Download Sec 6.2 Enthalpy - Okemos Public Schools

Chapter 17
Thermochemistry
Objectives:
17-1
The flow of Energy –Heat and Work
1. Define thermochemistry.
2. Understand what exothermic and endothermic reactions are.
3. Be familiar with the units energy is measured in and convert between them.
4. Understand the concept of specific heat and be able to calculate it.
17-2
Measuring and Expressing Enthalpy Changes
5. Understand the meaning of Enthalpy, ΔH. Calculate ΔH using ΔH = m ∙ c ∙ ΔT
6. Include a heat term in a thermochemical equation given ΔH and vice versa.
7. Given the value of ΔH for a reaction, calculate the heat involved for any amount of reactant used or
product produced.
17-3
Heat in Changes of State
8. Use heat of fusion and heat of vaporization values to calculate energy involved in phase changes.
9. Understand what heat of solution is.
17-4
Calculating Heats of Reaction
10. Use Hess’s Law to find ΔH for a reaction
11. Use Heats of formation to solve for ΔH.
Types of math problems covered in chapter 17
1) Convert between Joules and Calories
2) Heat added or removed from one substance.
3) Hot and cold things together
Q = m ∙C∙ ΔT
(-1) m ∙ C ∙ ΔT = m ∙ C ∙ ΔT
4) Given a thermochemical equation and a mass of a reactant, find energy released or absorbed.
Convert the mass to moles then covert that to energy using the mole ratio.
5) Given a mass of water, calculate the energy absorbed or released for a phase change.
Convert the mass to moles then multiply by the molar heat of fusion or vaporization.
6) Heat of solution – when a solid dissolves.
i. Convert grams to moles
ii. Calculate Q
iii. Divide Q by moles
7) Hess’ Law: Calculate ΔH given a few reactions and their ΔH values.
8) Calculate ΔH given heats of formation
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Vocabulary
Joule/calorie
Specific Heat
Heats of:
Thermochemistry
Heat capacity
Calorimetry
Enthalpy
Thermochemical Equation
a. Reaction
b. Fusion/solidification
c. Vaporization/Condensation
d. Solution
e. Formation
f. Hess’ Law
Concepts
Heat flows from warm objects to cold ones
Exothermic
a.
b.
c.
gives off heat
feels warm
has a negative ΔH
d. would be included in the equation on the right side
Endothermic
a.
absorbs heat
b. feels cold
c.
has a positive ΔH
d. would be included in the equation on the left side
Be able to take a ΔH value and put it on the correct side of the equation and vice versa.
Know that the temperature remains constant when a phase change occurs.
Observe and equation to determine if it is for the standard heat of formation.
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NOTES Chapter 17 Thermochemistry
17.1 The Flow of Energy
Thermochemistry – the study of energy changes that occur during chemical reactions and changes in state.
Heat – q
Energy that transfers from one object to another because of a temperature difference.
Heat flows from warmer objects to cooler ones.
System
vs.
Surroundings
Conservation of Energy
Endothermic
q is positive
Exothermic
q is negative
Units! 1 Calorie =
1 kilocalorie
=
1000 calories
Joule is the SI unit
4.184 J = 1 calorie
Heat Capacity – the amount of heat needed to raise the temperature of an object exactly 1 °C – It depends
on the mass of the object and the chemical composition.
Consider:
A 2 lb weight vs. a 10 lb weight
1 ft diameter puddle vs. 1 ft diameter manhole cover
Specific Heat – the amount of heat it takes to raise 1 g of a material 1°C
17.2 Measuring and Expressing Enthalpy Changes
Calorimetry – the precise measurement of the heat flow into or out of a system
Constant Pressure experiments – open to the atmosphere – often done in foam cups because they are such
good insulators.
At constant pressure, the heat that goes in or out can also be defined as the Enthalpy Change. ΔH
So for all of our experiments, q = ΔH
You can measure q ( or ΔH) using the formula:
q = CmΔT
m = mass, C = specific heat
Examples:
1) How many Joules of energy would be required to raise a 10.0 gram piece of iron from 21.0°C to
100.°C? The specific heat of iron is 0.460 J/g∙°C
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2) What is the specific heat of mercury if it takes 900. Joules of energy to raise the temperature of
161 g mercury by 40.0°C?
