Download Binnie Thermochemistry Practice ANSWERS - binnie

Chapter 6 THERMOCHEMISTRY
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Energy = capacity to do work or produce heat
Δ E = q + w (q > 0 heat transferred to system, q<0 heat transferred out of system)
(w>0 work done on system, w<0 work done by system on surroundings)
Energy is conserved (1st Law of thermodynamics)
Energy is a state function (does not depend on history of sample, only present
conditions)
Kinetic Energy = ½ m v2 energy of motion, measured in Joules
Potential Energy = stored energy
System vs. surrounding
Endothermic (+) = system absorbs heat, Exothermic (-) = system releases heat
Work = energy used to make an object move against a force w = F d
Heat = energy transferred from hotter to colder objects, causing temperature change
Enthalpy (H) = heat transfer under constant pressure Δ H = Hfinal – Hinitial = qp
Heat of reaction = enthalpy change that accompanies a reaction
Calorimetry q=mc ΔT (ΔT = Tfinal – Tinitial) measuring heat flow, heat lost by one object
must be gained by another
C = heat capactity (J/g) specific heat capacity (J/g°C) molar heat capacity (J/mol°C)
The magnitude of ΔH is proportional to the quantities of reactants/products
If a reaction is reversed the sign of ΔH is reversed.
Hess’s Law says that since enthalpy is a state function, it can be calculated by a series of
steps that add up to overall reaction
ΔH°f = enthalpy change for formation of one mole of a compound from its elements in
standard state (25°C, 1atm, 1M)
ΔH°f of an element in standard state = 0
H rxo   H of (Pr oducts)   H of Re ac tan ts
1. Classify each statement as talking about an [EXO]thermic or [ENDO]thermic reaction:
_EXO_ surroundings get hot
_EXO_ H is negative
_ENDO
PE diagram is uphill
_EXO_
PE diagram is downhill
_EXO_
energy is a product
_ENDO
surroundings get cold
_ENDO
H is positive
_ENDO
products have more energy
_EXO_
reactants have more energy
_ENDO
energy is a reactant
2. Determine the change in energy, E, for each system:
E=q+w
A system gives off 25.0 kJ of heat and has 15.0 kJ of work done on it. -25 kJ + 15 kJ = -10kJ
A system takes in 75.0 kJ of heat and has 25.0 kJ of work done on it. 75 kJ + 25 kJ = 100kJ
A system does 45.0 kJ of work and loses 80.0 kJ of heat. -80 kJ - 45 kJ = -125kJ
3. Calculate the change in the internal energy of the system for a process in which the system
absorbs 140 J of heat from the surroundings and does 85 J of work on the surroundings
E=q+w
140 J - 85 J = 55 J
4. When two solutions react the container “feels hot.” Thus,
a. the reaction is exothermic
c. the energy of the universe is increased
b. the reaction is endothermic
d. the energy of both the system & surroundings decrease
5. (a) How much heat is needed to warm 250 g of water (about 1 cup) from 22 °C (about room
temperature) to near its boiling point, 98 °C? The specific heat of water is 4.18 J/g-K.
q=mcT
q= (250g) (4.18 J/gK) (98°C - 22°C) = 79420 J
(b) What is the molar heat capacity of water?
Molar heat capacity is the heat capacity for 1 mole of the substance.
( 4.18 J/gK ) (18g/1mol) = 75.2 J/mol K
(c) Describe the direction of heat flow.
From the surroundings (air in the room) to the water
6. A 140.0-g sample of water at 25.0°C is mixed with 100.0 g of a certain metal at 100.0°C.
After thermal equilibrium is established, the (final) temperature of the mixture is 29.6°C.
How would you calculate the specific heat capacity of the metal, assuming it is constant
over the temperature range concerned?
(A)
(100.0 ) ● (29.6 - 25.0)
(C)
(140.0) ● (4.184) ● (100.0 - 29.6)
(140.0) ● (4.184) ● (100.0 - 29.6)
(100.0) ● (29.6 - 25.0)
(B)
(140.0)●(4.184) ● (29.6 - 25.0)
(100.0) ● (29.6-100)
(D)
(140.0) ● (4.184) ● (29.6 - 25.0)
(100.0) ● (29.6)
7. Which of the following is equal to the heat transferred at constant pressure?
a. Work
b. Internal Energy
c. Enthalpy
d. Entropy
e. Free Energy
8. (a) Large beds of rocks are used in some solar-heated homes to store heat. Assume that the
specific heat of the rocks is 0.82 J/g-K. Calculate the quantity of heat absorbed by 50.0 kg of
rocks if their temperature increases by 12.0 °C.
q=mcT
q= (50000 g) (0.82 J/gK) (12°C) = 492000 J
(b) What temperature change would these rocks undergo if they emitted 450 kJ of heat?
q=mcT T = q / mc = (-450000J) / (50000) (0.82 J/gK) = 10.9° drop in temperature
9. When a 94.3 g piece of gold (specific heat 0.13 J/g·°C) and a 94.3g piece of iron (specific heat
0.46 J/g·°C) each at 10OC absorb an equal amount of heat
(A) the gold will end up with the higher final temperature. ΔT = q_
(B) the iron will end up with the higher final temperature.
mc
(C) they will end up at the same final temperature.
