Unit 2: Equilibria and Thermodynamics
Chem 1B
Name__ANSWERS____
Chapters 16, 6, 17
Equilibrium Problems
1.
Of what does a buffer consist and why are buffers useful? Buffers are a special case of
the common ion effect, calculate the change in pH of a) 0.020 mole HCl added to 1.0
liter of water, and b) 0.020 mole HCl added to a 1.0 L solution of 0.10 M NaF and 0.10
M HF. Ka of HF = 7.2 x 10-4
Buffers consist of a weak acid and a salt that has its conjugate base or a weak
base and a salt that has its conjugate acid. Ratios are between 1:10 and 10:1.
Relatively high concentrations of each will make a stronger buffer capacity.
Buffers resist change in pH even when a strong acid or base is added.
a) change in pH after adding 0.020 mole of HCl to 1.0 L H2O = 5.30
pH 7.00 for neutral water to pH 1.70 for 0.020M [H+]
b) change in pH after adding 0.020 mole of HCl to buffer solution = 0.18
pH 3.14 for the buffer with 0.10 M HF and 0.10M F-1 to pH 2.96 for the
buffer with 0.12M HF and 0.08M F-1
2.
What is more basic a) 1.0 M NH3 solution or b) 1.0 M NH3 and 1.0 M NH4Cl solution?
Explain your answer before any calculations are done. Given that K b = 1.8 x 10-5
calculate the pH of each solution.
The 1.0 M NH3 solution in (a) is more basic since it contains only the weak
base and none of the conjugate acid. When a salt containing the conjugate
acid is added the pH will go down (lower pH, more acidic and closer to
neutral)
(a) pH = 11.63
(b) pH = 9.26
3.
Choose an appropriate weak acid and conjugate base pair to prepare 1.0 L of buffered
solution with a pH of 4.50. What molarity of the conjugate base is needed if 0.10 M of
the weak acid is used?
Several weak acids may be appropriate, I chose acetic acid, HC2H3O2, whose
pKa = 4.74.
molarity of NaC2H3O2 required will be 0.058 M
4.
a)
Describe or draw a simplified picture of a strong acid/strong base titration curve
where you start with 0.10 M NaOH, the y axis is pH and the x axis is the volume of
strong acid added.
b)
Describe or draw a picture for a weak acid/strong base titration curve. As the
strong base is added what initially happens? Is the equivalence point on the basic or
acidic side? after equivalence what happens? Be able to find pH when given initial
concentration of contents in a flask and the quantity of various amounts of solution
added from the buret and any time.
SA + SB
a)
WA + SB
b)
13
pH
equivalence
point at 7.0
equivalence point
above 7.0
pH
buffered
ml of acid added
ml of base added
In part b) a buffer is initially made when adding the strong base to a weak acid.
The equivalence point is basic since there is a conjugate base from the weak acid
that is not neglible. after the equivalent point the curve loses its buffer and a
sharp rise in pH occurs as in the strong acid strong base curve.
5.
An acid-base titration experiment begins with 10.0 ml of 0.30 M NH3 in the flask and
0.10 M HCl in the burette. Kb for the ammonia = 1.8 x 10-5 .
a)
What is the initial [OH-], pOH, and pH for the 0.30 M ammonia solution in the
flask before any acid is added? [OH-]=2.3 x 10-3 M, pOH = 2.63, pH= 11.37
b)
What is the pOH and the pH after 10.0 ml of 0.10 ml of HCl has been added to
the ammonia in the flask?
pOH = 4.44, pH= 9.56
The half equivalence pOH = 4.74, pH = 9.26 (not asked here)
c)
What total volume in ml of the 0.10 M HCl must be added to the 10.0 ml of 0.30
M NH3 solution to reach the equivalence point? 30.0 ml
d)
What is the Ka for NH4 ? Ka = 5.56 x 10-10 for NH4+1
e)
What is the numerical value of the pH at the equivalence point when the solution
contains only the salt and water? pH= 5.19 Is this basic, acidic or neutral?
acidic
f)
Which is the best indicator to use for this titration methyl orange (pH range 3-5),
or phenolphthalein (pH range 8-10)? methyl orange (pH range 3-5)
+
g)
6.
Draw a graph where pH is the y-axis and ml of HCl added is the x-axis. Sketch
the titration curve including points of the previously calculated numbers from
parts above. Indicate where there is a buffered solution if relevant. Inverse
of 4b
a)
If you would like to make a buffered solution with a pH of 3.60, which one of
the following weak electrolyte--salt combinations is best to make the buffer?
