Download Topic 5 Energetics File

Topic 5:Energetics
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Table of Contents
Topic 5:Energetics ......................................................................................................................................... 486
Definitions .................................................................................................................................................... 488
Temperature Conversion ........................................................................................................................ 490
Topic 5.1: Exothermic and Endothermic Reactions. ............................................................................ 490
Endothermic Reactions vs. Exothermic Reactions........................................................................... 490
Energy Diagrams ...................................................................................................................................... 491
Topic 5.2: Calculation of Enthalpy Changes........................................................................................... 496
Heat Transfer Problems ........................................................................................................................ 496
Chemistry Calorimetry Exercises ........................................................................................................ 498
Topic 5.3 Hess’s Law ................................................................................................................................... 499
Notes:Calculations using Hess's Law and Heats of Formation ..................................................... 499
Hess's Law Worksheet .......................................................................................................................... 501
Hess’s Law Challenge Worksheet ........................................................................................................ 504
Topic 5.4 Bond Enthalpies ......................................................................................................................... 506
Bond Enthalpies........................................................................................................................................ 506
Bond Enthalpies 2 .................................................................................................................................... 507
Using Hess’s Law and Bond Energy to calculate enthalpy change. ............................................... 508
Topic 5: Energetics Exam Questions ...................................................................................................... 510
Topic 5: Exam Questions Markscheme .................................................................................................. 517
Topic 15.1 Standard Enthalpy Changes of Reaction HL ...................................................................... 519
Enthalpy of Formation............................................................................................................................ 520
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Definitions
Average bond enthalpy: The average enthalpy change of breaking one mole of a bond in a
gaseous atom into its constituent gaseous atoms.
Born-Haber cycle: Energy cycles for the formation of ionic compounds. If there is little
agreement between the theoretical and experimental values, this could indicate a degree of
covalent character.
Electron affinity: Enthalpy change when an electron is added to an isolated atom in the gaseous
state.
Endothermic: A reaction in which energy is absorbed. ΔH is +. Reactants more stable than
products.
Enthalpy: The internal energy stored in the reactants. Only changes in enthalpy can be
measured.
Entropy: A measure of the disorder of a system. Things causing entropy to increase: 1) increase
of number of moles of gaseous molecules; 2) change of state from solid to liquid or liquid to gas;
3) increase of temperature
Exothermic: A reaction in which energy is evolved. ΔH is –. Products more stable than
reactants.
Gibb’s free energy: Must be negative for reaction to be spontaneous. ΔG = ΔH – TΔS
Hess’ law: Enthalpy change for a reaction depends only on difference between enthalpy of
products and enthalpy of reactants. It is independent of pathway.
Lattice enthalpy: The endothermic process of converting a crystalline solid into its gaseous
ions, or the reverse exothermic process. The lattice enthalpy increases with decreasing size of
the ions and increasing charge.
Spontaneous: A reaction that has a natural tendency to occur.
Standard conditions: 298 K and 1 atm.
Temperature: A measure of the average kinetic energy.
Standard enthalpy of vaporisation: The energy required to vaporise one mole of a liquid.
Enthalpy of atomisation: The energy required to produce one mole of gaseous atoms from an
element in its standard state.
Bond dissociation enthalpy: The energy change when one mole of a specific bond is broken or
created under standard conditions.
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Enthalpy of Combustion: The energy released when one mole of a compound is burned in excess
oxygen.
Standard enthalpy of formation: The energy change when one mole of a compound is formed
under standard conditions from its constituent elements in their standard states.
Standard enthalpy of solution: The energy change when one mole of a substance is dissolved in
an infinite amount of water under standard conditions.
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Temperature Conversion
Ko= Co + 273
Fo = 9/5 Co +32
Co = 5/9(Fo-32)
Convert the following to Fahrenheit
1) 10o C ________
2) 30o C ________
3) 40o C ________
4) 37o C ________
5) 0o C ________
Convert the following to Celsius
6) 32o F ________
7) 45o F ________
8) 70o F ________
9) 80o F ________
10) 90o F ________
11) 212o F ________
Convert the following to Kelvin
12) 0o C ________
13) -50o C ________
14) 90o C ________
15) -20o C ________
Convert the following to Celsius
16) 100o K ________
17) 200o K ________
18) 273o K ________
19) 350o K _________
Topic 5.1: Exothermic and Endothermic Reactions.
Endothermic Reactions vs. Exothermic Reactions
Classify each of the following changes as either exothermic or endothermic. Also, determine
whether each change is a chemical reaction or a physical change. Explain your reasoning for
each.
1.
2.
3.
4.
5.
6.
An ice cube melts after being left out on the table over an extended period of time.
Cracking a raw egg into a hot frying pan and cooking it.
A match burns/ignites after being struck against a rough surface.
The human body uses the energy provided from the digestion of food.
Morning dew forming on grass and plants.
Dynamite explodes in the destruction of a building.
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Energy Diagrams
450
400
350
(d)
300
(a)
Z
Potential
250
Energy
(kJ)
(c)
200
150
(e)
(f)
X+Y
100
(b)
50
Reaction Coordinate
1.
What is the potential energy of the products?
2.
What is the potential energy of the activated complex?
3.
What is the potential energy of the reactants?
4.
How much activation energy is required for this reaction to occur?
5.
What is the heat of reaction
6.
Is the reaction exothermic or endothermic?
7.
What is the activation energy of the reverse reaction?
8.
What is the heat of reaction of the reverse reaction?
9.
Is the reverse reaction exothermic or endothermic?
10.
If a catalyst were added, which quantities, if any, would change? (give letters)
11.
Would the activation energy increase, decrease, or remain unchanged?
12.
Would the heat of reaction increase, decrease or remain unchanged?
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Energy Diagrams K
900
800
700
600
X+Y
(d)
(b)
Potential
500
Energy
(kJ)
(a)
(e)
(c)
400
Z
300
200
(f)
100
Reaction Coordinate
1.
What is the potential energy of the products?
2.
What is the potential energy of the activated complex?
3.
What is the potential energy of the reactants?
4.
How much activation energy is required for this reaction to occur?
5.
What is the heat of reaction
6.
Is the reaction exothermic or endothermic?
7.
What is the activation energy of the reverse reaction?
8.
What is the heat of reaction of the reverse reaction?
9.
Is the reverse reaction exothermic or endothermic?
10.
If a catalyst were added, which quantities, if any, would change? (give letters)
11.
Would the activation energy increase, decrease, or remain unchanged?
12.
Would the heat of reaction increase, decrease or remain unchanged? 1
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Energy Diagrams O
90
80
70
(d)
60
(a)
Z
Potential
50
Energy
(kJ)
(c)
40
30
(e)
(f)
X+Y
20
(b)
10
Reaction Coordinate
1.
What is the potential energy of the products?
2.
What is the potential energy of the activated complex?
3.
What is the potential energy of the reactants?
4.
How much activation energy is required for this reaction to occur?
5.
What is the heat of reaction
6.
Is the reaction exothermic or endothermic?
7.
What is the activation energy of the reverse reaction?
8.
What is the heat of reaction of the reverse reaction?
9.
Is the reverse reaction exothermic or endothermic?
10.
If a catalyst were added, which quantities, if any, would change? (give letters)
11.
Would the activation energy increase, decrease, or remain unchanged?
12.
Would the heat of reaction increase, decrease or remain unchanged?
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Energy Diagrams S
1800
1600
1400
1200
X+Y
(d)
(b)
Potential
1000
Energy
(kJ)
(a)
(e)
(c)
800
Z
600
400
(f)
200
Reaction Coordinate
1.
What is the potential energy of the products?
2.
What is the potential energy of the activated complex?
3.
What is the potential energy of the reactants?
4.
How much activation energy is required for this reaction to occur?
5.
What is the heat of reaction
6.
Is the reaction exothermic or endothermic?
7.
What is the activation energy of the reverse reaction?
8.
What is the heat of reaction of the reverse reaction?
9.