10. When a student mixes 50 mL of 1.0 M HCl and 50 mL of 1.0 M NaOH in a coffee-cup
calorimeter, the temperature of the resultant solution increases from 21.0 °C to 27.5 °C.
Calculate the enthalpy change for the reaction in kJ/mol, assuming that the calorimeter
loses only a negligible quantity of heat, that the total volume of the solution is 100 mL, that
its density is 1.0 g/mL, and that its specific heat is 4.18 J/g-K.
Total volume = 100mL and density = 1.0 g/mL, so mass = 100 g
q=mcT
q = (100g) (4.18 J/gK) (6.5K) = 2717 J
moles of solution = 0.05 moles HCl + 0.05 = 0.1 mol solution
2717 J / 0.1 mol = 27170 J/mol = 27.170 kJ/mol
11. If 25 J are required to change the temperature of 5.0 g of substance A by 2.0C, what is the
specific heat of substance A?
a) 250 J/gC
b) 10. J/gC
c) 63 J/gC
d) 2.5 J/gC
12. Calculate the enthalpy change for: C(s) + 1/2O2(g)  CO(g).
Given:
C(s) + O2(g)  CO2(g)
ΔH = -393 kJ/mol
CO(g) + 1/2O2(g)  CO2(g) ΔH = -282 kJ/mol
As is
C(s) + O2(g)  CO2(g)
ΔH = -393 kJ/mol
Reversed CO2(g) CO(g) + 1/2O2(g) ΔH = +282 kJ/mol
Total = - 111 kJ/mol
14. Using the standard enthalpies of formation (you will have to look these values up), calculate
the standard enthalpy change for the reactions:
a. 2KClO3(s)  2KCl(s) + 3O2(g)
∆Hrx = (2xKCl + 3xO2) – (2xKClO3) = (2x-435.87 + 0) – (2x(-391.20)) = -89.34 kJ
b. 2Al(s) + Fe2O3(s)  Al2O3(s) + 2Fe(s)
∆Hrx = (1xAl2O3 + 2xFe) – (2xAl + 1xFe2O3) = (-1669.8 + 0) – (-822.2 + 0) = -847.6 kJ
15. Write standard formation reactions for each of the following:
a. Fe2O3 (s) ____2Fe(s) + 𝟑⁄𝟐O2(g) → Fe2O3(s)__________________________________
b. Mg(NO3)2(s) ___ Mg(s) + N2(g) + 3 O2(g) → Mg(NO3)2(s)__________________________
16. Write standard enthalpy of combustion reactions for the following
a. C3H8 (l) __________ C3H8 (l) + 5O2(g) → 3CO2(g) + 4H2O(l)______________
C6H5OH(s) + 7 O2(g)  6 CO2(g) + 3 H2O(l)
17. When a 2.000-gram sample of pure phenol, C6H5OH(s), is completely burned according to the
equation above, 64.98 kilojoules of heat is released. Use the information in the table below to
answer the questions that follow.
Standard Heat of
Formation, ΔH°f; at
Substance
25°C (kJ/mol)
CO2(g)
-393.5
H2O(l)
C6H5OH(s)
-285.85
?
(a) Calculate the molar heat of combustion of phenol in kilojoules per mole at 25°C.
-64.98 kJ / [2.000 g * (1 mol /94 g)] = -3058
(b) Calculate the standard heat of formation, ΔH°f, of phenol in kilojoules per mole at 25°C.
∆Horxn = ∑∆Hfoproducts - ∑∆Hforeactants
-3058 kJ = (6 ∆HfoCO2 + 3 ∆HfoH20 ) - ( ∆HfoC6H5OH )
-3058 kJ = (6 •-393.5 + 3 • -285.85) - (∆HfoC6H5OH)
∆HfoC6H5OH = -161 kJ/mol
18. Expanding gases, due to the combustion of gasoline, inside an automobile engine
displaces the piston, as diagramed below. Write the sign (+ or -) for each of the
thermodynamic functions. 2 points each.
qsys_-___ wsys _-___ ∆Esurroundings___+_____ ∆Esys____-_____
19. In order to measure the heat capacity of kerosene, a student added 150 g of kerosene at
81.2oC to 100 g water at 17.4oC in a styrofoam cup. After mixing, the temperature was
44.7oC. Assuming negligible heat capacity of the cup and negligible heat losses to the
surroundings, calculate the specific heat of kerosene.