3) If 1000. Joules of energy are added to a 5.00 g piece of iron that has an original temperature of
20.0°C, what will be the final temperature?
Now how about hot and cold things mixed together?
4) What will the final temperature be if 30.0 mL of water at 20.0°C is mixed with 70.0 mL of water at
95.0°C?
Sec 6.2 Enthalpy
Enthalpy can be defined as the heat content of a given substance in a specific condition. The letter H is used to
represent enthalpy.
The symbol H stands for Hproducts – Hreactants . It represents the amount of heat released or absorbed in a
chemical reaction that takes place at constant pressure.
In nature, substances always strive to attain the greatest stability, which must mean that there is a natural drive
toward minimum energy or towards minimum enthalpy. Note the two situations below:
Situation 1
Situation 2
Enthalpy Hreactants
Enthalpy
H
Hproducts
H
Hreactants
Hproducts
Reaction coordinate
reaction coordinate
H is ( +, -)
H is ( +, -)
H is ( favorable, unfavorable)
H is ( favorable, unfavorable)
enthalpy is (increasing, decreasing)
enthalpy is (increasing, decreasing)
heat (is given off, absorbed from)surroundings
heat (is given off, absorbed from) surroundings
reaction is (exothermic, endothermic)
reaction is (exothermic, endothermic)
We can indicate the magnitude of the energy change during chemical reactions in two ways:
1. heat as a reaction term
2. heat exchange shown in ∆H notation
H2(g) + ½ O2(g) →H2O (g) + 57.8 Kcal
H2(g) + ½ O2(g) →H2O (g) ∆H = -57.8 Kcal
Of course as the quantities of reactants and products change ∆H does also. Example
2 H2(g) + O2(g) → 2 H2O (g) ∆H = -115 Kcal
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Consider the reaction C(graphite) + H2O (g) →CO(g) + H2 (g) ∆H = + 31.4 Kcal
The reaction is (exothermic, endothermic). The energy of the reactants is (greater, less) than that of the products.
As this reaction proceeds from left of right heat must be _______________________________.
Rewrite the reaction and indicate the energy change as a reaction term.
C(graphite) + H2O (g) →CO(g) + H2 (g)
Sketch an energy diagram for the reaction:
C(graphite) + H2O (g) →CO(g) + H2 (g) ∆H = + 31.4 Kcal
Fill in the table:
Energy of products lower than that of
reactants
Energy of products higher than that of
reactants
Exothermic or endothermic
Stability increasing or
decreasing
Enthalpy increasing or
decreasing
Heat term on left or right in equation
∆H= + or ∆H = -
Thermochemical equation – an equation that includes a heat term.
1.
CH4 + 2 O2 → CO2 + 2H2O + 890 kJ
Endothermic / Exothermic
ΔH =
2.
241.8 kJ + 2H2O →
2H2 + O2
Endothermic / Exothermic
ΔH =
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3.
4.
Calculate the amount of heat needed to decompose 200. g of NaHCO3.
2NaHCO3 + 129 kJ → Na2CO3 + H2O + CO2
2 CO (g) + C (s) → C3O2 (g)
ΔH = 127.3 kJ
How much heat is absorbed when 14.0 grams carbon monoxide react?
17.3 Heat in Changes of State
Temperature remains constant during phase changes.
Molar Heat of Fusion = energy required to melt one mole
(Molar Heat of Solidification) = energy removed to freeze one mole.
ΔHfus for water is 6.01 kJ/mol
Example 1 : How many grams of ice at 0°C will melt if 2.25 kJ of heat are added?
Example 2: How much energy is required to melt 50.0 grams of ice?
The process is the same for boiling/condensing.
Molar Heat of Vaporization = Energy required to vaporize (boil off) 1 mole.
(Molar Heat of Condensation) = Energy removed to condense 1 mol
ΔHvap for water is 40.7 kJ/mol
Example 1: How much heat is absorbed when 24.8 g of H 2O at 100.°C is converted to steam at 100.°C?
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Example 2 : How many grams of water could be vaporized if 100. kJ of energy vaporized water?
Example 3 : How much energy will be needed to change 50.0 grams of water from 80.0°C to steam?