Is the reverse reaction exothermic or endothermic?
10.
If a catalyst were added, which quantities, if any, would change? (give letters)
11.
Would the activation energy increase, decrease, or remain unchanged?
12.
Would the heat of reaction increase, decrease or remain unchanged?
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Energy Diagrams W
180
160
140
(d)
120
(a)
Z
Potential
100
Energy
(kJ)
(c)
80
60
(e)
(f)
X+Y
40
(b)
20
Reaction Coordinate
1.
What is the potential energy of the products?
2.
What is the potential energy of the activated complex?
3.
What is the potential energy of the reactants?
4.
How much activation energy is required for this reaction to occur?
5.
What is the heat of reaction
6.
Is the reaction exothermic or endothermic?
7.
What is the activation energy of the reverse reaction?
8.
What is the heat of reaction of the reverse reaction?
9.
Is the reverse reaction exothermic or endothermic?
10.
If a catalyst were added, which quantities, if any, would change? (give letters)
11.
Would the activation energy increase, decrease, or remain unchanged?
12.
Would the heat of reaction increase, decrease or remain unchanged?
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Energy Diagrams A
360
320
280
240
X+Y
(d)
(b)
Potential
200
Energy
(kJ)
(a)
(e)
(c)
160
Z
120
80
(f)
40
Reaction Coordinate
1.
What is the potential energy of the products?
2.
What is the potential energy of the activated complex?
3.
What is the potential energy of the reactants?
4.
How much activation energy is required for this reaction to occur?
5.
What is the heat of reaction
6.
Is the reaction exothermic or endothermic?
7.
What is the activation energy of the reverse reaction?
8.
What is the heat of reaction of the reverse reaction?
9.
Is the reverse reaction exothermic or endothermic?
10.
If a catalyst were added, which quantities, if any, would change? (give letters)
11.
Would the activation energy increase, decrease, or remain unchanged?
12.
Would the heat of reaction increase, decrease or remain unchanged?
Topic 5.2: Calculation of Enthalpy Changes
Heat Transfer Problems
(These questions are more physics based, but does give you a better understanding of how thermal
energy moves around a system and practice the use of the specific heat equation)
1. What is the specific heat of a substance? What does it mean?
2. Surely a cotton bath mat and the tile floor in your bathroom are the same temperature. So
why is it comfortable to stand on the bath mat but not on the tiled bathroom floor, which
seems very cold?
3. After watching the video clip on pizza, why is it that the cheese portion of the pizza burns the
roof of your mouth but the crust doesn't burn your tongue?
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4. How many joules are required to raise the temperature of 100 g of water from 20 to 50 ºC?
(12,552 J)
5. How many joules are required to raise the temperature of 100 g of iron from 20 to 50 ºC?
(1350 J)
6. Why is there such a large difference in your answers between water and iron?
7. How many kilojoules are required to heat 6.3 kg of aluminum from -23 to 625 ºC ? (3633 kJ)
8. A 15.0 g block of aluminum at an initial temperature of 27.5°C absorbs 0.678 kJ of heat. What
is the final temperature of the block? (The specific heat of Al is 0.902 J g-1 C°.-1)
9. A sample of aluminum absorbed 9.86 J of heat, and its temperature increased from 23.2°C to
30.5°C. What was the mass of the aluminum sample? (The specific heat of aluminum is 0.902 J
g-1 C°.-1.)
10. What is the resulting temperature when 35 g of water at 75°C is mixed with 15 g of water at
15°C? (The specific heat of water is 4.184 J g-1 C°.-1.)
11. What is the resulting temperature when 100 g of water at 75°C is mixed with 30 g of ice at 0
°C? Assume that all of the ice has melted. The specific heat of water is 4.184 J g-1 C°.-1 and
the heat of fusion of ice is 335 Jg-1.
12. The specific heat of lead is 0.13 Jg-1C°-1. How many J of heat would be required to raise the
temperature of 15 g of lead from 22°C to 37°C?
13. How much heat is required to raise the temperature of 125 g of water from 2.2 C° to 48.6
C°? The specific heat of water is 4.184Jg-1C°-1.
14. What heat is needed to raise 3.4 kg of lead from 23oC to 58oC?
Some specific heats
(1.5E4 J)
15. If 23.0 kg of copper at 21.0oC absorbs 45.6 kJ of heat, what will (JoC-1kg-1)
be its final temperature? (26.1oC)
16. If some aluminum at 57.0oC, cools to 24.1oC, and gives off 13.4 kJ
H2O liquid
4186
of heat, what is its mass? (453 g)
2100
17. A 35.0 g of a mystery substance absorbs 314 J of heat and raises H2O ice
o
o its temperature by 2.14 C. What is its specific heat? (4190 J C
1
H2O steam 2010
kg-1)
Heat Lost = Heat Gained Problems:
Aluminum
900
1. 112 grams of a mystery liquid at 83.0 oC is mixed with 564 grams of
water initially at 22 oC. The final temperature of the mixture is 33.0
oC. What is the specific heat of the mystery liquid? (Assuming no
Iron
450
Copper
390
2.
3.
4.
5.
heat was lost to the surroundings) (4640 J kg-1 oC-1)
Lead
130
A piece of lead (c = 130 J/kgoC) at 82.0oC is mixed with 112 grams of water and an 87.5 g
aluminum (c = 900. g/kgoC) calorimeter cup initially at 25.0oC. The final temperature of the
system is 56.0oC. What is the mass of the piece of lead? (Assuming no heat was lost to the
surroundings) (5.02 kg)
89.2 g of a mystery substance is at 99.20 oC, and it is placed in a 95.0 g iron container holding 216
ml of water both at 21.01oC. The final temperature is 23.38oC. What is the specific heat of the
substance? (332 J/kg/oC)
A 347 g piece of copper at 98.0oC is placed in a Styrofoam cup containing 259 ml of water at 18.0
o
C. What will be the final temperature of equilibrium? (Ignore the Styrofoam) (26.9oC)
A 13.5 g piece of aluminum at 93.9oC is placed in an 82.0 g iron calorimeter containing 203 g of
water both at 23.0oC. What will be the final temperature? (24.0oC)
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6. If you drop a 16 g ice cube at 0.0 oC into a Styrofoam cup containing 241 ml of water at 20.0 oC
what will be the final temperature? (13.8oC)
7. You take an ice cube out of the freezer at -17.0oC, and drop it into a 67.0 g aluminum cup
containing 308 g of water at 23.0 oC. The final temperature is observed to be 12.7oC. What is
the mass of the ice cube? (33.0 g)
Chemistry Calorimetry Exercises
1. When 12.29 g of finely divided brass (60% Cu, 40% Zn) at 95.0oC is quickly stirred into 40.00 g
of water at 22.0oC in a calorimeter, the water temperature rises to 24.0oC. Find the specific
heat of brass.
Hints:
• The heat lost by the brass is gained by the surroundings (the water plus the calorimeter). What
relation can you therefore write between qbrass and qsurr?
• Since no information is given about the heat capacity of the calorimeter, you should assume it is
negligible.
• The final temperature of the brass is the same as the final temperature of the water. The
specific heat of water, s(H2O), is 4.184 J g-1 oC-1.
(Answers: qsurr = q(H2O) = 334.7 J; qbrass = - 334.7 J; sbrass = 0.38 J g-1 oC-1).
2. In an experiment, 400. mL of 0.600 M HNO3(aq) is mixed with 400. mL of 0.300 M Ba(OH)2(aq)
in a constant-pressure calorimeter having a heat capacity of 387 J/oC. The initial temperature
of both solutions is the same at 18.88oC, and the final temperature of the mixed solution is
22.49oC. Calculate the heat of neutralization in kJ per mole of HNO3.
Hints:
 The heat evolved in the neutralization reaction is gained by the surroundings (the mixed solution
plus the calorimeter). What relation can you therefore write between qrxn and qsurr?
 There are two contributions to qsurr. What are they? What assumptions (if any) need to be made
in calculating these contributions?