Qwater = (100g)(4.18 J/g°C)( 44.7-17.4°C) = 11411.4 J
ckerosene = _____11411.4 J___
(150g)(81.2-44.7°C)
= 2.1 J/g°C
20. The bombardier beetle uses an explosive discharge as a defensive measure. The chemical
reaction involved is the oxidation of hydroquinone by hydrogen peroxide to produce quinine
and water:
C6H4(OH)2 (aq) + H2O2 (aq)  C6H4O2 (aq) + 2H2O (l)
Calculate the ΔH for this reaction using the following:
C6H4(OH)2 (aq)  C6H4O2 (aq) + H2 (g)
ΔH = +177.4 kJ
H2 (g) + O2 (g)  H2O2 (aq)
ΔH = -191.2 kJ
H2 (g) + ½ O2 (g)  H2O (g)
ΔH = -241.8 kJ
H2O (g)  H2O (l)
ΔH = -43.8 kJ
when you reverse a reaction, also reverse the sign of ΔH
when you double a reaction, also double the ΔH
reaction 1
C6H4(OH)2(aq) -------------------->C6H4O2(aq) + H2(g)
reaction 2 REVERSED H2O2(aq) ----------------------------->H2(g) + O2(g)
reaction 3 DOUBLED 2H2(g) + O2(g) --------------------->2H2O(g)
reaction 4 DOUBLED 2H2O(g) ------------------------------->2H2O(l)
ΔH = +177.4 kJ
ΔH = +191.2 kJ
ΔH = 2(-241.8) kJ
ΔH = 2(-43.8) kJ
Add them all together
ΔH reaction = (+177.4 +191.2 + 2(-241.8) + 2(-43.8)) kJ = -203 kJ/mole hydroquinone
(exothermic)
21. How much heat is released when 34.7 grams of aluminum is completely burned in excess
oxygen. 4Al(s) + 3O2(g) → 2Al2O3(s) ∆H = –3351 kJ/mol
-3351 kJ / 4 moles of reaction = 837.8 kJ per mole of Al
34.7 g (1 mol/27 g) = 1.3 mol Al
1.3mol(-837.8kJ/mol) = - 1089 kJ released
22. For a particular process q = 20 kJ and w = 15 kJ. Which of the following statements is true?
A) Heat flows from the system to the surroundings.
B) The system does work on the surroundings.
C) ΔE = 35 kJ
D) All of the above are true.
E) None of the above are true.
Substance
∆Hof (kJ/mol)
Substance
∆Hof (kJ/mol)
Al2O3 (s)
-1676.0
H2O2(l)
−187.8
CH3OH(l)
-239.0
KCl(s)
-436.7
CO (g)
-110.5
KClO3(s)
−391.4
CO2 (g)
-393.5
NH3 (g)
-46.19
Fe2O3 (s)
-822.1
NO (g)
90.37
H2O(g)
-241.8
NO2 (g)
33.85
H2O(l)
-285.8
O3 (g)
142.0
23. Using the table above calculate the change in enthalpy for the following reaction:
2CH3OH(l) + 3 O2(g) → 2 CO2(g) + 4 H2O(l)
∆H = ?
o
o
o
∆H rxn = ∑∆Hf products - ∑∆Hf reactants
∆Horxn = ( 4 ∆HfoH20 + 2 ∆HfoCO2) - ( 2 ∆HfoCH3OH )
∆Horxn = (4 •-285.8 + 2 • -393.5) - ( 2 • -239.0)
∆Horxn =
-1452.2 kJ/mol
C2H5OH(l) + 3O2(g)  2CO2(g) + 3H2O(l), ∆H = –1.37 x 103 kJ
24. For the combustion of ethyl alcohol as described in the above equation, which of the
following is true?
I. The reaction is exothermic.
II. The enthalpy change would be different if gaseous water was produced.
III. The reaction is not an oxidation–reduction one.
IV. The products of the reaction occupy a larger volume than the reactants.
A) I, II
B) I, II, III
C) I, III, IV
D) III, IV
E) only I
25. On a cold winter day, a steel metal fence post feels colder than a wooden fence post of identical size
because:
A) The specific heat capacity of steel is higher than the specific heat capacity of wood.
B) The specific heat capacity of steel is lower than the specific heat capacity of wood.
C) Steel has the ability to resist a temperature change better than wood.
D) The mass of steel is less than wood so it loses heat faster.
E) Two of the above statements are true.
26. If a student performs an endothermic reaction in a calorimeter, how does the calculated value of ∆H
differ from the actual value if the heat exchanged with the calorimeter is not taken into
account?
A) ∆Hcalc would be more negative because the calorimeter always absorbs heat from the reaction.
B) ∆Hcalc would be less negative because the calorimeter would absorb heat from the reaction.
C) ∆Hcalc would be more positive because the reaction absorbs heat from the calorimeter.
D) ∆Hcalc would be less positive because the reaction absorbs heat from the calorimeter.
E) ∆Hcalc would equal the actual value because the calorimeter does not absorb heat.
27. Which of the following does not have a standard enthalpy of formation equal to zero at 25°C and 1.0
atm?
A) F2(g)
B) Al(s)
C) H2O(l)
D) H2(g)
E) They all have a standard enthalpy equal to zero.