Download Honors Chemistry 2 Chapter 10 Test Review

Honors Chemistry 2
Chapter 10 Test Review Sheet
Cunnings, Fall, 2014
IMPORTANT NOTES:




There are MORE THAN ENOUGH math practice problems on this review sheet. Your test won’t be
nearly this long, but for those wanting extra practice, I’ve spent many extra hours making practice
problems. If you think you’ve mastered Hess’s Law, for example, then I fully understand if you don’t do
all of the Hess’s Law examples I’ve included. However, I strongly recommend you do as many math
practice problems as possible to prepare for your test!
Please use the data tables from Chapter 10 of your textbook and the appendix for all calculations. I will
be providing all of that data to you on your test, but please know how to use it!!
When using the data tables and appendix, Don’t forget:
o ΔH°f = ENTHALPY OF FORMATION
o S° = STANDARD ENTROPY
o ΔG°f = STANDARD FREE ENERGY OF FORMATION
o CP = MOLAR HEAT CAPACITY AT CONSTANT PRESSURE
If you can do every single problem on this review sheet, it will be pretty hard NOT to get an “A” on the
test, and almost impossible to score lower than a “B.” To legitimately ACE the test, also do the
following:
o Study your notes packet
o Look over all in-class problems and worksheets (e.g., Hess’s Law problems, molar heat capacity,
etc.)
o Go through your textbook and pay special attention to all equations. I will NOT be giving you
equations on this test, so know them!!
Review Problems:
1) Define enthalpy. Your notes definition works very well on the test, FYI…
2) What are the differences between HEAT and TEMPERATURE?
3) What does the symbol “Δ” mean in physical science?
4) What is the equation for converting Celsius temperatures to Kelvin temperatures? Know the equation
AND be able to convert values:
a. 32°F = __________ °C = _____________ K
b. 212°F = _________°C = _____________ K
c. 22.1°C = _________ K
d. -11.5°C = _________ K
e. 298.15 K = ________°C
f. 155 K = _________ °C
g. 0 K = _________ °C
5) What is considered the “standard” thermodynamic temperature? Express this temperature in both Kelvin
and degrees Celsius.
6) For the expression q = ncΔT, be able to identify what q, n, c, and ΔT represent. ALSO know all of the
units typically associated with each quantity (e.g., T = K), and know how to solve the equation for all of
the variables (e.g., solve the equation for n).
7) The molar heat capacity of tin (Sn) is 27.0 J/mol*K. How much energy does a 150.0g sample of tin
(MW = 118.7 g/mol) gain if it is heated from 21.2°C to 89.3°C?
8) The molar heat capacity of tin (Sn) is 27.0 J/mol*K. How much energy does a 150.0g sample of tin
(MW = 118.7 g/mol) LOSE if it is COOLED from 89.3°C to 0.00°C?
9) 7.40 moles of liquid ethanol is heated from 5.9°C to 78.0°C, and in the process the sample absorbed
59.76 kJ of energy. What is the molar heat capacity of liquid ethanol? Give this molar heat capacity in
J/mol*K.
10) A sample of 4.7 moles of nitrogen gas releases 3176 J of energy as it is cooled. If nitrogen gas has a
molar heat capacity of 29.1 J/mol*K, and the initial temperature of the gas was 34.0°C, what was the
final temperature of the gas?
11) A sample of liquid mercury metal (C = 27.8 J/mol*K) is heated from 1.00°C to 100.0°C, and in the
process it absorbs 6.925 kJ of energy. What is the mass, in grams, of the mercury sample? Note that the
molar mass of Hg = 200.59 g/mol.
12) What is the study of thermodynamics? What topics are covered by the study of thermodynamics?
13) What is the equation for molar enthalpy change? What does each variable represent?
14) What is the molar enthalpy change when a block of titanium is heated by 54.0K, knowing that the molar
heat capacity of titanium is 25.0 J/mol*K? Your final unit should be in J/mol whenever we are finding
“molar enthalpy change.”
15) What is the molar enthalpy change when a block of titanium is cooled by 54.0K, knowing that the molar
heat capacity of titanium is 25.0 J/mol*K?
16) You find a solid metal sample made entirely of a pure substance, and you want to determine which
substance it is. You heat a 1.000 mole sample of the metal to a temperature of exactly 84.96°C, and then
you cool it gradually to a temperature of 52.75°C. During this process, the metal sample released
847.7672 J of energy. From the table below, which metal is your sample made of?
Name of metal
C (J/k*mol)
Potassium (s)
29.58
Magnesium (s)
24.89
Iron(s)
25.10
Beryllium(s)
16.44
Barium (s)
28.07
Calcium (s)
25.31
Titanium(s)
25.02
Manganese (s)
26.32
Tungsten (s)
24.27
Chromium (s)
23.35
Lead (s)
26.44
Nickel (s)
26.07
*Note: All values for this table were taken from “Principles of Modern Chemistry: 4th Edition”
17) Given the table in the previous question, determine the molar change in enthalpy for a 1.000 mole
sample of magnesium metal as it’s cooled from 36.22°C to 28.09°C.