 Is this a limiting reagent problem, or are reactants supplied in the stoichiometric ratio given by
the equation? (Why do we care about this?)
 We want our answer in kJ per mole of HNO3. How do we calculate that?
(Answers: ΔT = 3.61oC; qsolution = 12083.4 J; qcalorimeter = 1397.1 J; qsurr = 13481 J; qrxn = 13481 J; ΔHneut(kJ/mol HNO3) = - 56.2 kJ/ mol HNO3).
3. The reaction of 1.0 mol of C to form carbon monoxide in the reaction 2 C(s) + O2(g)  2
CO(g) releases 113 kJ of heat. How much heat will be released by the combustion of 100 g of
C according the above information? (940 kJ)
4. When 1.0 mole of ethanol (C2H5OH) is burned, it releases -1367 KJ. How much heat is released
if 50 g of ethanol burn? (1500 kJ)
5. Octane (C8H18) is a component in gasoline used to fuel a car. If the heat of combustion of
Octane is -5471.5 kJ/mol of octane, how much energy is provided by burning 10.0 gallons? The
density of octane is 0.7028 g/mL. (1.28x106 kJ)
6. 26 g of water at 18oC are mixed with 49 g of water at 70oC. Find the final temperature of the
system. (52 oC)
7. 84 g of water at 22oC are mixed with 150 g of ethanol (C = 2.44 J/gC) at 88oC. Find the final
temperature of the system. (55.7 oC)
8. 75.0 g of water at 30oC are mixed with 83.8 g of a solid metal at 600oC. The final temperature
of the system is 50oC. What is the metal? (W)
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9. When 1.00 dm3 of 1.00 mol dm-3 Ba(NO3)2 solution at 25.0C is mixed with 1.00 dm3 of 1.00 mol
dm-3 Na2SO4 solution at 25.0C in a calorimeter, the white solid BaSO4 forms and the
temperature of the mixture increases to 28.1C. Assuming that the calorimeter absorbs only a
negligible quantity of heat, that the specific heat capacity of the solution is 4.18 J/C-g, and
that the density of the final solution is 1.0 g/cm3, calculate the enthalpy changer per mole of
BaSO4 formed. (25.9 kJ/mol)
10. In a coffee-cup calorimeter, 1.60 g of NH4NO3 is mixed with 75.0 g of water at an initial
temperature of 25.00C. After dissolution of the salt, the final temperature of the
calorimeter contents is 23.34C. Assuming the solution has a heat capacity of 4.18 J/C-g and
assuming no heat loss to the calorimeter, calculate the enthalpy change for the dissolution of
NH4NO3 in units of kJ/mol. (26.0 kJ/mol)
Phase Diagram Worksheet
Refer to the phase diagram below when answering the questions on this worksheet:
1. What is the normal freezing point of this substance?
2. What is the normal boiling point of this substance?
3. What is the normal freezing point of this substance?
4. If I had a quantity of this substance at a pressure of 1.25 atm and a temperature of 3000 C
and lowered the pressure to 0.25 atm, what phase transition(s) would occur?
5. At what temperature do the gas and liquid phases become indistinguishable from each other.
6. If I had a quantity of this substance at a pressure of 0.75 atm and a temperature of -1000 C,
what phase change(s) would occur if I increased the temperature to 6000 C? At what
temperature(s) would they occur?
Topic 5.3 Hess’s Law
Notes:Calculations using Hess's Law and Heats of Formation
Enthalpy of reaction values have been determined experimentally for numerous reactions, and these ∆H
values may be used to calculate ∆H values for other reactions involving the same chemical species. The reason
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this is possible is that enthalpy H is a state property so ∆H is independent of path. (Similarly, the height of a
mountain above sea level is independent of the path you follow to climb the mountain.) Because ∆H is
independent of path, we can determine the enthalpy of foods by burning them in a bomb calorimeter in the
laboratory to produce the same products that are obtained by the complicated metabolic pathways in our body!
There are two principle methods used to calculate ∆H values for a reaction, both of which are based on the
idea that ∆H for a reaction is independent of the path used to go from reactants to products. The first makes
o
use of Hess's Law while the second employs tabulated heats of formation H f (kJ/mol).
Use of Hess's Law to Calculate ∆H
Hess's Law states that ∆H for a reaction can be found indirectly by summing ∆H values for any set of
reactions which sum to the desired reaction. Usually before reactions are added together, some of them must
be reversed and/or multiplied by a factor n in order that they sum to the desired reaction. In this process the
rules are:
•
Whenever you multiply a reaction by n, ∆H for the reaction is also multiplied by n.
•
If you reverse a reaction, ∆H changes sign.
Problem (1) below is an example of how this procedure is used.
Use of Tabulated Heats of Formation to Calculate ∆H
o
The Standard Heat of Formation H f (kJ/mol) for a compound is the heat absorbed (or released) in
forming one mole of the compound from its elements in their standard states at 1 bar
o
(≈ 1 atm) pressure and the specified temperature (usually 25 oC). Thus H f (kJ/mol) for acetone CH3COCH3
is the heat of the reaction
3 C(s) + 3 H2(g) + 1/2 O2(g)  CH3COCH3(l)
∆H = – 246.8 kJ
o
o
and so H f (CH3COCH3(l)) = – 246.8 kJ/mol). By definition, H f (kJ/mol) = 0 for any element in its standard
state at 25oC and 1 bar.
These tables may be used to calculate the standard enthalpy change,
o
, for any reaction for which the
Hrxn
heats of formation of all reactants and products are known:
o
Hrxn 
n
products
o
H f ( prod) 
prod
n
react
o
H f (react)
reactants
This equation tells us to sum the enthalpies of formation of each product multiplied by its stoichiometric
coefficient in the reaction equation and then to subtract the enthalpy of formation of each reactant multiplied
by its stoichiometric coefficient. We use this equation to work problem (2).
o
applies for a balanced equation with specific stoichiometric amounts. If a different number of
Hrxn
moles reacts, the heat absorbed or evolved will change proportionately (problem 3).
1. From the following heats of reaction
2 SO2(g) + O2(g)  2 SO3(g)
∆H = – 196 kJ
(a)
2 S(s) + 3 O2(g)  2 SO3 (g)
∆H = – 790 kJ
(b)
∆H = ? kJ
(c)
calculate the heat of reaction for
S(s) + O2(g)  SO2(g)
Method: Use Hess’s Law to solve this problem:
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•
Identify a species in the target equation (c) which is on the correct side in only one of the listed equations, (a) or (b).
Multiply the entire equation, and its ∆H value by the factor n necessary to make the stoichiometric coefficient for the
species identical to that in equation (c).
•
Reverse a listed equation, (a) or (b), and change the sign of its ∆H value if it contains a species which is on the wrong
side of the target equation (c); next multiply the entire reversed equation by the factor n necessary to make the
stoichiometric coefficient for the species identical to that in equation (c). The ∆H value for the rewritten equation is (
- n) times that of the original equation. (Ignore any species present in both equations (a) and (b).)
•
Test to see if your rewritten equations now sum to the desired equation (c). If they do, the ∆H value for equation (c)
is the sum of the ∆H values of the rewritten equations.
Answer: (1/2) Eq(b) + ( - 1/2) Eq(a) = Eq(c); Thus, (1/2)H (b) + ( – 1/2)H (a) = H (c) so
H(c) =(1/2)( – 790 kJ) + ( – 1/2)( – 196 kJ) =  kJ
2. Calculate the standard reaction enthalpy for the photosynthesis reaction,
o
Hrxn = ? kJ
6 CO2(g) + 6 H2O(l)  C6H12O6(s) + 6 O2(g)
Note: Search for the heat of formation of glucose, C6H12O6(s)
o
o
o
Answer: Use Eq (2) with H f (CO2(g)) = – 393.5 kJ/mol, H f (H2O(l)) = – 285.8 kJ/mol, H f (C6H12O6(s)) = – 1274.5
o
kJ/mol, and H f (O2(g)) = 0 kJ/mol. Thus Ho = 2801.3 kJ ≈ 2801 kJ for the photosynthesis reaction.