1)
2)
H2SO3 and NaHSO3
CH3NH2 and CH3NH3Br
3)
4)
HF and NaF
H3AsO4 and NaCl
Ka of H2SO3 =
1.2 x 10-2
Kb of CH3NH2 = 5.0 x 10-4
Ka of HF =
7.2 x 10-4
Ka of H3AsO4 = 2.5 x 10-4
-4
b)
#3 HF and NaF, Ka of HF = 7.2 x 10 , pKa = 3.14 (close to the desired
pH) and the acid is monoprotic.
Using your chosen combination, what molarity of the weak acid is needed to
make this a buffer of pH 3.60 if the conjugate base used has a concentration of
0.75 M?
0.26 M HF
7.
Choose an appropriate weak acid and conjugate base pair to prepare 1.0 L of buffered
solution with a pH of 4.50. What molarity of the conjugate base is needed if 0.12 M of
the weak acid is used?
Repeat of question #3 with a change in concentration.
Several weak acids
may be appropriate, I chose acetic acid, HC2H3O2, whose pKa = 4.74. The
molarity of NaC2H3O2 required will be 0.069 M
8.
If 2.0 mol NH3 and 1.50 mol Ca(NO3)2 are added to enough water to make 1.0 liter
-5
-6
solution, will Ca(OH)2 precipitate out? Kb = 1.8 x 10 for NH3 and Ksp = 6.5 x 10
for Ca(OH)2.
If a precipitate forms, how many moles of NH4Cl must be added to prevent
precipitation of the calcium hydroxide?
[OH-1] from 2.0 M NH3 = 6.0 x 10-3 M OH-1
[Ca+2] from Ca(NO3)2 = 1.50 M
Qsp for Ca(OH)2 = 5.4 x 10-5
Qsp>Ksp so solid Ca(OH)2 does form, but you cannot see it with the human eye
since the Qsp/Ksp ratio is less than 1000
Maximum of 2.1 x 10-3 M OH-1 allowed to avoid precipitation
A minimum of 0.017 moles of NH4Cl must be added
9.
Experimental evidence shows 1.00 liter of saturated CaF2 solution contains 0.0167 g
CaF2 at 25°C calculate the molar solubility and the Ksp.
0.0167g/L /x(1mol/78.08 g) = 2.1 x 10-4 M
Ksp = (x) (2x)2 = 3.9 x 10-11
10.
Calculate the molar solubility of Mg3(AsO4)2, [Mg ], and [AsO4 ]. Ksp= 2.1 x 10
Ksp = (3x)3 (2x)2 = 2.1x 10-20 = 108x5
Molar solubility x = 4.5 x 10-5M
[Mg+2] = 1.4 x 10-4 M;
[AsO4-3] = 9.0 x 10-5M
11.
Adding a common ion decreases the solubility. Ksp for lead (II) chloride, 1.7 x 10 .
Calculate and compare the molar solubility, (x), of PbCl2 for
a) a saturated solution of PbCl2 in just water
(x) = 0.016 M
b) A solution of 1.0 M Pb(NO3)2 in which excess solid PbCl2 is added to saturate.
Molar solubility = 2.1 x 10-3 M
12.
What is Kd? What is Kf? How are they different from Ksp?
+2
-3
-20
-5
Kd is the dissociation constant of a complex ion
Kf is the formation constant of a complex ion and is the inverse of Kd
Ksp is for the dissociation of a solid and the solid activity is 1
13.
Calculate the concentration of ammonia and cobalt(II) ions in 0.30 M [Co(NH3)6]+2
solution. Kd=1.3 x 10-5
Kd = [Co+2][NH3]6/[Co(NH3)6+2]= (x)(6x)6/(0.30)
Molar solubility and [Co+2] = 0.036 M
6x = [NH3] = 0.22 M
14.
An acid-base titration experiment begins with 20.0 ml of 0.300 M HF in the flask and
0.150 M NaOH in the buret. Ka for HF = 6.8 x 10-4.
a) Draw the expected titration graph where pH is the y axis and ml of NaOH added
is the x axis, Sketch and label the titration curve including WA or WB, SA or
SB, salt, equivalence pt, pH at equiv pt <, =, or > 7, buffer, where color change
of indicator should occur.