18) If the molar change in enthalpy for a sample of solid calcium is -639.0775 J/mol, what was the samples
change in temperature? (use the C value given within the table in problem 16).
19) If the enthalpy of 1 mol of a compound decreases by 350 J when it is cooled by 15.554 K, what is its
molar heat capacity?
20) Explain how a calorimeter works.
21) Is the energy of the products HIGHER or LOWER than the energy of the reactants if an EXOTHERMIC
reaction occurs?
22) Is the energy of the products HIGHER or LOWER than the energy of the reactants if an
ENDOTHERMIC reaction occurs?
23) If the total enthalpy of the products in a chemical reaction is 854 kJ, and the total enthalpy of the
reactants is 989 kJ, what is ΔH for the chemical reaction?
24) What is the equation for the change in enthalpy for a chemical reaction? Be able to state this equation in
words, too!!
25) Define entropy.
26) Define Gibb’s Free Energy. Write the equation for Gibb’s free energy and be able to label all of the
symbols, including proper units.
27) Be able to identify factors that INCREASE entropy (e.g., changes in temperature, state)
28) Is a mixture of many gases more or less disordered than a mixture of just a single gas?
29) For which process below would you expect entropy to INCREASE? WHY?
a. Na+(g) + Cl-(g)  NaCl(s)
b. 2H2(g) + O2(g)  2H2O(l)
c. Zn(s)  Zn(l)
30) A reaction is considered spontaneous when ΔG is ___________________ (either positive or negative?)
31) Calculate ΔG for a reaction that occurs at 315K, with ΔH = -223 kJ/mol and ΔS = +3.11 kJ/mol*K. Is
this reaction spontaneous?
32) Calculate ΔG for a reaction that occurs at 285K, with ΔH = +396 kJ/mol and ΔS = +0.11 kJ/mol*K. Is
this reaction spontaneous?
33) For a process in which ΔH = -315kJ and ΔS = 32.5 J/K, calculate the change in free energy at -5.3°C.
Express your answer in both J and kJ.
34) Considering the equation for Gibb’s free energy, does entropy play a larger role in determining the
Gibbs energy of a reaction at HIGHER temperatures or at LOWER temperatures? How do you know?
35) What is the molar change in enthalpy when 1.00kg of ice is heated from -19.55°C to -3.88°C? Note that
Cice = 37.4 j/mol*K.
36) Using Hess’s Law, calculate the change in enthalpy for the reaction C(s, gr) + ½ O2 (g)  CO(g) given
the information below. PS…if you need practice with Hess’s Law, there is PLENTY to come!!
C(s, gr) + O2(g)  CO2(g)
ΔH = -393.5 kJ
CO(g) + ½ O2(g)  CO2(g)
ΔH = -283.0 kJ
37) The alkali metals react with chlorine gas to give salts, as shown in the reactions below:
2Li(s) + Cl2(g)  2LiCl(s)
2Na(s) + Cl2(g)  2NaCl(s)
2K(s) + Cl2(g)  2KCl(s)
2Rb(s) + Cl2(g)  2RbCl(s)
2Cs(s) + Cl2(g)  2CsCl(s)
Using the back of your book, calculate ΔSo for each reaction. Is there a trend that’s noticeable?
38) Calculate the enthalpy change for the following reaction using enthalpy data provided in the back of
your book. Classify the reaction as either endothermic or exothermic. Remember, you should be using
the ΔH°f values for ENTHALPY changes, and S° values for ENTROPY changes!
a. 2H2S(g) + 3O2(g)  2H2O(l) + 2SO2(g)
b. MgCl2(s) + H2O(l)  MgO(s) + 2HCl (g)
c. CaO(s) + H2O(l)  Ca(OH)2 (s)
d. From part C, calculate how much heat is released when 100 grams of calcium oxide react with
liquid water to form calcium hydroxide.
e. MgO(s) + H2O(l)  Mg(OH)2(s)
39) Using the data tables in your book for standard free energies of formation, determine the free energy
change in the following reaction and report whether the reaction will proceed spontaneously:
NO(g) + H2O(l) NO2(g) + H2(g)
40) For the following reaction, ΔG is known to be -16.6 kilocalories (calories are just another unit for
energy, just as the Joule is, so don’t freak out!!). If I have a solid jar of NO, O2, and NO2 gases mixed
together, what will happen?
2NO(g) + O2(g)  2NO2(g)
ΔG = -16.6 kcal
a. NO and O2 gas will react to form more NO2 (reaction proceeds spontaneously)
b. More NO and O2 will form (reaction proceeds in the opposite direction)
c. No reaction occurs spontaneously
41) Given the data table below, calculate the standard entropy change for the reaction below at 25.0°C:
Substance
CaCl2(s)
Ca2+(aq)
Cl-(aq)
CaCl2(s)  Ca2+(aq) + 2Cl-(aq)
S° (25.0°C) in J/K*mol
+104.6
-53.1
+56.5
42) Boron Trichloride (BCl3) is the starting material in the manufacturing of elemental Boron. It can be
produced by reacting Diborane (B2H6) with chlorine gas. This reaction also produces hydrochloric acid
in the gaseous state. Boron trichloride can also be produced by reacting boric acid and hydrochloric
acid, as shown in the reaction below. Given the following chemical equations and ΔH values, determine
the change in enthalpy (give your answer in kJ) for the reaction:
B2H6(g) + 6 Cl2(g)
2 BCl3(g) + 6 HCl(g)
ΔH = ???
