Sophomore Chemistry
Unit 10 – Stoichiometry
Name:
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Period:
Unit 11: Energetics
Driving Questions


Why do some reactions produce heat while other reactions require it?
How can energy of a reaction be calculated to design real-world
products?
Connections to Past/ Future Units


Experimental design skills will be used to build a procedure
Problem solving skills will be used to solve multi-step energy
problems including Stoichiometry
Objectives








SWBAT differentiate between exothermic and endothermic reactions
in equations and energy diagrams
SWBAT calculate the enthalpy of a reaction using stoichiometry
SWBAT calculate the enthalpy of a reaction using bond energies
SWBAT model chemical reactions to visibly demonstrate changes in bond energies
SWBAT use and design a lab to investigate energy using a calorimeter
SWBAT calculate the energy necessary to change the temperature and phase of a substance
SWBAT interpret heating and cooling curves during phase change
SWBAT use Hess’s Law to calculate the enthalpy of a reaction
Essential Vocabulary
thermodynamics
temperature
system
surroundings
endothermic
exothermic
activation energy
enthalpy
bond energy
specific heat capacity
heat of fusion
heat of vaporization
calorimeter
Hess’s Law
PERSONAL OBJECTIVE: Looking at the objectives above, what more do you want to learn this unit?
REFLECTION: Did you accomplish your personal objective? What further questions do you have about this unit?
Sophomore Chemistry
Unit 11 - Energetics
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When you eat an orange, the sugar it contains reacts in your body with oxygen to ultimately form CO2 and H2O. But
more occurs than merely this conversion of chemicals – energy is released. The food you eat is the fuel that your body
uses to operate your muscles and to maintain proper body temperature. This example illustrates a general point:
chemical reactions involve changes in energy. Some reactions, like the oxidation of sugar, give off energy. Others, like
the electrolysis of water, take in energy. At the present time, over 90 percent of the energy produced in our society
comes from chemical reactions, principally from the combustion of coal, petroleum products, and natural gas.
The study of energy and its transformations is known as thermodynamics. This area of study began during the Industrial
Revolution as the relationships among heat, work, and the energy content of fuels were studied in an effort to maximize
the performance of steam engines. Thermodynamics is important not only to chemistry but to other areas of science
and to engineering as well. It touches our daily lives as we use energy for manufacturing, travel, and communications.
Thermodynamics relates to such diverse topics as the metabolism of foods, the operation of batteries, and the design of
foods.
To begin our study of thermodynamics, we should examine the ways we measure energy in the lab. By now, you are
undoubtedly aware of temperature. Temperature is a measurement of the average kinetic energy of atoms or
molecules in a sample. The units for measuring temperature are degrees Celsius, or Kelvins. Hopefully, you remember
the conversions between these two scales:
°C + 273 = K
K − 273 = °C
It is important to recognize that temperature is not a direct measurement of heat. Heat is energy that is transferred
from one body to another. Temperature determines the direction of this heat transfer.
To study energy changes we focus our attention on a limited and well-defined part of the universe. The portion that we
single out for study is called the system. Everything else outside the system is called the surroundings. A reaction that
results in the evolution of heat is exothermic – energy flows out of the system and into its surroundings. Combustion
reactions are exothermic. Reactions that absorb heat from their surroundings are endothermic. The melting of ice is an
endothermic process.
EXAMPLE: You place a small strip of magnesium into a test tube containing HCl. When you pick up the test tube after it
has reacted, your hand feels hot.
System =
Surroundings =
Explain how the heat is traveling in this example
Is this reaction endothermic or exothermic?
EXAMPLE: You decide that you want to make scrambled eggs for breakfast. After turning on the stove, you pour in your
egg mixture and notice the egg whites start to solidify.
System =
Surroundings =
Explain how the heat is traveling in this example
Is this reaction endothermic or exothermic?