Heat of Solution = the energy released or absorbed when a substance is dissolved
Example 1:
How much heat is involved when 100. g of NaOH is dissolved in water if ΔHsol for NaOH is 445.1 kJ/mol. Is it absorbed or released?
Example 2:
Determine ΔHsol for CaCl2 if 100. grams of it are added to 300. g of water and the
temperature goes from 20.0°C to 64.6°C.
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17.4
Calculating Heats of Reaction
Hess’ Law
You will be given 2 or more equations with their ΔH values. You will use them to calculate the ΔH for
another reaction.
You will try to make the given equations “add up” to the one with the missing ΔH. You can manipulate the
equations to do this. There are two manipulations you can do:
1) You can reverse an equation. You must flip the sign of ΔH if you do this.
2) You can multiple the equation through by a number. You must multiply the ΔH by this same
number.
Things on opposite sides of the arrow will cancel out if they are the same.
When the given equations add up to the one you want, their ΔH values will add up to the one you want.
Hess’s Law Examples
#1
H2S + 3/2 O2 → H2O + SO2
ΔH = -563 kJ
CS2 + 3O2 →
ΔH = -1075 kJ
CO2 + 2SO2
CS2 + 2H2O →
CO2 + 2H2S
ΔH = ?????
#2
OF2 + H2O → O2 + 2HF
SF4 + 2H2O → SO2 + 4HF
S + O2 → SO2
ΔH = -277 kJ
ΔH = -828 kJ
ΔH = -297 kJ
2S + 2OF2 →
ΔH = ?????
SO2 + SF4
Standard Heats of Formation = Another way to find ΔH for a reaction
Standard heat of Formation is the change in enthalpy that accompanies the formation of one mole of a
compound from its elements in their standard states.
Like this:
Na(s) + ½ Cl2(g) → NaCl(s)
But not this:
Na(g) + Cl2(g) →
Na(s) +
2 NaCl(s)
½ Cl2(l) →
NaCl(s)
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The heat of formation for an element in its standard state is 0.
ΔHrxn = Σ np ΔHf° (products) – Σ nr ΔHf° (reactants)
n = # of moles
Example: Calculate the value of ΔH° for the following reaction using the given data.
CH4 (g) + 4 Cl2 (g) → CCl4 (l) + 4 HCl (g)
ΔH = ?
ΔHf of CH4 (g) = -74.80 kJ/mol
ΔHf of CCl4 (l) = -139.3 kJ/mol
ΔHf of HCl (g) = -92.30 kJ/mol
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1.
Chapter 17 Practice
Thermochemistry is the study of __________ changes that occur during chemical reactions and
changes in state.
2.
Heat always flows from a __________ object to a_________ object.
3. An ___________________ process absorbs energy from the surroundings.
4. A ____________________ process releases energy into the surroundings.
5. An exothermic reaction in a test tube would feel _______ if you touched it whereas an
endothermic reaction would feel __________.
6. Energy is measured in the following units: ___________________________________________
7. 23 J = ? calories
8. The specific heat of a substance is the amount of ______________ required to raise 1
_________ of a substance _________°C.
9. Enthalpy is: _________________________________________________________________
10. How much heat is released when 50.0g water at 50.0°C is cooled to 30.0°C?
11. How much heat is needed to raise 65.0g of water from 22.0°C to 45.0°C?
12. What is the temperature change for 25.0 g water if 2000. J of energy is absorbed?
13. What will be the new temperature of an 80.0g sample of water that starts out at 25.0°C and
absorbs 16,000. J of energy?
14. What is the specific heat of Aluminum if a 10.0 g sample is warmed 20.0 °C and absorbs 178J of
energy?
15. If a 15.0 g piece of aluminum is heated from 50.0 °C to 85.0 °C, how much energy did it absorb?
16. A 5.00 g piece of iron with a specific heat of 0.450 J/°C · g is at room temperature (25.0°C). If
250. J of energy are removed, what will be the new temperature?
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17. If 20.0 g of water at 20.0 °C are mixed with 45.0 g of water at 67.0 °C, what will be the final
temperature of the mixture?
18. If a 12.0 g sample of iron at 89.0°C is put in a 50.0g sample of water at 25.0°C, what will be the
final temperature of the water and iron mixture?
19. What is the specific heat of Carbon if a 12.0g sample at 100.°C is put into a 28.5g sample of water
at 25.0°C and the final temperature is 30.0°C?