BCl3(g) + 3H2O(l)  H3BO3(s) + 3HCl(g)
ΔH = -112.5 kJ
HCl(g)  ½ H2(g) + ½ Cl2(g)
ΔH = +92.3 kJ
B2H6(g) + 6 H2O (l)  2H3BO3(s) + 6H2(g)
ΔH = -493.4 kJ
43) Phosphoryl Chloride (POCl3) is used in the chemical manufacturing of flame retardants, as well as in a
variety of important organic chemical reactions. Given the following chemical equations and ΔH values,
determine the change in enthalpy (give your answer in kJ) for the reaction:
P4O10(g) + 6 PCl5(g)
10 POCl3(g)
ΔH = ???
¼ P4(s) + 3/2 Cl2(g)  PCl3(g)
ΔH = -306.4 kJ
P4(s) + 5 O2(g)  P4O10(g)
ΔH = -2967.3 kJ
PCl3(g) + Cl2(g)  PCl5(g)
ΔH = -84.2 kJ
PCl3(g) + ½ O2(g) POCl3(g)
ΔH = -285.7 kJ
44) Calculate the heat of combustion of propane gas (C3H8) using the following information:
Write a balanced equation for combustion of propane! (note that combustion reactions always
produce water vapor and carbon dioxide gas!):
3C(s, gr) + 4H2(g)  C3H8 (g)
ΔH = -103.8 kJ
C(s, gr) + O2  CO2 (g)
ΔH = -393.5 kJ
H2(g) + ½ O2(g)  H2O(g)
ΔH = -241.8 kJ
45) Dinitrogen Pentoxide (N2O5) is a hazardous chemical that can form explosive mixtures with ammonium
salts, and it’s also used as a reagent in chloroform (fyi, a “reagent” is basically a substance used to
produce other substances). Dinitrogen pentoxide can be prepared by reacting nitrogen gas and oxygen
gas. A secondary means of preparation involves nitric acid. Given the following chemical equations and
ΔH values, determine the change in enthalpy for the reaction:
2 N2(g) + 5 O2(g) 2 N2O5(g)
ΔH = ???
2H2(g) + O2(g)  2H2O (l)
ΔH = -571.6 kJ
N2O5(g) + H2O(l)  2HNO3 (l)
ΔH = -73.7 kJ
N2(g) + 3O2(g) + H2(g)  2HNO3 (l)
ΔH = -348.2 kJ
46) Hydrazine (N2H4) is highly toxic and dangerously unstable unless it’s kept in a solution. It is often used
in the manufacturing of polymer foams and pharmaceutical drugs. Additionally, it’s used to make rocket
fuels and to make the chemicals responsible for deploying airbags in cars. Hydrazine can be prepared by
reacting nitrogen gas with liquid water in a reaction that also produces hydrogen peroxide. Hydrazine
can also be produced by reacting different quantities of nitrogen gas and water in a reaction that does not
produce hydrogen peroxide. Given the following chemical equations and ΔH values, determine the
change in enthalpy for the reaction:
N2(g) + 4H2O(l)  N2H4(l) + 2H2O2(g)
ΔH = ???
H2O2(l)  H2(g) + O2(g)
ΔH = +187.8 kJ
H2(g) + ½ O2(g)  H2O (l)
ΔH = -285.8 kJ
N2(g) + 2H2O(l)  N2H4(l) + O2(g)
ΔH = +622.3 kJ
47) Calcium sulfate is a very common chemical with a wide variety of uses and applications. In differing
hydrate forms, it is known as “plaster of paris,” and it also occurs naturally in a mineral form known as
gypsum. Other common uses include drywall (in building/construction), and it’s even used as a
coagulant in foods such as tofu. Given the following chemical equations and ΔH values, determine the
change in enthalpy for the reaction:
CaO (s) + SO3 (g)  CaSO4 (s)
ΔH = ???
H2(g) + ½ O2(g)  H2O (l)
ΔH = -285.8 kJ
Ca(s) + ½ O2(g)  CaO (s)
ΔH = -634.9 kJ
Ca(s) + H2SO4(l)  CaSO4(s) + H2(g)
ΔH = -602.5 kJ
H2SO4(l)  SO3(g) + H2O(l)
ΔH = +132.5 kJ
48) Do I seriously need to type you background information on why glucose is important?!? I’m getting
lazy, so how about you just calculate the change in enthalpy for the formation of glucose:
6C(s, gr)+6H2(g)+3O2(g)C6H12O6(s)
using the following three equations:
C(s, gr) + O2(g)CO2 (g)
H2 (g)+ ½ O2(g) H2O(l)
C6H12O6(s) + 6O2 (g) 6CO2 (g)+ 6H2O(l)
H= -393.51 kJ
H= -285.83 kJ
H= -2803.02 kJ