3. Is the photosynthesis reaction above endothermic or exothermic? How much heat is absorbed or
evolved if 11.0 g of CO2(g) reacts completely with excess water to form glucose and oxygen gas?
Answers: Endothermic. 117 kJ of heat is absorbed.
Hess's Law Worksheet
For each problem below, write the equation and show your work. Always show units.
1. Consider the following hypothetical reactions:
A  B
∆H = +30 kJ
B  C
∆H = +60 kJ
Use Hess’s law to calculate the enthalpy change for reaction A  C.
Construct an enthalpy diagram for substances A, B, and C, and show
how Hess’s law applies.
2. Suppose you are given the following hypothetical reactions:
XY
∆H = -40 kJ
X  Z
∆H = -95 kJ
Use Hess’s law to calculate the enthalpy change for the reaction Y  Z.
Construct an enthalpy diagram for substances X, Y, and Z, and discuss its relevance to Hess’s
law.
3. From the following heats of reaction:
2H2 (g) + O2 (g)  2H2O (g)
3O2 (g)  2O3 (g)
Calculate the heat of the reaction
3H2 (g) + O3 (g)  3H2O (g)
∆H = -483.6 kJ
∆H = +284.6 kJ
4. From the following enthalpies of reaction:
H2 (g) + F2 (g)  2HF (g)
∆H = - 537kJ
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C (s) + 2 F2 (g)  CF4 (g)
2C (s) + 2 H2 (g)  C2H4 (g)
∆H = - 680 kJ
∆H = + 52.3 kJ
Calculate the ΔH for the reaction of ethylene with F2.
C2H4 (g) + 6F2 (g)  2CF4 (g) + 4HF(g)
5. Given the following data:
N2 (g) + O2 (g)  2NO (g)
2NO (g) + O2 (g)  2NO2 (g)
2N2O (g)  2N2 (g) + O2 (g)
∆H = + 180.7 kJ
∆H = - 113.1 kJ
∆H = - 162.3 kJ
Use Hess’s law to calculate ΔH for the reaction.
N2O (g) + NO2 (g)  3NO (g)
(Answers: 1a) 90 kJ 2a) -55 kJ3) -867.7 kJ4)-2486.3 kJ5) 156.1 kJ)
6. Calculate ∆H for the reaction: C2H4 (g) + H2 (g) C2H6 (g), from the following data.
C2H4 (g) + 3O2 (g)  2 CO2 (g) + 2H2O(l) ∆H = -1411. kJ
C2H6(g) + 3½O2 (g)  2CO2 (g) + 3H2O(l) ∆H = -1560. kJ
H2 (g) + ½O2 (g)  H2O(l)
∆H = -285.8 kJ
7.
Calculate ∆H for the reaction 4 NH3 (g) + 5 O2 (g) 4 NO (g) + 6 H2O (g), from the following
data.
N2 (g) + O2 (g)  2NO (g)
N2 (g) + 3H2 (g)  2NH3 (g)
2H2 (g) + O2 (g)  2H2O (g)
∆H = -180.5 kJ
∆H = -91.8 kJ
∆H = -483.6 kJ
8. Find ∆H for the reaction 2H2(g) + 2C(s) + O2(g)  C2H5OH(l), using the following
thermochemical data.
C2H5OH (l) + 2O2 (g)  2CO2 (g) + 2H2O (l)
C (s) + O2 (g)  CO2 (g)
H2 (g) + ½O2 (g)  H2O (l)
∆H = -875. kJ
∆H = -394.51kJ
∆H = -285.8 kJ
9. Calculate ΔH for the reaction CH4 (g) + NH3 (g)  HCN (g) + 3 H2 (g), given:
N2 (g) + H2 (g)  2NH3 (g)
C (s) + 2H2 (g)  CH4 (g)
H2 (g) + 2C (s) + N2 (g)  2HCN (g)
∆H = -91.8 kJ
∆H = -74.9 kJ
∆H = +270.3 kJ
10. Calculate ∆H for the reaction 2 Al (s) + 3 Cl2 (g) 2 AlCl3 (s) from the data.
2Al (s) + 6HCl (aq)  2AlCl3 (aq) + 3H2 (g) ∆H = -1049. kJ
HCl (g)  HCl (aq)
∆H = -74.8 kJ
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H2 (g) + Cl2 (g) 2 HCl (g)
AlCl3 (s)  AlCl3 (aq)
∆H = -1845. kJ
∆H = -323. kJ
11. Calculate ΔHº for Mg(s) + ½ O2(g)  MgO(s) given the equations:
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g) ΔHº = –462 kJ
MgO(s) + 2HCl(aq)  MgCl2(aq) + H2O(l)
ΔHº = –146 kJ
2H2(g) + O2(g)  2H2O(l)
ΔHº = –572 kJ
Then draw the Enthalpy Diagram for the overall reaction.
12. Calculate ΔHº for C2H4(g) + 2H2(g)  2CH4(g) given the equations:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
ΔHº = –890 kJ
C2H4(g) + 3O2(g)  2CO2(g) + 2H2O(l)
ΔHº = –1411 kJ
2H2(g) + O2(g)  2H2O(l)
ΔHº = –572 kJ
Then draw the Enthalpy Diagram for the overall reaction.
ΔH = -602 kJ
ΔH = -203 kJ
Calculate ΔH (kJ) for the reaction: NH3(g) + 5/2 O2(g)  2 NO(g) + 3 H2O(g) given the following:
½ N2(g) + ½ O2(g)  NO(g)
ΔH= 90.4 kJ mol-1
½ N2(g) + 3/2 H2(g)  NH3(g)
ΔH= -46.0 kJ mol-1
H2(g) + ½ O2 (g)  H2O(g)
ΔH= -241.8 kJ mol-1
ΔH = -452.6 kJ
13.
Calculate ΔH (kJ) for the reaction: Ca2+(aq) + 2OH-(aq) + CO2(g)  CaCO3(s) + H2O(l) given
the following:
CaCO3(s)  CaO(s) + CO2(g)
ΔH = 178.1 kJ mol-1
CaO(s) + H2O(l)  Ca(OH)2(s)
ΔH = -64.8 kJ mol-1
Ca(OH)2(s)  Ca2+(aq) + 2OH-(aq)
ΔH = -11.7 kJ mol-1
ΔH = -101.6 kJ
14.
Calculate ΔH (kJ) in kJ for the reaction: 2C(s) + H2(g)  C2H2(g) given the following:
C2H2(g) + 5/2O2 (g)  2CO2(g) + H2O(l)
ΔH= - 1300 kJ mol-1
C(s) + O2(g)  CO2(g)
ΔH= -394 kJ mol-1
H2(g) + ½ O2(g)  H2O(l)
ΔH= -286 kJ mol-1
ΔH = 226 kJ
15. The bombardier beetle uses an explosive discharge as a defensive mechanism. The chemical
reaction involved is the oxidation of hydroquinone by hydrogen peroxide to produce quinone and
water.
C6H4(OH)2(aq) + H2O2(aq)  C6H4O2(aq) + 2H2O(l)
Calculate ΔH in kJ for the reaction given the following:
C6H4(OH)2(aq)  C6H4O2(aq) + H2(g)
H2(g) + O2(g)  H2O2(aq)
H2(g) + ½ O2(g)  H2O(g)
H2O(g)  H2O(l)
ΔH = +177.4 kJ mol-1
ΔH = -191.2 kJ mol-1
ΔH= -241.8 kJ mol-1
ΔH = -43.8 kJ mol-1
16. The enthalpies of these reactions at 25°C are given in kJ.
C4H9OH (l) + 6 O2(g)
(C 2H5)2O(l) + 6 O2(g)


4 CO2 (g) + 5 H2O (g)
4 CO2 (g) + 5 H2O (g)
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ΔH = -202 kJ
∆H° in kJ mol-1
-2456
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From a consideration of the above information, calculate the change in enthalpy for
C4H9OH (l)  (C 2H5)2O (l)
17. The enthalpies of three reactions at 25°C are given in kJ.
∆H° in kJ mol-1
2 H2 (g) + O2 (g)  2 H2O (l)
-572
2 C2H6 (g) + 7 O2 (g)  4 CO2(g) + 6 H2O(l)
- 3120
C2H4 (g) + 3 O2 (g)  2 CO2 (g) + 2 H2O (l)
- 1411
Use this information to calculate the change in enthalpy for
C2H6 (g)  C2H4 (g) + H2 (g)
Is the reaction endothermic or exothermic? Explain.