See picture in 4b
b) What are the initial [H+], [OH-], pH, pOH, and % ionization for the 0.300 M HF
in the flask? Clearly identify answers. Ka for HF = 6.8 x 10-4 [H+] = 0.014 M;
[OH-] = 7.0 x 10-13 M , pH = 1.85, pOH = 12.15, and 4.7 % ionization
c) What is the pH of the titrated solution at the half equivalence point? What
volume of NaOH must be added to reach this point?
pH = pKa = 3.17 volume = 20.0 ml
d) What is the pH after 18.0 ml of 0.150 M NaOH has been added to the flask
containing 20.0 ml of 0.300 M HF?
pH = 3.08
e) What volume of 0.150 M NaOH must be added to the original flask of 20.0 ml
of 0.300 M HF to reach the equivalence point?
volume = 40.0 ml
f) Write the balanced whole, complete ionic, and net ionic equations of hydrolysis
of the salt sodium fluoride (NaF) and indicate whether the salt is acidic, basic, or
neutral.
All species aqueous unless otherwise indicated
NaF + H2O (l)  NaOH + HF
Na+1 + F- + H2O (l)  Na+1 + OH- + HF
F- + H2O (l)  OH- + HF
Basic
g) Solve for the numerical Kb value for the net ionic equation in the previous step.
Kb = 1.5 x 10-11
h) What is the pH of the solution in the flask at the equivalence point?
[F-1] = 6.0 mmol/(20.0+40.0 ml) = 0.10 M F-1
[OH-1] = 1.2 x 10-6 M ; pOH = 5.91; pH = 8.09
i) Which is the best indicator to use for determining the endpoint of this titration
Bromocresol green (pH 4-6) or phenolphthalein (pH 8-10)?
phenolphthalein (pH 8-10)
j) What is the pH after 45.0 ml of 0.150 M NaOH has been added to the original
flask containing 20.0 ml of 0.300 M HF?
0.75 mmol OH-1/65.0 ml = 0.0115 M [OH-1]; pOH = 1.94; pH = 12.06
k) Check your pH answers and compare to the expected pH titration curve you
predicted in part (a). Fill in your calculated values.
Thermochemistry Problems
15.
Give a brief description of each of the laws of thermodynamics. (Zeroth, first, second,
third)
Laws of thermodynamics are useful in predicting outcomes, but unlike a theory
do not explain the expected behavior.
Zeroth Law: Temperature Concept: Temperature measures the intensity of hotness or
coldness of an object. When two objects are brought together heat always
flows spontaneously from a hotter object to a colder object until thermal
equilibrium is reached. Think of the heat flow direction when placing
your hand in ice water verses hot water.
First Law: Law of Conservation of Energy: Energy cannot be created or destroyed.
The energy of the universe is constant.
This law works for ordinary chemical and physical processes. It was
later combined with another (conservation of matter) to include nuclear
reactions after E =mc2. It was changed to the Law of Conservation of
Matter and Energy
Second Law: Entropy Principle: Every spontaneous change increases the entropy of
the universe. This does not mean local decreases in entropy cannot take
place. For the entropy of a system to decrease, the entropy of the
surroundings must increase more.
Third Law: Zero Entropy Established: The entropy of a pure, perfect, crystalline
substance (no disorder) is zero at the temperature of absolute zero, 0
Kelvin.
It is impossible to reach absolute zero. It has been tried in the lab and
gets close.
16.
A certain system change releases 3462 J of heat and at 1.00 atm pressure expands its
volume from 12.9 L to 46.8 L. Solve for the E, H, and w for the system. (include
the correct sign +/-, and units)
.
H = -3462 J, w = -33.9 L atm must convert to joules = -3434 J,
17.
When gasoline burns in an automobile engine, the heat released causes the CO 2 and
H2O vapor (this is gas) produced to expand, which pushes the pistons outward.
Excess heat is removed by the car's cooling system. Determine the change in internal
energy if the expanding gases do 451J of work on the pistons and the system looses
325 J to the surroundings as heat.
H = -325 J, w = -451 J,
18.
E = -6896 J
E = -776 J
Before the days of refrigerators and icemakers, drinking water in hot countries was
often cooled by storage in clay pots. Evaporation of water through the pot cooled the
remaining water. How many grams of water must evaporate to cool 3790 grams (1
gallon) of water from 34.8°C to 21.0°C? For water, specific heat is 4.184 J/g°C; Hvap
is 44.0 kJ/mol.