20. Iron (30.0g) at 95.0°C is added to water at 15.0°C and the final temperature of the mixture is
18.0°C. What was the mass of the water?
21. Add the heat value to the correct side of the equation:
a)
A
b)
+
G
C →
→
H
D
+
+
F
J
ΔH = -345 kJ
ΔH = + 24.3 kJ
22. Determine ΔH from the given equations:
a)
b)
K
+
L →
30.5 J
+ R
M
+ 450 J
ΔH = ___________
→
S +T
ΔH = ___________
23. How much heat will be released when 6.44 g of sulfur reacts with excess O 2 according to the
following equation?
2S + 3O2 → 2SO3
∆H = -791.4 kJ
24. How much heat will be released when 14.9 g of ammonia reacts with excess O 2 according tio the
following equation? 4NH3 + 5 O2 → 4NO + 6H2O ∆H = -1170 kJ
25. How much heat will be absorbed when 28.0 g of NO is formed according to the following equation?
N2 + O2 → 2NO
∆H = 180 kJ
26. How much heat will be released when 1.48 g of chlorine reacts with excess phosphorous according
to the following equation?
2P + 5Cl2 → 2PCl5 ∆H = -886 kJ
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27. Determine the amount of heat needed to vaporize 250. g of H 2O at 100. C
28. Calculate the heat needed to convert 40. g of ice at 0.0°C to water at 0.0°C.
29. Calculate the total energy needed to take 10.0 grams of water from 15.0°C to steam at 100.°C.
30. How much heat in (kJ) is released when 0.677 mol NaOH is dissolved in water? ΔH sol = -445.1
kJ/mol
31. When a 12.8 gram sample of KCl dissolves in 75.0 g of water in a calorimeter, the temperature
drops from 31.0° C to 21.6°. Calculate ∆H for the process.
KCl(s) → K+(aq) + Cl-(aq)
32. When a 25.7 gram sample of NaI dissolves in 80.0 g of water in a calorimeter, the temperature
rises from 20.5° C to 24.4°. Calculate ∆H for the process.
NaI(s) → Na+(aq) + I-(aq)
33. When a 16.9 gram sample of NaOH dissolves in 70.0 g of water in a calorimeter, the temperature
rises from 22.4° C to 86.6°. Calculate ∆H for the process.
NaOH(s) → Na+(aq) + OH-(aq)
34. From the following enthalpy changes,
C + ½ O2 → CO
∆H = -110.5 kJ
CO + ½ O2 → CO2
∆H = -283.0 kJ
Calculate the value of ∆H for the reaction: C + O 2 → CO2
35. From the following enthalpy changes,
2P + 3 Cl2 → 2PCl3
∆H = -640. kJ
2P + 5 Cl2 → 2PCl5
∆H = -886 kJ
Calculate the value of ∆H for the reaction: PCl3 + Cl2 → PCl5
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36. From the following enthalpy changes,
C2H5OH + 3 O2 → 2CO2 + 3H2O
∆H = -1234.7 kJ
CH3OCH3 + 3 O2 → 2CO2 + 3H2O
∆H = -1328.3 kJ
Calculate the value of ∆H for the reaction: C2H5OH → CH3OCH3
37. From the following enthalpy changes,
Cu + Cl2 → CuCl2
∆H = -206 kJ
2Cu + Cl2 → 2CuCl
∆H = -136 kJ
Calculate the value of ∆H for the reaction: CuCl2 + Cu → 2CuCl
38. From the following enthalpy changes,
H2 + F2 → 2HF
∆H = -542.2 kJ
2H2 + O2 → 2H2O
∆H = -571.6 kJ
Calculate the value of ∆H for the reaction: 2F2 + 2H2O → 4HF + O2
39. From the following enthalpy changes,
Xe + F2 → XeF2
∆H = -123 kJ
Xe + 2F2 → XeF4
∆H = -262 kJ
Calculate the value of ∆H for the reaction: XeF2 + F2 → XeF4
40. From the following enthalpy changes,
4NH3 + 5 O2 → 4NO + 6H2O
∆H = -1170 kJ
4NH3 + 3 O2 → 2N2 + 6H2O
∆H = -1530 kJ
Calculate the value of ∆H for the reaction: N 2 + O2 → 2NO
41. From the following enthalpy changes,
2Al + 3/2 O2 → Al2O3
∆H = -1601 kJ
2Fe + 3/2 O2 → Fe2O3
∆H = -821 kJ
Calculate the value of ∆H for the reaction: 2Al + Fe2O3 → 2Fe + Al2O3
42. From the following enthalpy changes,
H2O2 → H2O + ½ O2
∆H = -94.6 kJ
H2 + ½ O2 → H2O
∆H = -286.0 kJ
Calculate the value of ∆H for the reaction: H2 + H2O2 → 2H2O
43. From the following enthalpy changes,
C + O2 → CO2
∆H
H2 + ½ O2 → H2O
∆H
2C2H2 + 5O2 → 4CO2 + 2H2O
∆H
Calculate the value of ∆H for the reaction: 2C +
= -393.5 kJ
= -285.8 kJ
= -2598.8 kJ
H 2 → C2H2
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44. Which reaction shows the standard enthalpy of formation? For those that don’t what changes
would have to be made?