18. The enthalpies of three reactions at 25°C are given in kJ.
∆H° in kJ mol-1
Mg (s) + 1/2 O2(g)  MgO (s)
-601.7
C(s) + O2 (g)
 CO2 (g)
-393.5
Mg (s) + C (s) + 3/2 O2 (g)  Mg CO3 (g)
-1095.8
Use this information to calculate the change in enthalpy for
MgO (s) + CO2 (g)  MgCO3 (s)
19. The enthalpies of three reactions at 25°C are given in kJ.
C(s) + 1/2 O2 (g)  CO (g)
C(s) + O2 (g)  CO2 (g)
CO (g) + 1/2 O2 (g)  CO2 (g)
H2 (g) + 1/2 O2 (g)  H2O (g)
Use this information to calculate the change in enthalpy for
C(s) + H2O (g)  CO (g) + H2
∆H° in kJ mol-1
- 110.5
- 393.5
-283.0
-241.8
(g)
20. The enthalpies of three reactions at 25°C are given in kJ.
N2 (g) + 2O2 (g)  2 NO2 (g)
2 NO (g) + O2 (g)  2 NO2 (g)
∆H° in kJ mol-1
+ 33.2
- 57.1
From a consideration of the above information, calculate the change in enthalpy for
N2 (g) + O2 (g)  2 NO (g)
Hess’s Law Challenge Worksheet
1.
Use the data below, as appropriate, to answer these questions.
Cl2(g)  2 Cl(g)
H2(g  2 H(g)
HCl(g)  H(g) + Cl(g)
Hz(g) + 1/2 O2(g)  HzO(g)
O2(g)  2 O(g)
∆H= +242kJmol-1
∆H = + 436 kJ mol-1
∆H = + 431 kJ mol-1
∆H = -242kJmol'1
∆H = + 496 kJ mol-1
a) State Hess's Law.
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b) i. Calculate the standard enthalpy change for the following reaction.
H2(g) + Cl2(g)  2HCl(g)
ii. ln tne dark, the reaction between hydrogen and chlorine starts when the mixture is
heated to above 600'C. Suggest why heat is necessary in order to start the reaction.
iii. What happens in the first step in the reaction between hydrogen and chlorine? Give
a reason why this occurs first.
c) Calculate a value for the mean bond enthalpy of the H-O bond in water.
2. a) Define the term standard enthalpy of formation (∆Hf).
b) Give the chemical equation for which he enthalpy change is the standard enthalpy of formation
of methane.
c) Use the following data to calculate a value for the enthalpy of formation of methane
C(s) + O2(g)  CO2(g)
∆H = -394 kJ mol-1
H2(g) + ½ O2(g)  H2O(g)
∆H = -242kJmol-1
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
∆H = - 802 kJ mol-1
d) Use the following data to calculate the mean bond enthalpy values for the C-H and C-C bonds.
CH4(g)  C(g) + 4H(g)
∆H = + 1648 kJ mol-1
C2H6(g) 2C(g) + 6H(g)
∆H = + 2820 kJ mol-1
3)
a) Write a chemical equation, including state symbols, for the reaction which is used to define
the enthalpy of formation of aluminium chloride (AlCl3).
b) State the additional conditions necessary if the enthalpy change for this reaction is to be the
standard enthalpy of formation, Hf at 298 K.
c) Use the standard enthalpies of formation given below to calculate a value for the standard
enthalpy change for the following reaction.
AlCl3(s) + 6 H2O(l)  AlCl3.6H2O(s)
4) a) Explain the meaning of the term standard enthalpy of
combustion of a compound.
b) Write an equation to represent the enthalpy of combustion of
methanol.
c) The apparatus shown was used to find the enthalpy of
combustion of methanol. In this experiment, 0.75 grams of methanol
was burned. The water in the apparatus, which had a volume of
100cm3, rose in temperature from 18.5oC to 50.3oC.
i. Calculate the heat released by the 0.75 grams of methanol,
given that the specific heat capacity of water is 4.2Jg-1 K-1.
ii. Calculate the enthalpy of combustion of methanol
iii. The actual value of the enthalpy of combustion of methanol is
-726 kJ mol-l. Account for the difference between this value and the experimental value.
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Topic 5.4 Bond Enthalpies
Bond Enthalpies
The energy required to break a covalent bond in a gaseous substance is called the bond energy for
the chemical bond. For example, for the diatomic molecule HCl is defined as the standard enthalpy
change for the reaction:
HCl (g)

H (g) + Cl (g)
∆H° = 431 kJ mol-1
The enthalpy change for the equation
H2O (g)  O (g) + 2 H (g)
is ∆Hreaction = + 925 kJ. Since water is made up of two O-H covalent bonds, the value of ∆Hreaction is the energy
required to break these bonds. Since ∆Hreaction is positive, at constant pressure 925 kJ of energy is required
to break two moles H-O covalent bonds in 1 mole of H2O. The average molar bond enthalpy for a hydrogen to
oxygen bond is equal to one half of the total energy input as heat required to break two moles of hydrogen to
oxygen bonds or about 463 kJ mol-1.
It has been found that the hydrogen to oxygen molar bond enthalpy in a variety of compounds is about the same
as the H-O bond enthalpy in water. Bond enthalpies are a function of the particular atoms forming the bond
rather than the particular substance in which they are incorporated.
Some typical molar bond enthalpies are listed below:
Bond
Bond
enthalpy
Bond
Bond
enthalpy
Bond
Bond
enthalpy
O-H
O-O
C-O
O=O
C=O
C-C
C=C
C= C
C-H
464
142
351
502
730
347
615
811
414
C-F
C-Cl
C-Br
C-N
C=N
C= N
N-H
N-N
N=N
439
439
276
293
615
890
390
159
418
N= N
F-F
Cl-Cl
Br-Br
H-H
H-F
H-Cl
H-Br
H-S
945
155
243
192
435
565
431
368
364
In general it is helpful to picture a chemical reaction as taking place in two steps. First the chemical bonds in
the reactants are broken down into their constituent atoms and then the atoms are rejoined to form the
product molecules. The first step requires an input of energy to break the chemical bonds of the reactants.
The second step releases energy as new chemical bonds are formed. The enthalpy change for the reaction,
∆Hrxn, is the difference between the energy required to break the chemical bonds of the reactants and the
energy released in the formation of the chemical bonds for the products.
Example 1: Calculate the enthalpy of reaction for the following reaction:
CH4 (g) + Cl2 (g)
 CH3Cl (g) + HCl(g)
In the above reaction it is necessary to break one C-H bond and one Cl-Cl bond and to form one C-Cl bond and
one H-Cl bond.
Reactants: (Bonds broken)
4
C-H
414 kJ x 4
= 1656 kJ
1
Cl-Cl
243 kJ x 1
= 243 kJ
Total = 1899 kJ (absorbed ∆H is +)
Products (Bonds formed)
3
C-H
414 kJ x 3
= 1242 kJ
1 C-Cl 439 kJ x 1
= 439
1 H-Cl
431 kJ x 1 =
431
Total = 2112 kJ (Released ∆H is -)
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Change in enthalpy = - 213 kJ
Exercises:
1.
The enthalpy change for the reaction:
ClF3 (g)  Cl (g) + 3 F (g) is 514 kJ.
Using this information and the bond enthalpies on the previous page, calculate the average Cl-F bond
energy in ClF3.