Qrxn = -218.8 kJ heat removed from the 1 gallon of water
4.97 moles of water, 89.5 g of water must evaporate
19.
Is there a difference between the heat (q) determined in a coffee-cup calorimeter and
the heat determined in a bomb calorimeter? Explain your answer.
Heat in a coffee-cup calorimeter is under constant pressure and one measures
change in enthalpy, in a bomb calorimeter the volume is constant and you
measure internal energy change.
20.
50.0 ml of 1.00 M NaOH (aq) at 25.4 °C was added to 50.0 ml of a 0.50 M H 2SO4
(aq) solution also at 25.4°C in a coffee-cup calorimeter with a heat capacity of 20 J/°C.
The 2 solutions were mixed thoroughly and the temperature rose to 31.9°C. What is
the heat of neutralization of 1 mole of H2SO4?
(Assume the density of both liquids to be 1.00 g/ml and the specific heat to be 4.184
J/g°C)
Qcal + water = 2719.6 J + 130 J = 2849.6J, reverse sign for the reaction and divide by
moles of H2SO4
H = -114 kJ/mol H2SO4
21.
A manufacturer claims that its new dietetic dessert has fewer than 10 Calories per
serving. An independent laboratory uses a bomb calorimeter with a heat capacity of
8.151 kJ/K to test the claim. When one serving of the dessert is burned in excess O 2,
the temperature rises from 22.153 °C to 27.080°C. Is the manufacturer's claim correct?
YES
22.
The major source of Al in the world is bauxite (mostly Al2O3)
Al2O3 (s)  2Al (s) + 3/2 O2 (g);
Hrxn = 1676 kJ
How many grams of Al can be produced utilizing 1000. kJ of heat? 32.2 g Al
23.
Using Hess' Law and the 2 equations below, solve for the Hrxn of
P4 (s) + 10 Cl2 (g)  4 PCl5 (s)
1)
P4 (s) + 6Cl2 (g)  4 PCl3 (l)
H = -1277.2 kJ
2)
PCl5 (s)  Cl2 (g) + PCl3 (l)
H = +123.8kJ
H = -1772.4 kJ
24.
x1
x -4
Write the balanced formation equation of solid aluminum bicarbonate Al(HCO3)3.
Al (s) + 3/2 H2 (g) + 3 Cgraphite (s) + 9/2 O2 (g)  Al(HCO3)3 (s)
25.
What is a state function? Give several examples of both state functions and those that
are not.
State Function: a property that depends only on the present state and not the
path the system took to reach the state. In Thermochemistry state functions are
generally capitalized. Some examples of state functions include: E, H, V, P, T.
Time and work are not state functions.
26.
Why is it important to show the physical phase in thermochemical equations?
Different states of matter have different energies.
27.
The first step in the production of nitric acid is the oxidation of ammonia.
4NH3 (g) + 5O2 (g)  4NO (g) + 6H2O (g)
Calculate Hrxn from the Hf values found in the appendix.
Hrxn = 4(90.37) + 6(-241.82) – 4(-46.19) – 5(0) = -904.68 kJ/mol
28.
A 48.3 g sample of NH3 is produced according to the following reaction done under
constant pressure and 450°C. Solve for the amount of work done on the system.
N2 (g) + 3 H2 (g)  2 NH3 (g)
w = - nRT = +12.0 kJ/mole reaction as written, for 48.3 g NH3, work = 17.0 kJ
29.
Estimate the heat of reaction below using the bond energies given :
CH4 (g) + 3Cl2 (g)  CHCl3 (g) + 3 HCl (g)
Bond Energies: C - H = 414 kJ/mol: Cl-Cl = 243 kJ/mol: C-Cl = 330 kJ/mol: H Cl = 431 kJ/mol
Hrxn = -312 kJ/mol rxn
30.
Given the Hf° of ammonia gas is - 46.11 kJ/mol, and the following bond energies:
N - H is 389 kJ; H - H is 435 kJ/mol. Estimate the bond energy of a nitrogen
nitrogen triple bond.
Bond Energy for the triple bond in N2 = 937 kJ/mol
31.
List and explain the factors (such as phase changes) that help in predicting the sign of
the change in entropy of a system. Try to think of a type of example for each. (these
factors interrelate so the same example may be used more than once with a combined
explanation.)