a.
2 Na (s) + ½ O2 (g) → Na2O (s)
b. 2 K (l) + Cl2 (g) → 2 KCl (s)
c.
C6H12O6 (s) → 6 C (graphite) + 6 H2 (g) + 3 O2 (g)
45. Calculate the enthalpy change for the combustion of 1 mol of ethanol.
C2H5OH (l) + 3 O2 (g) → 2 CO2 (g) + 3 H2O (l)
Given: ΔHf° C2H5OH (l) = -277.7 kJ/mol,
ΔHf° CO2 (g) = -393.5 kJ/mol
ΔHf° H2O (l) = -285.8 kJ/mol
and
46. Calculate ΔH for the following reaction: CaCO3(s) → CaO(s) + CO2(g)
ΔHf° CaCO3 = -1207.0 kJ/mol, ΔHf° CaO = -635.1 kJ/mol, ΔHf° CO2 = -393.5 kJ/mol
Page | 14
Page | 15
Extra Chapter 17 practice on Sections 1-3
1. What is the difference between heat capacity and specific heat?
2. Joules and Calories are both units of ______________.
3. Use the given facts to determine the number of Joules in the awesome lime Oiko yogurt I had for lunch
yesterday. It had 160 Calories.
1 calorie = 4.184 J and 1000 calories = 1 Calorie
4. What is ΔH?
5. Given the reaction: CaCO3 + CO2 →CaC2 + 2 ½ O2
ΔH = 1538 kJ
a. What side would you write the heat term on if you included it in the reaction?
b. If this reaction was going on in a beaker you were holding, how would it feel?
c. Is the reaction endothermic or exothermic?
d. If 100.0 g of CaCO3 reacted with excess CO2, how much heat would be released or absorbed?
e. Fill in the general shape of the enthalpy diagram. Draw an
arrow with the label ΔH.
H
6. Which has more chemical potential energy, gasoline or water?
7. Why does a puddle heat up in the sun more slowly than a manhole
cover of the same size?
Reaction progression
8. Which diagram is correct?
a.
cold
cold
hot
b.
hot
cold
d
9. When 435 J of heat is added to 3.4 g of olive oil at 21 ⁰C , the temperature increases to 85. What is
the specific heat of olive oil?
10. Write and balance the equation for the following reaction: Magnesium metal burns in air. If it is
known that 1 mole of Magnesium burns to produce 602 kJ of energy, figure out the correct number to
add to your equation. Be sure to put in on the correct side!
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11. How are “molar heat of fusion” and “molar heat of solidification” similar and different?
12. How many grams of ice could be melted if 4.5 kJ of heat are added?
13. How are “molar heat of condensation” and “molar heat of vaporization” similar and different?
14. How much heat is absorbed when 24.8 g of liquid water at 100 ºC and standard pressure is converted to
steam at the same temperature?
15. How much heat is evolved when 120.0 g of CaCl 2 dissolve in water? ΔHsoln = -82.8 kJ/mol
16. When a 120. g sample of a solid with a molar mass of 54 g/mol dissolves in 75.0 g water, the
temperature goes up from 22⁰ C to 29⁰ C. Calculate ΔH soln by following these steps:
a. Use Q = m ∙ c ∙ ΔT to find the about of Joules gained by the water. (Remember to use the mass
of the entire solution!)
b. Calculate the moles of the solid using it’s mass and molar mass.
c. CalculateΔHsoln in kJ/mol by diving the energy in part “a” by the moles in part”b”. Remember to
put the correct sign on it!