2. Use the bond enthalpies on the previous page to calculate ∆Hreaction for the following reaction:
CH4 (g) + 2 Cl2 (g)  CH2Cl2 (g) + 2 HCl (g)
3. Hydrazine, N2H4 and its derivatives are used as rocket fuels. Use the molar bond enthalpies on the
previous page to calculate the molar enthalpy of formation of N2H4 (g).
4. The formation of water from oxygen and hydrogen involves the reaction:
2 H2 (g) + O2 (g)  2 H2O (g)
Use the bond enthalpies on the previous page to calculate the enthalpy of formation of H 2O (g).
5. The formation of ammonia from nitrogen and hydrogen involves the reaction:
3 H2 (g) + N2 (g)  2 NH3 (g)
Use the bond enthalpies on the previous page to calculate the enthalpy of formation of NH 3 (g) .
Bond Enthalpies 2
1.
2.
3.
4.
Which process releases energy: breaking a bond or forming a bond?
Which process requires energy: breaking a bond or forming a bond?
Define bond energy.
If the energy used to break bonds is greater than the energy released in the
formation of new bonds, is the reaction endothermic or exothermic?
5. If the energy used to break bonds is less than the energy released in the
formation of new bonds, is the reaction endothermic or exothermic?
6. Look at the table of selected bond energies to the right.
a. Which is the strongest single bond on this table?
b. Which is the weakest single bond on this table?
7. (Note: This is not a question. It is an explanation of how to calculate the total
energy from a reaction, based on bond energies. There are supposed to be empty
spaces in the table.)
Let’s examine the electrolysis of water (fig 4 on page 587). The general reaction is
2H2O → 2H2 + O2 or H−O−H + H−O−H → H−H + H−H + O=O
The overall heat of reaction can be calculated as follows:
Thus, this is an endothermic reaction (energy required) that absorbs 466 kJ.
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The thermochemical equation is 2H2O + 466 kJ → 2H2 + O2
Calculate the heat of reaction for H2 + Cl2 → 2HCl (by completing a table similar to the one above. Write the
thermochemical equation for this reaction.
8. Calculate the theoretical value molar heat of reaction for burning paraffin (C25H52) in oxygen.
a. Draw Lewis structures for O2, H2O, and CO2.
b. Write the balanced equation for the combustion of C25H52.
c. Fill in this table to calculate the theoretical molar heat of combustion for C25H52
9. Calculate the molar heat of combustion & the specific heat of combustion
C2H5OH (shown right).
for
Using Hess’s Law and Bond Energy to calculate enthalpy change.
1. Given the enthalpy changes for the reactions below
2 H2O2 (aq)  2 H2O (l) + O2 (g) ∆H = -200 kJ mol-1
2 H2 (g) + O2 (g)  2 H2O (l)
∆H = -600kJ mol-1
what will be the enthalpy change for : H2 (g) + O2 (g)  H2O2
(aq)
?
A. -200 kJ mol-1
B. -400 kJ mol-1
C. -600 kJ mol-1
D. -800 kJ mol-1
2. Iron and chlorine react directly to form iron(III) chloride, not iron(II) chloride. Therefore it is not possible
to directly measure the enthalpy change for the reaction
Fe(s) + Cl2 (g)  FeCl2 (s)
The enthalpy changes for the formation of iron(III) chloride from the reaction of chlorine with iron and with
iron(II) chloride are given below. Use these to calculate the enthalpy change for the reaction of iron with
chlorine to form iron (II) chloride.
2 Fe (s) + 3 Cl2 (g)  2 FeCl3 (s) ΔH = -800 kJ mol-1
2 FeCl2 (s) + Cl2 (g)  2 FeCl3 (s) ΔH = -120 kJ mol-1
3. The Romans used calcium oxide (CaO) as mortar in stone structures. The CaO was mixed with water to give
Ca(OH)2, which slowly reacted with CO2 in the air to give limestone.
Ca(OH)2 (s) + CO2 (g)  CaCO3 (s) + H2O (l)
Use the enthalpies of the following reactions to calculate the enthalpy change for the reaction above.
Ca (s) + O2 (g) + H2 (g)  Ca(OH)2 (s)
ΔH = -986.1 kJ/mol
C (s) + O2 (g)  CO2 (g)
ΔH = -394.1 kJ/mol
H2(g) + 1/2 O2(g)  H2O(l)
ΔH = -285.8 kJ/mol
Ca (s) + C (s) + 3/2 O2 (g)  CaCO3(s)
ΔH = -1206.9 kJ/mol
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4. Which of the following equations is equivalent to the bond enthalpy of the carbon-oxygen bond in
carbon monoxide?
A. CO (g)  C(g) + O (g)
B. CO (g)  C(s) + O (g)
C. CO (g)  C(s) + ½ O2 (g)
D. CO (g)  C(g) + ½ O2 (g)
5. The bond enthalpy of the N-O bond in nitrogen dioxide is 305 kJ mol-1. If that of the bonds in the
oxygen molecule and the nitrogen molecule are 496 kJ/mol and 944 kJ/mol respectively, what
will be the enthalpy change for the following reaction?
N2 (g) + 2 O2 (g)  2 NO2 (g)
A. 716 kJ mol-1
B. 1135 kJ mol-1
C. 1326 kJ mol-1
D. 1631 kJ mol-1
6. The enthalpy change for the following reaction is +688 kJ mol-1.
N2 (g) + 3 Cl2 (g)  2 NCl3 (g)
Calculate the bond enthalpy of the N-Cl bond, given that the bond enthalpies of the nitrogen molecule and the
chlorine molecule are 944 kJ mol-1 and 242 kJ mol-1 respectively.
7. Use bond energy data to calculate the enthalpy change when cyclopropane reacts with hydrogen to form
propane. The actual value is –159 kJ mol-1. Give the reasons why you think this differs from the value you have
calculated.
[Bond energies in kJ mol-1: C-C = 348; C-H = 412; H-H = 436]
The reaction is shown in the following chemical equation:
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Topic 5: Energetics Exam Questions
1.
When some solid barium hydroxide and solid ammonium thiosulfate were reacted
together, the temperature of the surroundings was observed to decrease from 15 ºC to – 4
ºC. What can be deduced from this observation?
A.
B.
C.
D.
The reaction is exothermic and ∆H is negative.
The reaction is exothermic and ∆H is positive.
The reaction is endothermic and ∆H is negative.
The reaction is endothermic and ∆H is positive.
(Total 1 mark)
2.
Which process represents the C–Cl bond enthalpy in tetrachloromethane?
A.
B.
C.
D.
CCl4(g) → C(g) + 4Cl(g)
CCl4(g) → CCl3(g) + Cl(g)
CCl4(l) → C(g) + 4Cl(g)
CCl4(l) → C(s) + 2Cl2(g)
(Total 1 mark)
3.
Some water is heated using the heat produced by the combustion of magnesium metal.
Which values are needed to calculate the enthalpy change of reaction?
I.
The mass of magnesium
II.
The mass of the water
III. The change in temperature of the water
A.
B.
C.
D.
I and II only
I and III only
II and III only
I, II and III
(Total 1 mark)
3. In a reaction that occurs in 50 g of aqueous solution, the temperature of the reaction mixture
increases by 20 °C. If 0.10 mol of the limiting reagent is consumed, what is the enthalpy
change (in kJ mol–1) for the reaction? Assume the specific heat capacity of the solution
4. = 4.2 kJ kg–1 K–1.
A.
B.
C.
D.
–0.10 × 50 × 4.2 × 20
–0.10 × 0.050 × 4.2 × 20
 50  4.2  20
0.10
 0.050 4.2  20
0.10
(Total 1 mark)
5.
Use the average bond enthalpies below to calculate the enthalpy change, in kJ, for the
following reaction.
H2(g) + I2(g) → 2HI(g)
Bond
Bond energy / kJ mol–1
H–H
440
I–I
150
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H–I
A.
B.
C.
D.