Predicting the sign for S:
1) Phase changes: s  l  g… + S H2O (l)  H2O (g)
2) Temp: warming…
+ S O2 (g) 25°C  O2 (g) 80°C
3) Volume: expanding…
+ S 10 ml gas  1.0 liter of same gas
4) Mixing: dissolving ...
+ S NaCl (s) + H2O (l)  NaCl (aq)
5) Rearrange to create more ways of being…
+ S H2 (g) + Cl2 (g)  2 HCl (g)
6) # parts: increase…
+ S 10 g rock  10 g sand
7) More moles gas…
+ S N2O4 (g)  2 NO2 (g)
32.
Choose the one with the higher entropy in each of the following pairs and justify your
choice.
a) 1 mole of SO2 gas
or
1 mole of SO3 gas
more atoms
b) 1 mole CO2 solid
or
1 mole CO2 gas
gas vs solid
c) 3 moles O2 gas
or
2 moles O3 gas
more moles gas
d) 1 mole KBr crystal or
1 mole KBr aqueous
more mixed
e) Sea water at 2°C
or
Sea water at 23°C
higher temp
f) 1 mole HF gas
or
1 mole HI gas
more subatomic particles
33.
Calculate S°rxn for the combustion of 1 mole of methane, CH4, at 25.0°C using S° in
Appendix
S°rxn = nS°products nS°reactants
CH4 (g) + 2 O2 (g)  CO2 (g) +2 H2O (l)
S°rxn = 213.6 + 2(69.91) – 186.3 – 2(205) = -242.9 J/mol K
34.
Estimate the normal boiling point (the temperature at which liquid and vapor are in
equilibrium with each other and G = 0) for CCl4 at 1 atm pressure. ( H° and S° do
not vary much with temperature, so often they may be used to estimate G at
temperatures other than 25°C)
CCl4(l)  CCl4 (g)
H°= 32.6 kJ/mol, S° = 95.0 J/mol K, Teq= 343K, 70°C
35.
Use H°f and S° values to calculate G° at 25 °C for the reaction
2 CH3OH (l) + 3 O2 (g)  2 CO2 (g) + 4 H2O (l)
Under which conditions is this reaction spontaneous?
H°= -1453.1 kJ/mol, S° = -161.76 J/mol K, G° = -1404.9 kJ/mol,
Teq= 8983 K, 8710 °C, This reaction is spontaneous under 8710°C, in other
words, for all reasonable temperatures this will be a spontaneous reaction.
36.
Answer positive, negative or zero for each blank
Exothermic reactions have a negative enthalpy.
The change in entropy of a spontaneous reaction is usually positive , this means more
disorder.
A reaction that is nonspontaneous has a
positive
G, while a system at
equilibrium has a zero G, and a spontaneous reaction has a negative
G.
37.
What is the value of Keq when G = -98.0 kJ/mol at 25°C?
Keq = 1.5 x 1017
38.
In the Haber process NH3 is manufactured from N2 and H2. At 25°C, Kp = 3.6 x 108.
To increase the rate of reaction the reaction is run at a higher temperature. Use the
van't Hoff equation to calculate Kp at 450°C
N2 (g) + 3 H2 (g)  2NH3 (g)
H° = -92.2 kJ
Van't Hoff eq. ln(KT2/KT1) = H/R (1/T1 - 1/T2)
At 450°C, Kp = 0.11
39.
For the following question on this problem refer to the equation below.
2 Cl2 (g) + 2 H2O (g)  4 HCl (g) + O2 (g) ; H° = +115 kJ
At 25°C, the equilibrium partial pressures in a 5.00 liter container are as follows;
[P(Cl2)] = 3.1 atm, [P(H2O)] = 3.1 atm, [P(HCl)] = 1.3 x 10-3 atm and [P(O2)] = 1.5
atm.
a)
Calculate Kp at 25°C
Kp = 4.6 x 10-14
b)
c)
d)
e)
40.
Estimate G° from the Kp value
G° = 76.1 kJ/mol
Using the values of Kp and G° explain your prediction of whether products or
reactants are favored. Reactants are favored, forward is nonspontaneous
Is this reaction spontaneous at 25°C? NO
You solved for Kp in the original reaction at 25°C . Use the van't Hoff equation
to estimate Kp at 400°C. Kp = 7.9 x 10-3
The reaction for the enthalpy of formation of H2O (g) has H°f = -242 kJ/mol.
a) Write the balanced formation equation.