17. Identify each enthalpy change by name and classify each change as exo or endothermic.
a. 1mol C3H8(l) → 1 mol C3H8(g)
c. 1 mole NH3(g) → 1 mole NH3(l)
b. 1mol Hg(l) → 1 mole Hg(s)
d. 1 mole NaCl(s) + 3.88 kJ/mol →1 mole NaCl(aq)
e. 1 mole NaCl(s) → NaCl(l)
Page | 17
Thermochemistry Practice Problems
Name: _________________
1.
Calculate the quantity of heat gained or lost in the following changes:
The heat of fusion of water is 6.01 kJ / mole.
The heat of vaporization of water is 40.67 kJ / mole.
a. 3.5 mol of water freezes at 0.0C.
b. 0.44 mol of steam condenses at 100.C.
2.
What is H for this reaction:
PbCl2 + Cl2  PbCl4
Pb + 2Cl2  PbCl4
Pb + Cl2  PbCl2
H=-329.2 kJ
H= -359.4 kJ
3.
A sample of water with a mass of 40.0g is originally at 25.0C and has a final temperature of
32.0°C. How much heat does the water absorb?
4.
When 55.0g of water at 67.0°C is added to a beaker of water at 44.0°C the final temperature ends
up being 58.0°C. What was the mass of water in the beaker?
5.
An orange contains 445 KJ of energy. What mass of water could this same amount of energy raise
from 25.0C to 100.C?
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6.
Calculate the amount of heat given off when 4.79g of C 2H4 reacts with excess oxygen.
C2H4 + 3O2  2CO2 +2H2O
ΔH= -1.39 x 103 kJ/mol
7.
If a 70.0 g piece of aluminum (specific heat 0.900 J/g°C) at 80.0°C is added to 200.g of water at
35.0°C what would be the final temperature?
8.
When a 16.0g sample of NaCl dissolves in 60.0g of water the temperature rises from 45.0°C to
49.0°C. Calculate the ΔHsoln for the process.
9.
What is the standard heat of formation (ΔH°) for the decomposition of hydrogen peroxide?
2H2O2  2H2O + O2
ΔHf° H2O2 = -187.8 kJ/mol
ΔHf° H2O = -285.8 kJ/mol
10.
What will be the new temperature of a 78.0g sample of water that starts out at 60.0°C and loses
7,600.J of energy?
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Thermochemistry Review
Name: __________________
1. What is the standard heat of formation (∆H˚) for the following?
4NH3 + 5O2  4NO + 6H2O
∆Hf˚ NH3 = -46.19kJ/mol
∆Hf˚ NO = 90.37kJ/mol
∆Hf˚ H2O = -241.8kJ/mol
2. Calculate the amount of heat given off when 12.5g of ethanol burns.
C2H5OH +3O2  2CO2 +3H2O
∆H = -1368 kJ/mol
3. How much heat does a 32g piece of silver absorb when it is heated from 25˚C to 80.˚? Specific
heat silver = 0.24 J/g˚C
4. Calculate the quantity of heat gained or lost in the following changes:
Heat of fusion water = 6.01 kJ/mol
Heat of vaporization water = 40.67 kJ/mol
a) 63.7g of ice melts at 0˚C.
b)
188g of steam condenses at 100˚C
5. What was the original temperature of a 310.g sample of water that has a final recorded
temperature of 33.0˚C and lost 5,887 J of energy?
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6. What is the ∆H for the reaction:
3
Ca + C + O2  CaCO3
2
Ca + 2C  CaC2
CO2  C + O2
1
CaCO3 + CO2  CaC2 + 2 O2
2
H = -62.8 kJ
H = 393.5 kJ
H = 1538 kJ
7. If a 17g piece of iron (specific heat = 0.46 J/g˚C) at 22˚C is dropped into a flask containing 95g of
water at 98˚C what would the final temperature be?
8. When a 34g sample of MgCl2 is dissolved in 101g of water the temperature drops from 73C to
55C. Calculate the Hsoln for this reaction.
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