300
+290
+10
–10
–290
(Total 1 mark)
6.
Which substance does not conduct electricity?
A.
B.
C.
D.
Solid zinc
Molten zinc
Solid zinc chloride
Molten zinc chloride
(Total 1 mark)
7.
1.0 g of sodium hydroxide, NaOH, was added to 99.0 g of water. The temperature of the
solution increased from 18.0 ºC to 20.5 ºC. The specific heat capacity of the solution is
4.18 J g–1 K–1.
Which expression gives the heat evolved in kJ mol–1?
A.
2.5 100.0  4.18 1000
40.0
B.
2.5 100.0  4.18
1000 40.0
C.
2.5 100.0  4.18  40.0
1000
D.
2.5 1.0  4.18  40.0
1000
(Total 1 mark)
8.
Which is true for a chemical reaction in which the products have a higher enthalpy than
the reactants?
Reaction
∆H
A.
endothermic
positive
B.
endothermic
negative
C.
exothermic
positive
D.
exothermic
negative
(Total 1 mark)
9.
Consider the reaction between magnesium and hydrochloric acid. Which factors will affect
the reaction rate?
I.
The collision frequency of the reactant particles
II.
The number of reactant particles with E ≥ Ea
III. The number of reactant particles that collide with the appropriate geometry
A.
B.
C.
D.
I and II only
I and III only
II and III only
I, II and III
(Total 1 mark)
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10.
What is the energy, in kJ, released when 1.00 mol of carbon monoxide is burned
according to the following equation?
2CO(g) + O2(g) → 2CO2(g)
A.
B.
C.
D.
ΔHo = –564 kJ
141
282
564
1128
(Total 1 mark)
11.
Which of the following reactions are exothermic?
A.
B.
C.
D.
I.
CH4 + 2O2 → CO2 + 2H2O
II.
NaOH + HCl → NaCl + H2O
III. Br2 → 2Br
I and II only
I and III only
II and III only
I, II and III
(Total 1 mark)
12.
The specific heat of iron is 0.450 J g–1 K–1. What is the energy, in J, needed to increase
the temperature of 50.0 g of iron by 20.0 K?
A.
9.00
B.
22.5
C.
45.0
D.
450
(Total 1 mark)
13.
Two students were asked to use information from the Data Booklet to calculate a value for
the enthalpy of hydrogenation of ethene to form ethane.
C2H4(g) + H2(g) → C2H6(g)
John used the average bond enthalpies from Table 10. Marit used the values of
enthalpies of combustion from Table 12.
(a)
Calculate the value for the enthalpy of hydrogenation of ethene obtained using the
average bond enthalpies given in Table 10.
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(2)
(b)
Marit arranged the values she found in Table 12 into an energy cycle.
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Calculate the value for the enthalpy of hydrogenation of ethene from the energy
cycle.
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(1)
(c)
Suggest one reason why John’s answer is slightly less accurate than Marit’s
answer.
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(1)
(d)
John then decided to determine the enthalpy of hydrogenation of cyclohexene to
produce cyclohexane.
C6H10(l) + H2(g) → C6H12(l)
(i)
Use the average bond enthalpies to deduce a value for the enthalpy of
hydrogenation of cyclohexene.
...........................................................................................................................
...........................................................................................................................
(1)
(ii)
The percentage difference between these two methods (average bond
enthalpies and enthalpies of combustion) is greater for cyclohexene than it
was for ethene. John’s hypothesis was that it would be the same. Determine
why the use of average bond enthalpies is less accurate for the cyclohexene
equation shown above, than it was for ethene. Deduce what extra information
is needed to provide a more accurate answer.
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
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...........................................................................................................................
(2)
(Total 7 marks)
14.
Two students were asked to use information from the Data Booklet to calculate a value for
the enthalpy of hydrogenation of ethene to form ethane.
C2H4(g) + H2(g) → C2H6(g)
John used the average bond enthalpies from Table 10. Marit used the values of
enthalpies of combustion from Table 12.
(a)
Calculate the value for the enthalpy of hydrogenation of ethene obtained using the
average bond enthalpies given in Table 10.
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(2)
(b)
Determine the value for the enthalpy of hydrogenation of ethene using the values for
the enthalpies of combustion of ethene, hydrogen and ethane given in Table 12.
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(2)
(c)
Suggest one reason why John’s answer is slightly less accurate than Marit’s answer
and calculate the percentage difference.
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(2)
(d)
John then decided to determine the enthalpy of hydrogenation of cyclohexene to
produce cyclohexane.
C6H10(l) + H2(g) → C6H12(l)
(i)
Use the average bond enthalpies to deduce a value for the enthalpy of
hydrogenation of cyclohexene.
...........................................................................................................................
...........................................................................................................................
(1)
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(ii)
The percentage difference between these two methods (average bond
enthalpies and enthalpies of combustion) is greater for cyclohexene than it
was for ethene. John’s hypothesis was that it would be the same. Determine
why the use of average bond enthalpies is less accurate for the cyclohexene
equation shown above, than it was for ethene. Deduce what extra information
is needed to provide a more accurate answer.
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(2)
(Total 9 marks)
15.
The standard enthalpy change of three combustion reactions is given below in kJ.
2C2H6(g) + 7O2(g) → 4CO2(g) + 6H2O(l)
2H2(g) + O2(g) → 2H2O(l)
C2H4(g) + 3O2(g) → 2CO2(g) + 2H2O(l)
∆HO = –3120
∆HO = –572
ΔHO = –1411
Based on the above information, calculate the standard change in enthalpy, ∆HO, for the
following reaction.
C2H6(g) → C2H4(g) + H2
...............................................................................................................................................
.
...............................................................................................................................................
.
...............................................................................................................................................
.
...............................................................................................................................................
.
...............................................................................................................................................
.
...............................................................................................................................................
.
...............................................................................................................................................
.
...............................................................................................................................................
.
(Total 4 marks)
16.
In some countries, ethanol is mixed with gasoline (petrol) to produce a fuel for cars called
gasohol.
(i)
Define the term average bond enthalpy.
(2)
(ii)
Use the information from Table 10 of the Data Booklet to determine the standard
enthalpy change for the complete combustion of ethanol.
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CH3CH2OH(g) + 3O2(g) → 2CO2(g) + 3H2O(g)
(3)
(iii)
The standard enthalpy change for the complete combustion of octane, C8H18, is
–5471 kJ mol–1. Calculate the amount of energy produced in kJ when 1 g of ethanol
and 1 g of octane is burned completely in air.
(2)
(iv)
Ethanol can be oxidized using acidified potassium dichromate, K2Cr2O7, to form two
different organic products.
CH3CH2OH
A
B
State the structural formulas of the organic products A and B and describe the
conditions required to obtain a high yield of each of them.
(4)
(v)
Deduce and explain whether ethanol or A has the higher boiling point.
(2)
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(vi)
Ethene can be converted into ethanol by direct hydration in the presence of a
catalyst according to the following equation.
C2H4(g) + H2O(g)
CH3CH2OH(g)
For this reaction identify the catalyst used and state one use of the ethanol formed
other than as a fuel.
(2)
(Total 15 marks)
Topic 5: Exam Questions Markscheme
1.
D
[1]
2.
B
[1]
3.
D
4.
D
5.
C
6.
C
7.
C
8.
A
9.
D
10.
B
11.
A
12.
D
[1]
[1]
[1]
[1]
[1]
[1]
[1]
[1]
[1]
[1]
13.
(a)
energy required = C=C + H–H/612 + 436 and
energy released = C–C + 2(C–H)/347 +2(413) /
energy required = C=C + H–H + 4(C–H)/612 + 436 + 4(413) and
energy released = C–C + 6(C–H)/347 + 6(413);
∆H = (1048 – 1173)/(2700 – 2825) = –125 kJ mol–1;
2
(b)
∆H = –1411 + (–286) – (–1560) = –137 kJ mol–1;
1
(c)
the actual values for the specific bonds may be different to the average
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values / the combustion values referred to the specific compounds / OWTTE;
(d)
(i)
–125 kJ mol–1;
(ii)
average bond enthalpies do not apply to the liquid state / OWTTE;
the enthalpy of vaporization/condensation of cyclohexene and
cyclohexane / OWTTE;
1
1
2
[7]
14.