H2 (g) +1/2 O2 (g)  H2O (g)
b)
Is the change in entropy expected to be positive or negative. Explain
Negative, since fewer moles of gas are in products.
c)
How much PV work is done in kJ/mol and what is the value of E in kJ/mole
for the formation equation of H2O (g) formed, given that when 0.50 moles of
H2 gas reacts with 0.25 mol of O2 gas at 1.00 atmospheric pressure the
volume compresses by 5.6 L.
Work = -P V = +5.6 L atm = 567 J for 0.50 moles H2
w = 1.13 kJ/mol H2O
E = -242 kJ + 1.13 kJ = -240.9 kJ/mol
d)
Under what conditions will this be spontaneous? Explain.
Spontaneous at temperatures below the equilibrium temp since bothe
enthalpy and entropy changes are negative.
41.
Specific Heat
A 7.82 g sample at 99.2 C is added to 30.0g of water at 22.10 C in a coffee-cup
calorimeter. The final temperature is 24.20 C. The specific heat capacity of water is
4.184 J/g C.
a) Solve for the specific heat of the metal
0.449 J/g C
b) Determine which metal it is. (Al = 0.903 J/g C, Fe = 0.449 J/g C, Pb = 0.128 J/g C)
c) Test the Law of Dulong and Petit which states that the molar heat capacity of
elements are all approximately 25 J/mol C by solving for the molar heat capacities of
Al, Fe, Pb.
Al = 0.903 J/g C x 27.0 g/mol = 24.4 J/mol C
Fe = 0.449 J/g C x 55.85 g/mol = 25.1 J/mol C
Pb = 0.128 J/g C x 207.2 g/mol = 26.5 J/mol C
The law seems to work, answers come to approximately 25 J/mol C
42.
Heat of Reaction
50.0 ml of 0.500 M NaOH and 50.0 ml of 0.500 M HC2H3O2 both initially at 20.00 C are
added together in a coffee-cup calorimeter and the final temperature is found to be
23.37 C. Assume the volumes are additive, the density of the solutions are 1.00 g/ml, the
specific heat of the solution is 4.18 J/g C and it is an isolated system in which no matter
or energy is exchanged outside of the calorimeter.
a) Solve for the Qcalorimeter and the Qrxn.
Qcalorimeter = 1409 J , Qrxn = -1409 J
b) Solve for the molar enthalpy of neutralization in kJ/mol, Hrxn
Hrxn= -56.3 kJ/mol
43.
Bomb Calorimeter Calibration
A known compound that generates 15.1 kJ of energy is burned in a bomb calorimeter and
the temperature increased from 21.80 C to 23.28 C.
a) Solve for the Qcalorimer
Qcalorimeter = +15.1 kJ
b) Is this experiment under constant pressure or constant volume? Is this H or U
Constant volume, measuring internal energy, written as E or U
c) Solve for the Heat Capacity of this bomb calorimeter in kJ/ C.
Heat capacity = 10.2 kJ// C
44.
Bomb Calorimeter Problem
For the combustion of 1 mole of ethanol in kJ/mol use the following data and the above
calibrated bomb calorimeter.
1.000 ml of ethanol (C2H5OH) is pipeted into the sample chamber of the bomb
calorimeter and then 25 atm pressure of oxygen are added to insure complete
combustion. The initial temperature is 22.185 C, the final temperature is 24.469 C. The
density of the ethanol is 0.789 g/ml.
a) Write the balanced combustion reaction for 1 mole ethanol
C2H5OH (l) + 3 O2 (g)  2 CO2 (g) +3 H2O (l)
b) Solve for Qcalorimeter and the Qrxn.
Qcalorimeter = 23.3 kJ , Qrxn = -23.3kJ
c) Solve for internal energy
E = -29.5 kJ/g , -1358 kJ/mol
d) Solve for work and enthalpy in kJ/mol under standard conditions.
w = 2.48 kJ/mol,
Hrxn= -1356 kJ/mol
45.
Work
Predict the sign for work, determine if the system is expanding or compressing and state
if work is done on or by the system.
a) 2 NO (g) + O2 (g)

2 NO2 (g)
compression, (+)
b) PCl5 (g)

PCl3 (g) + Cl2 (g)
expansion, (-)
c) 2 C4H10 (g) + 13 O2 (g)  10 H2O (l) + 8 CO2 (g) compression, (+)