Sophomore Chemistry
Unit 11 - Energetics
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We can also graphically represent endothermic and exothermic reactions with a graph of Potential Energy vs. Reaction
time. Molecules contain potential chemical energy in the bonds between atoms.
EXOTHERMIC REACTIONS:
A=
B=
C=
D=
∆H =
1. Which of the letters a–f in the diagram represents the
potential energy of the products? _______
2. Which letter indicates the potential energy of the
activated complex? ________
3. Which letter indicates the potential energy of the
reactants? ________
4. Which letter indicates the activation energy? _____
5. Which letter indicates the heat of reaction? ______
6. Is the reaction exothermic or endothermic? ______
7. Which letter indicates the activation energy of the
reverse reaction? ________
8. Which letter indicates the heat of reaction of the
reverse reaction? ________
9. Is the reverse reaction exothermic or endothermic? ___
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Unit 11 - Energetics
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Draw an energy diagram for a reaction. (label the axes)

Potential energy of reactants = 350 KJ/mole

Activation energy = 100 KJ/mole

Potential energy of products = 250 KJ/mole
1. The potential energy of the reactants of the forward reaction is about
______ kilojoules.
2. The potential energy of the products of the forward reaction is about
_______kilojoules.
3. The potential energy of the activated complex of the forward reaction is
about ______ kilojoules.
4. The activation energy of the forward reaction is about ______ kilojoules.
5. The heat of reaction (ΔH) of the forward reaction is about ______ kilojoules.
6. The forward reaction is _______________ (endothermic or exothermic).
7. The potential energy of the reactants of the reverse reaction is about
________ kilojoules.
8. The potential energy of the products of the reverse reaction is about
_______ kilojoules.
9. The potential energy of the activated complex of the reverse reaction is
about _______kilojoules.
10. The activation energy of the reverse reaction is about _______ kilojoules.
11. The heat of reaction (ΔH) of the reverse reaction is about _______
kilojoules.
12. The reverse reaction is __________________ (endothermic or
exothermic).
∆H = ENTHALPY =
+∆H = ENDOTHERMIC = Heat is a reactant (Heat needs to be put in for a reaction to occur)
-∆H = EXOTHERMIC = Heat is a product (Heat comes out when the reaction to occurs)
Rewrite the following reactions with ∆H as a product or reactant.
1. H2  2H
∆H = 436 kJ/mol
Exothermic or Endothermic?
Sophomore Chemistry
2. 2F  F2
Unit 11 - Energetics
∆H = - 158 kJ/mol
3. H2(g) + F2(g)  2 HF
∆H = -542 kJ/equation
4. C2H5OH(l)  2 C(s) + 3 H2(l) + 0.5 O2(g) ∆H = 228 kJ/mol
STOICHIOMETRY OF ENERGY
1) 2S + 3O2  2SO3
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Exothermic or Endothermic?
Exothermic or Endothermic?
Exothermic or Endothermic?
∆Ho = -791.4 kJ
a. How much heat will be released when 6.44 g of sulfur reacts?
b. Is this reaction endothermic or exothermic?
2) C + O2  CO2
∆Ho = -393.5 kJ
a. How many grams of carbon would be required to liberate 284 kJ of thermal energy?
b. Draw and energy diagram for this reaction.
3) H2 + Br2 + 72.80 kJ  2HBr
a. How much heat will be absorbed when 38.2 g of bromine reacts?
b. Is this reaction endothermic or exothermic?
4) 2P + 5Cl2  2PCl5
∆Ho = -886 kJ
a. Determine the amount of heat produced if 1.98 L of chlorine gas are reacted at STP.
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Unit 11 - Energetics
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b. Rewrite the equation with ∆H as a product or reactant.
5) C2H5OH + 3O2  2CO2 + 3H2O
∆Ho = -1366.7 kJ
a. If this reaction gives off 951 kJ of heat, how many molecules of water will be produced?
b. Will this reaction cause the surroundings to increase or decrease in temperature?
