting reactant when the methane and chlorine gases are combined. Justify your answer
with a calculation.
Cl2 is the limiting reactant because, in order to react with the given amount of CH4 , more moles of Cl2 are required than
the 2.58 moles of Cl2 that are present.
25.0 g CH4 × 1 mol CH4/16.04 g CH4 x 2 mol Cl2/1 mol CH4 = 3.12 mol Cl2
(ii) Calculate the total number of moles of CH2Cl2(g) in the container after the limiting reactant has been
totally consumed.
1.29 mol CH2Cl2
Initiating most reactions involving chlorine gas involves breaking the Cl–Cl bond, which has a bond energy
of 242 kJ mol-1.
b. Calculate the amount of energy, in joules, needed to break a single Cl–Cl bond.
4.02 × 10−19 J
c. Calculate the longest wavelength of light, in meters, that can supply the energy per photon necessary to
break the Cl–Cl bond.
4.9 × 10−7 m
42. Write a balanced net ionic reaction
a. A sample of solid iron(III) oxide is reduced completely with solid carbon.
(i)
Balanced equation:
2 Fe2O3 + 3 C → 4 Fe + 3 CO2
OR
Fe2O3 + 3 C → 2 Fe + 3 CO
(ii) What is the oxidation number of carbon before the reaction, and what is the oxidation number of
carbon after the reaction is complete?
The oxidation number of C before the reaction is 0, and the oxidation number of C after the reaction is +4
(needs to be consistent with i)
b. Solid mercury(II) oxide decomposes as it is heated in an open test tube in a fume hood.
(i) Balanced equation:
2 HgO → 2 Hg + O2
(ii) After the reaction is complete, is the mass of the material in the test tube greater than, less than, or
equal to the mass of the original sample? Explain.
The mass of the contents of the test tube will decrease
owing to the loss of O2 gas to the atmosphere.
42. Answer the following questions related to sulfur and one of its compounds.
a. Consider the two chemical species S and S2 - .
(i)
Write the electron configuration (e.g., 1s2 2s2 . . .) of each species.
S : 1s2 2s2 2p6 3s2 3p4
S2− : 1s2 2s2 2p6 3s2 3p6
Note: Replacement of 1s2 2s2 2p6 by [Ne] is acceptable.
(ii)
Explain why the radius of the S2− ion is larger than the radius of the S atom.
The nuclear charge is the same for both species, but the eight valence electrons in the sulfide ion
experience a greater amount of electron-electron repulsion than do the six valence electrons in the
neutral sulfur atom. This extra repulsion in the sulfide ion increases the average distance between
the valence electrons, so the electron cloud around the sulfide ion has the greater radius.
(iii)
Which of the two species would be attracted into a magnetic field? Explain.
The sulfur atom would be attracted into a magnetic field. Sulfur has two unpaired p electrons,
which results in a net magnetic moment for the atom. This net magnetic moment would interact
with an external magnetic field, causing a net attraction into the field. The sulfide ion would not
be attracted into a magnetic field because all the electrons in the species are paired, meaning that
their individual magnetic moments would cancel each other
b.The S2− ion is isoelectronic with the Ar atom. From which species, S2− or Ar, is it easier to remove an electron? Explain.
It requires less energy to remove an electron from a sulfide ion than from an argon atom. A valence electron in the
sulfide ion is less attracted to the nucleus (charge +16) than is a valence electron in the argon atom (charge +18).