(a)
(b)
(c)
(d)
energy required = C=C + H–H/612 + 436 and
energy released = C–C + 2(C–H)/347 + 2(413) /
energy required = C=C + H–H + 4(C–H)/612 + 436 + 4(413) and
energy released = C–C + 6(C–H)/347 + 6(413);
∆H = (1048 – 1173)/(2700 – 2825) = –125 kJ mol–1
2
∆H = –1411 + (–286) – (–1560)/correct energy cycle drawn;
= –137 kJ mol–1;
Award [1 max] for incorrect or missing sign.
2
the actual values for the specific bonds may be different to the average
values / the combustion values referred to the specific compounds / OWTTE;
(137  125)
(percentage difference) =
× 100 = 8.76 %;
137
(137 125)
Accept
× 100 = 9.60 %.
125
(i)
–125 kJ mol–1;
(ii)
average bond enthalpies do not apply to the liquid state / OWTTE;
the enthalpy of vaporization/condensation of cyclohexene and
cyclohexane / OWTTE;
2
1
2
[9]
15.
(C2H6(g) + 3 12 O2(g) → 2CO2(g) + 3H2O(l))
∆HO = –1560;
(H2O(l) → H2(g) +
∆HO = +286;
1
2
O2(g))
(2CO2(g) + 2H2O(l) → C2H4(g) + 3O2(g))
(C2H6(g) → C2H4(g) + H2(g))
Allow other correct methods.
Award [2] for –137.
Allow ECF for the final marking point.
∆HO = +1411;
∆HO = +137(kJ);
4
[4]
16.
(i)
energy required to break (1 mol of) a bond in a gaseous molecule/state;
Accept energy released when (1 mol of) a bond is formed in a gaseous
molecule/state / enthalpy change when (1 mol of) bonds are made or
broken in the gaseous molecule/state.
average values obtained from a number of similar bonds/
compounds / OWTTE;
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(ii)
(iii)
(iv)
(v)
(vi)
Bonds broken
(1)(C–C) + (1)(O–H) + (5)(C–H) + (1)(C–O) + (3)(O=O)
= (1)(347) + (1)(464) + (5)(413) + (1)(358) + (3)(498) = 4728(kJ);
Bonds formed
(2 × 2)(C=O) + (3 × 2)(O–H)
= (4)(746) + (6)(464) = 5768 (kJ);
∆H = 4728 – 5768 = –1040 kJ mol–1 / –1040 kJ;
Units needed for last mark.
Award [3] for final correct answer.
Award [2] for +1040 kJ.
3
Mr(C2H5OH) = 46.08 / 46.1 and Mr(C8H18) = 114.26/114.3;
1 g ethanol produces 22.57 kJ and 1 g octane produces 47.88 kJ;
Accept values ranges of 22.5–23 and 47.8–48 kJ respectively.
No penalty for use of Mr = 46 and Mr = 114.
2
A: CH3CHO;
B: CH3COOH/CH3CO2H;
Accept either full or condensed structural formulas but not the
names or molecular formulas.
A: distillation;
B: reflux;
4
ethanol/CH3CH2OH;
hydrogen bonding (in ethanol);
Award second point only if the first is obtained.
2
(concentrated) H3PO4 /(concentrated) phosphoric acid / H2SO4/sulfuric acid;
dyes / drugs / cosmetics / solvent / (used to make) esters / (used in)
esterification/disinfectant;
2
[16]
Topic 15.1 Standard Enthalpy Changes of Reaction HL
The enthalpy of a chemical reaction can be determine from standard enthalpy of formation tables The standard
enthalpy of formation is the energy change that would result from the formation of a compound from its elements.
These values have been compiled for a wide variety of substances and are tabulated. Since enthalpy values may vary
with temperature and pressure of a gas, the standard tables are calibrated at 25 oC (298K) and 1 atmosphere.
Enthalpy values also depend on the physical state of a compound. Since enthalpies of formation are defined as the
difference in enthalpies of a compound and its elements, the standard enthalpies of formations for elements are
assumed to be zero.
The enthalpy change for a chemical reaction can be computed from this equation
Hreaction =
 H products
-

H reactants
Example
Use the standard enthalpy of formation table to determine the enthalpy of reaction for this reaction:
Ca(OH)2 (s) + 2 HCl (g )  CaCl2 (s) + 2 H2O (g)
Enthalpies of formation: Ca(OH)2 (s) = - 986.2 kJ mol-1;
241.8 kJ mol-1;
HCl (g ) = -92.30 kJ mol-1; CaCl2 (s) = -795.8 kJ mol-1; H2O (g) -
Hreaction = [ 2(-241.8) + (-795.8)] -- [ (-986.2) + 2(-92.3)],kJ
= (-483.6 -795.8) -(986.2-184.6)
= (-1279.4) -(-1170.8)
= -108.6 kJ
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1.
Use the standard enthalpy table to find the enthalpy of reaction for the following fuels
A. Ethanol
C2H5OH (g) + 3 O2 (g)  2 CO2 (g) + 3 H2O(g)
B.
Propane
C.
Benzene
D.
2
C3H8 (g)
+
C6H6 (l)
5 O2 (g)
+ 7.5 O2 (g)
 3 CO2 (g) +
 6 CO2 (g) +
4 H2O (g)
3 H2O (g)
On a per gram basis which of the above fuels produces the most heat?
Calcium carbonate decomposes on heating as follows
CaCO3 (s)  CaO (s) + CO2 (g)
Calculate the enthalpy of for this reaction
3
Hydrogen peroxide decomposes to water and oxygen according to the following reaction
2 H2O2 (l) 
2 H2O (l) + O2 (g)
Use this information and the enthalpies of formation from the table to calculate the enthalpy of this
reaction.
4
Calcium carbide, CaC2, reacts with water to form acetylene C2H2, and Ca(OH)2
CaC2 (s) + 2 H2O (l)  Ca(OH)2 (s) + C2H2 (g)
∆Hreaction = -127.2 kJ
Use this information and the enthalpies of formation from the table to calculate the enthalpy of formation
or calcium carbide.
Enthalpy of Formation
Calculate the enthalpy change for each of the following reactions. Label the reaction as
exothermic or endothermic.
1) NH4NO3(s)  NH4
+
(aq)
+ NO3- (aq)
–132 kJ
2+
2) CaCl2(s)  Ca
(aq)
–542.8 kJ
28.1 kJ Endothermic
–205 kJ
+ 2 Cl (aq)
–82.2 kJ Exothermic
–167.2 kJ
3) 2 H2(g) + O2(g)  2 H2O(l)
–571.6 kJ Exothermic
4) n-C4H10(g) + 13/2 O2(g)  4CO2(g) + 5H2O(l)
-2878 kJExothermic
Enthalpy of Formation
1) Calculate the enthalpy change for each of the following reactions. Label the reaction as exothermic or
endothermic.
a) C6H12O6(s)  2C2H5OH(l) + 2CO2(g)
–74.4 kJ Exothermic
b) n-C4H10(g) + 13/2 O2(g)  4CO2(g) + 5H2O(l)
–2878 kJ Exothermic
2) Calculate the ΔHº for the reaction as well as when a 28 g sample of glucose, C 6H12O6(s), is burned to
form CO2(g) and H2O(l) in a reaction at constant pressure and std pressure.
ΔHrxn = –2808 kJ –434.4 kJ Exothermic
3) Calculate the ΔHº for the reaction as well as the amount of heat evolved or absorbed when a 0.1045 g
sample of benzene, C6H6(l), is burned in excess oxygen to form CO2(g) and H2O(l) in a reaction at
constant pressure and std. pressure.
ΔHrxn = –3267.58 kJ –4.38 kJ Exothermic
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