6) N2 + O2 + 180 kJ  2NO
a. How much heat is required to react 13.7 g of nitrogen.
b. Draw an energy diagram for this reaction
7) 3Fe + 2O2  Fe3O4 + 1120.48 kJ
a. How many grams of iron would be required for this reaction to liberate 7985 kJ of thermal energy?
b. Write this equation in ∆H notation.
One can solve for ΔH for any reaction by considering bond energies, as well. The bond energy is the amount of energy
required to break a covalent bond. First, the individual bonds in a structure must be identified (using the Lewis Dot
Structures you learned about in Unit 07). Then, the bond energies of the reactants are added to a final total, and the
bond energies of the products are subtracted. The total will be equal to ΔH of the overall reaction. The values of bond
energies must be found in a table, and the equation for using bond energies to solve for enthalpy of a reaction is:
ΔHreaction = B.E.reactants − B.E.products
Note that the order of reactants and products is the opposite from the equation used for heat of formation. A complete
list of bond energies is found in the table below.
Bond
Bond Energy
(kJ/mole)
Bond
Bond Energy
(kJ/mole)
Bond
Bond Energy
(kJ/mole)
Sophomore Chemistry
Unit 11 - Energetics
C−C
C−H
C−O
C=O
C≡O
307
393
356
736
1074
O−O
O=O
O−H
H−Cl
Cl−Cl
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118
433
464
436
242
C−Cl
H−H
N−H
N≡N
156
446
393
928
Use this information to solve the problems in this section. The first problem has been done for you.
EXAMPLE: Determine the enthalpy for the following reaction using bond energies.
___ CH4 + ___ O2 → ___ CO2 + ___ H2O
Step 1: Balance the equation.
_1_ CH4 + _2_ O2 → _1_ CO2 + _2_ H2O
Step 2: Envision the Lewis Dot Structures for each of the species involved.
Reactants:
H
|
H― C ―H
|
H
Products:
O=O
H―O―H
→
+
O=C=O +
O=O
H―O―H
Step 3: Substitute the correct bonds into the equation ΔHreaction = B.E.reactants − B.E.products. Apply the correct mole ratio.
ΔHreaction = B.E.reactants − B.E.products
ΔHreaction = 4(C―H) + 2(O=O) − 2(C=O) − 4(O―H)
Step 3: Substitute the correct numerical values into the equation.
ΔHreaction = 4∙(393) + 2∙(433) − 2∙(736) − 4∙(464)
Step 4: Mathematically solve the problem.
ΔHreaction = −890 kJ
Use this example to guide you as you attempt the following problems. Also, be sure to check your answers against the
heat of formation problems. The equations are all the same. The answers should be very close to each other as well.
Show all work.
1) Determine the enthalpy for the following reaction using bond energies.
___ C2H6 + ___ O2 → ___ CO2 + ___ H2O
Sophomore Chemistry
Unit 11 - Energetics
2) Determine the enthalpy for the following reaction using bond energies.
___ C3H8 + ___ O2 → ___ CO2 + ___ H2O
3) Determine the enthalpy for the following reaction using bond energies.
___ H2 + ___ O2 → ___ H2O
4) Determine the enthalpy for the following reaction using bond energies.
___ H2O2 → ___ H2O + ___ O2
5) Determine the enthalpy for the following reaction using bond energies.
___ N2 + ___ H2 → ___ NH3
6) Determine the enthalpy for the following reaction using bond energies.
___ H2 + ___ Cl2 → ___ HCl
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Sophomore Chemistry
Unit 11 - Energetics
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7) Determine the enthalpy for the following reaction using bond energies.
___ CCl4 → ___ C + ___ Cl2
If you had the choice to touch 100 °C metal or °C water, what would it be?
What you’ve likely intuited is the concept of specific heat. The specific heat capacity of an object is the amount of
energy required to raise the temperature of one gram of a substance by one degree Celsius. The specific heat of liquid
water is 4.18 J/g°C, while the specific heat of solid copper (the metal used in a typical hot plate heating element) is 0.385
J/g°C. This means that a given amount of energy will be more effective at raising the temperature of copper than an
equivalent sample of water. The relationship between specific heat, energy, mass and temperature is expressed in the
following equation. Use it as you consider the problems below:
q = energy (J)
c = specific heat capacity (J/ g°C)
m = mass (g)
ΔT = change in temperature (°C)
q = c∙m∙ΔT
1) How many joules would it take to raise the temperature of a 50.0 g sample of H2O from 35.0°C to 75.0°C. The
specific heat of liquid water is 4.18 J/g°C. Show your work.
2) How many joules would it take to raise the temperature of a 50.0 g sample of Cu from 35.0°C to 75.0°C. The
specific heat of solid copper is 0.385 J/g°C. Show your work.
3)
Calculate the number of joules released when 75.0
o
o
grams of water are cooled from 100.0 C to 27.5 C
A calorimeter is a device used to measure the quantity of
heat transferred to or from an object. While it sounds very
fancy, in most chemistry labs this consists of a Styrofoam
cup filled with water, and a lid with room for a stirrer and
thermometer. The water will change its temperature
when it gains or loses energy based on the equation from
the previous page. So if the mass of water and the
temperature change of the water in the coffee cup
calorimeter can be measured, the quantity of energy
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Unit 11 - Energetics
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gained or lost by the water can be calculated. The assumption behind the science of calorimetry is that the energy
gained or lost by the water is equal to the energy lost or gained by the object under study. A bomb calorimeter has a
tiny chamber inside of a large volume of water where a certain amount of material can be combusted. Again, all of the
energy from the reaction is transferred to or from the water, and a calculation can be made.
Qreaction = - Qwater
EXAMPLES:
1) When a 25.7 g sample of NaI dissolves in 80.0 g of water in a calorimeter, the temperature rises from 20.5 oC to
24.4 oC. Calculate H for the process.
NaI(s)  Na+(aq) + I-(aq)
H = ?
2) What is the specific heat of silicon if the temperature of a 4.11 g sample of silicon is increased by 3.8 oC when
11.1 J of heat is added?
3) A 2.50 g sample of zinc is heated, then placed in a calorimeter containing 65.0 g of water. Temperature of water
increases from 20.00 oC to 22.50 oC. The specific heat of zinc is 0.390 J/goC. What was the initial temperature of
the zinc metal sample? (final temperatures of zinc and water are the same)
PRACTICE
4) When a 16.9 g sample of NaOH dissolves in 70.0 g of water in a calorimeter, the temperature rises from 22.4 oC
to 86.6oC. Calculate H for the process.
NaOH(s)  Na+(aq) + OH-(aq)
H = ?
5) What is the specific heat of aluminum if the temperature of a 28.4 g sample of aluminum is increased by 8.1 oC
when 207 J of heat is added?
6) A 28.4 g sample of aluminum is heated to 39.4 oC, then is placed in a calorimeter containing 50.0 g of water.
Temperature of water increases from 21.00 oC to 23.00 oC. What is the specific heat of aluminum?
Sophomore Chemistry
Unit 11 - Energetics
Page 11
7) A 2.50 g sample of zinc is heated, then placed in a calorimeter containing 65.0 g of water. Temperature of water
increases from 20.00 oC to 22.50 oC. The specific heat of zinc is 0.390 J/goC. What was the initial temperature of
the zinc metal sample? (final temperatures of zinc and water are the same)
8) A 13.5 g sample of gold is heated, then placed in a calorimeter containing 60.0 g of water. Temperature of
water increases from 19.00 oC to 20.00 oC. The specific heat of gold is 0.130 J/goC. What was the initial
temperature of the gold metal sample?
o
o
9) Calculate the number of joules released when 75.0 grams of water are cooled from 100.0 C to 27.5 C.
o
10) The specific heat of gold is 0.128 J/gC . How much heat would be needed to warm 250.0 grams of gold from
o
o
25.0 C to 100.0 C?
11) A 100.0 gram sample of water at 50.0oC is mixed with a 50.00 gram sample of water at 20.0oC. What is the final
temperature of the 150.0 grams of water?
12) Calculate the specific heat of a metal if 2.36 x 102 grams of it at 99.5oC is added to 125.0 mL of water at 22.0oC.
The final temperature of the system is 25.4oC.
13) A lump of chromium (Cr) has a mass of 95.3 grams and a temperature of 90.5oC. It is placed into a calorimeter
with 75.2 mL of water at 20.5oC. After stirring, the final temperature of the water, Cr metal, and calorimeter is
28.6oC. What is the specific heat of Cr metal? (Hint: the final temperature of BOTH the metal and the water is
28.6oC)
14) When a 12.8 g sample of KCl dissolves in 75.0 g of water in a calorimeter, the temperature drops from 31.0 oC to
21.6 oC. Calculate H for the process.
KCl(s)  K+(aq) + Cl-(aq)
H = ?
Sophomore Chemistry
Unit 11 - Energetics
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You can now calculate the heat necessary to change the temperature of a substance in one phase, but what if that
temperature change occurs over different phases?
The following labeled diagram is called a heating curve. As heat is added (and we move to the right in the diagram), the
way matter is put together changes.
What is happening in the
two areas that are
horizontal? Why does the
temperature not increase?
The amount of heat required to turn one gram of ice into liquid water at 0°C is called the heat of fusion (ΔHfus). The
heat of fusion of water is 334 J/g. How many joules are necessary to turn 50.0 g ice into water?
q = m ∙ ΔHfus =
The amount of heat required to turn one gram of liquid water into steam at 100°C is called the heat of vaporization
(ΔHvap). The heat of vaporization of water is 2260 J/g. As in most substances, the heat of vaporization of water is much
greater than the heat of fusion. Why do you think this is so?
How many joules are necessary to turn 50.0 g liquid water into steam? Input correct numbers and solve.
q = m ∙ ΔHvap =
Alright, now that we have spent some time learning how a heating curve is put together, we have to figure out how to
solve problems using one. Consider the following problem:
How much energy is needed to raise the temperature of 35.0 g of water
from 75°C to 110°C?
First, estimate where you start and end on the heat curve to the right.
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Unit 11 - Energetics
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Then, list the steps that it takes to raise the temperature of the water and select the correct equation.
1.
2.
3.
Finally, add all your values together to achieve your final answer.
This is the method by which you will solve heating curve problems. Problems may have as few as one step or as many as
five, depending on the identity of the substance undergoing heating, and the range of temperatures involved.
Use the following information about three substances as you solve heating curve problems.
Measurement
Specific Heat of Solid (J/g°C)
Melting Point (°C)
Heat of Fusion (J/g)
Specific Heat of Liquid (J/g°C)
Boiling Point (°C)
Heat of Vaporization (J/g)
Specific Heat of Gas (J/g°C)
Variable
csolid
Tfus
ΔHfus
Cliquid
Tvap
ΔHvap
Cgas
WATER
2.06
0
334
4.18
100
2260
1.7
NICKEL
0.444
650
450
0.750
1450
1067
1.050
EXAMPLES:
1) How much energy must be added to change the temperature of a sample
of 53 g of Water from –56 degrees Celsius to 160 degrees Celsius?
2) How much energy must be added to change the temperature of a sample
of 1560 g of Nickel from 2500o C to 50o C? (notice declining temperature)
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Unit 11 - Energetics
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3) How much energy is required to change the temperature of a 10.0 g sample of water from 15oC to 300oC?
4) How much energy is required to change the temperature of a 76 g sample of Nickel from 25oC to 700oC?
5) How much energy is required to change the temperature of a 50 g sample of Water from 99 degrees C to 101 degrees
C? Which of the three steps involved requires the most energy?
6) If 4155 calories of heat is added to 12 g of solid Nickel at 5o C, what will the final temperature be?
7) What would the final temperature be if 50 kJ of energy was added to a 200 g sample of water at 25o C?
8) If 1500 J of energy is added to 20 g of solid water at -20o C, what will the final temperature be? (this is more complex
than it sounds – plan it out, first)
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Unit 11 - Energetics
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Enthalpy, the measurement of energy of a thermodynamic system, is a state function. This means that the change in
enthalpy in going from some initial state to some final state is independent of the pathway. In going from a particular
set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in
one step or in a series of steps. This principle is known as Hess’s Law.
EXAMPLE:
The oxidation of nitrogen to produce nitrogen dioxide. The overall reaction can be written in one step:.
N2(g) + 2O2(g) → 2NO2(g)
ΔH1 = 68 kJ
The reaction also can be carried out in two distinct steps, with enthalpy changes designated by ΔH2 and ΔH3:
N2(g) + O2(g) → 2NO(g)
ΔH2 = 180 kJ
2NO(g) + O2(g) → 2NO2(g)
ΔH3 = −112 kJ
In order to use Hess’s Law to compute enthalpy changes for reactions, it will be helpful to apply the following rules:
1. Label each reaction, and treat it as a piece of a larger puzzle.
2. If a reaction is reversed, the sign of ΔH is reversed.
3. The magnitude of ΔH is directly proportional to the quantities of reactants and products in a reaction. If the
coefficients in a balanced reaction are multiplied by an integer or fraction, the value of ΔH is multiplied by the same
integer or fraction.
WE DO:
Target:
X + ½Z → W
ΔH = ?
The “puzzle pieces” are given as follows:
Piece “A”
Piece “B”
X + Y → W
2Y → Z
ΔH = −55 kJ
ΔH = −16 kJ
YOU DO:
Given the following data:
A + B → C
2B → D
Calculate ΔH for the reaction:
A + ½D → C
ΔH = − 10 kJ
ΔH = − 5 kJ
ΔH = ?
Sophomore Chemistry
Unit 11 - Energetics
1) Given the following data:
E + F → G
ΔH = − 20 kJ
E + G → H
ΔH = − 50 kJ
Calculate ΔH for the reaction:
2E + F → H
ΔH = ?
2) Given the following data:
K + 2L → M
4Q → R
Q + L → P + N
Calculate ΔH for the reaction:
K + 2N + 2P → M + ½ R
ΔH = − 40 kJ
ΔH = − 15 kJ
ΔH = − 25 kJ
ΔH = ?
3) Given the following data:
S(s) + 3/2 O2(g) → SO3(g)
2 SO2(g) + O2(g) → 2 SO3(g)
Calculate ΔH for the reaction:
S(s) + O2(g) → SO2(g)
ΔH = − 395.2 kJ
ΔH = − 198.2 kJ
ΔH = ?
4) Given the following data:
2 O3(g) → 3 O2(g)
O2(g) → 2 O(g)
NO(g) + O3(g) → NO2(g) + O2(g)
Calculate ΔH for the reaction:
NO(g) + O(g) → NO2(g)
ΔH = − 427 kJ
ΔH = + 495 kJ
ΔH = − 199 kJ
ΔH = ?
5) Given the following data:
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 =
ΔH =
ΔH =
ΔH =
+ 177.4 kJ
− 191.2 kJ
− 241.8 kJ
− 43.8 kJ
Calculate ΔH for the reaction:
C6H4(OH)2(aq) + H2O2(aq) → C6H4O2(aq) + 2 H2O(l) ΔH = ?
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Sophomore Chemistry
6) Given the following data:
H2(g) + ½ O2(g) → H2O(l)
N2O5(g) + H2O(l) → 2 HNO3(l)
½ N2(g) + 3/2 O2(g) + ½ H2(g) → HNO3(l)
Calculate ΔH for the reaction:
2 N2(g) + 5 O2(g) → 2 N2O5(g)ΔH = ?
Unit 11 - Energetics
ΔH = − 285.8 kJ
ΔH = − 76.6 kJ
ΔH = − 174.1 kJ
7) Given the following data:
2X + 4Z → 2Q + R
2W → Y + ½ R
2W → Z
Calculate ΔH for the reaction:
X + Y + Z → Q
ΔH = − 173 kJ
ΔH = − 122 kJ
ΔH = − 43 kJ
ΔH = ?
8) Given the following data:
P4(s) + 6 Cl2(g) → 4 PCl3(g)
P4(s) + 5 O2(g) → P4O10(s)
PCl3(g) + Cl2(g) → PCl5(g)
PCl3(g) + ½ O2(g) → Cl3PO(g)
Calculate ΔH for the following:
P4O10(s) + 6 PCl5(g) → 10 Cl3PO(g)
ΔH =
ΔH =
ΔH =
ΔH =
− 1225.6 kJ
− 2967.3 kJ
− 84.2 kJ
− 285.7 kJ
ΔH = ?
9) Given the following data:
2 B + 3/2 O2 → B2O3
B2H6 + 3 O2 → B2O3 + 3 H2O(g)
H2 + ½ O2 → H2O(l)
H2O(l) → H2O(g)
Calculate ΔH for the following:
2 B + 3 H2 → B2H6
ΔH =
ΔH =
ΔH =
ΔH =
− 1273 kJ
− 2035 kJ
− 286 kJ
+ 44 kJ
ΔH = ?
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Sophomore Chemistry
Unit 11 - Energetics
Unit 11 – Energetics – Take Away Sheet
Page 18
NAME: __________________________
Answer the following questions. Show your work. Box your answers.
1. A reaction has an enthalpy, ΔH, of −141.2 kJ/mole. Is it endothermic or exothermic? Explain.
2. Water has a melting point of 0°C and a boiling point of 100°C. Solid water has a specific heat of 2.06 J/g°C. Liquid
water has a specific heat of 4.18 J/g°C. Gaseous water has a specific heat of 1.7 J/g°C. The Heat of Vaporization of
water is 2260 J/g. How much energy is needed to raise the temperature of a 60.5 g sample of water from 55°C to
115°C?
3. A 35 g solid sample of Element X gives off 8500 J as it cools from 56°C to 34°C. Calculate the specific heat of Element
X, assuming that it doesn’t undergo any phase changes.
4. H2 + Br2 + 72.80 kJ  2HBr
a. How much heat will be absorbed when 50.0 g of bromine (Br2) reacts?
b. Sketch an energy diagram for this reaction. Assume the reactants have a potential energy = to 20 kJ.
Sophomore Chemistry
Unit 11 - Energetics
Page 19
5. Use Hess’s Law to determine the value of ΔH for reaction D.
Reaction A:
N2(g) + 3 H2(g) → 2 NH3(g)
ΔH = –91.8 kJ
Reaction B:
4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g)
ΔH = –906.2 kJ
Reaction C:
H2(g) + ½ O2(g) → H2O(g)
ΔH = –241.8 kJ
Reaction D:
½ N2(g) + ½ O2(g) → NO(g)
ΔH = ?
6. Use Hess’s Law to determine the value of ΔH for reaction 3.
Reaction 1:
Reaction 2:
Reaction 3:
C + O2 → CO2
CO + ½ O2 → CO2
C + ½ O2 → CO
ΔH = –393.5 kJ
ΔH = –283.0 kJ
ΔH = ?
Use the following information to solve problems 7 and 8.
(Hint: Some of the information might not be useful.)
Covalent
Bond
Bond
Energies (kJ/mol)
---------------------------------------------H–H
432
Cl–Cl
239
H–Cl
427
7. For the equation:
½ H2 + ½ Cl2 → HCl
determine Enthalpy using bond energies: