Chemistry 30
Unit 2:
Energy Changes in
Chemical Reactions
1
Energy:
In science, energy is often defined as “the ability to do _________________” or “the capacity to produce
_____________________________”.

_____________________________ energy is energy possessed by a body because of its ________________________

_____________________________ energy is the energy of _____________________________
o Temperature measures the _____________________________ kinetic energy of the substance, not the
total energy, since _____________________________ energy is not measured.
o Temperature is measured in Celsius or Kelvin. To change Celsius to Kelvin
________________________ and to change Kelvin to Celsius, _____________________________ 273.
Ex. Determine the temperature change in Celsius and in Kelvin if a container of water cools
from 98oC to 23oC.
∆t = tf – ti
Notice that the change in temperature is the _____________________________ if we use Celsius OR Kelvin.

_____________________________ energy is the ____________________ of the kinetic energy and potential energy
an object has.
o If you have a cup of water at 80oC and a bathtub of water at 30oC, which has more KE and
which has more PE? Which would have more thermal energy?

_____________________________ has more kinetic energy

_____________________________ has more potential energy

_____________________________ has more thermal energy
o _____________________________ energy is the _____________________________ of thermal energy from
one object to another. Heat always flows from the _____________________________ object to the
_____________________________ object. The transfer of energy can be detected by measuring the
resulting _____________________________ change.
o When energy is transferred to molecules, the molecules move _________________________, hit
each other more _________________________, and transfer more _________________________ to one
another.

_____________________________ energy is energy transmitted through space as _____________________________
waves. The most familiar form is _____________________________.

_____________________________ energy is the energy necessary to keep atoms joined by
___________________________. During a chemical reaction, chemical energy may be stored, released as
_________________________, converted to other forms of energy, or converted to ___________________________.
Whatever changes take place, energy is always _________________________. It is a fundamental law of nature
(the law of conservation of energy) that energy can neither be _________________________ nor destroyed but
may be converted from one _______________________ to another. Energy in the universe is _______________________.
The unit used to measure energy is the Joule: 1 J =
1𝑘𝑔∙𝑚2
𝑠2
1 J = 2.78 x 10-7 kW∙h
1 J = 2.39 x 10-4 kcal
2
Heat:



Heat is thermal energy in _____________________________.
If you touch something hot, energy is _____________________________ from the hot object to your hand. If
you touch something cold, heat energy passes from your hand to the cold object.
We should only use the term heat to describe energy that is _____________________________ from one
substance to another because of a temperature difference.
Specific Heat:


The _____________________________
of a substance is the amount of heat energy required to raise the
temperature of 1 gram of a substance by 1 degree Celsius or 1 kelvin.
Units used to measure specific heat = J•g−1•K−1 or J•g−1•°C−1
or
𝐽
𝑔∙𝐾

or
𝐽
𝑔∙℃
Different substances have different specific heats, this is because substances vary in their ability to
_________________________ energy, and they also gain or lose heat at different rates.
Ex. Comparing water to aluminum. It takes ______________ to raise the temperature of liquid water 1°C,
whereas for aluminum only _________________ are required. For the same increase in temperature, liquid
water requires over 4 times as much energy as the same mass of aluminum.
Why do you think water would be used as a coolant instead of another liquid (like mercury)?
Substance Specific heat (J/g·°C)
Al(s)
0.903
Fe(s)
0.449
Hg(l)
0.139
H2O(l)
4.184
O2(g)
0.917
He(g)
5.19
CO2(g)
0.843
Substance Specific heat (J/g·°C)
Au(s)
0.129
Na(s)
1.24
Cu(s)
0.385
H2(g)
14.4
H2O(g)
1.86
Ag(s)
0.235
H2O(s)
2.087
Heat Required to Change Temperature:

The formula to calculate heat required to change the temperature of a substance is:
Q = (m) (c) (∆t)
Where Q= __________________ amount (Joules)
m = _________________ (g)
∆t = tfinal – tinitial
*Losing heat would give you a negative Q*
c = _________________ heat capacity (J/g∙oC)
∆t = _____________________ change (oC or K)
3
Ex. Determine the heat required to raise the temperature of 100.0 mL of water from 298.0 K to 373.0 K.
Q = (m) (c) (∆t)
Heat capacity:

Heat capacity is the heat energy required to raise the temperature of a _____________________________ of
a substance by _____________ degree Celsius or one kelvin.
*Heat capacity is not the same as _____________________________ heat capacity*
Heat capacity (J/°C) = specific heat × mass
∴ Heat Capacity = Q/∆t
Ex. Determine the heat capacity of 1 cup of water (250mL).
*Remember the density of water is ~1g/mL*
Ex. Calculate the heat capacity of a piece of iron that releases 3500J of heat into a container of water, and
the temperature of the iron drops from 100oC to 24oC.
Molar Heat Capacity:

Molar heat capacity of a substance is the amount of energy required to raise the temperature of one
_____________________________of a substance by ____________ Celsius degree or one kelvin.
Molar heat capacity (J/mol•°C) = specific heat × molar mass
Ex. Calculate the molar heat capacity of water, given that the specific heat of water is 4.18 J/g•°C.
The molar mass of water is
Heat Capacity Assign
4
Molecular Motion:
There are three types of motion that molecules undergo:

_____________________________ Motion: Molecules are free to move along _____________________________
pathways from one place to another

_____________________________ Motion: When a molecule rotates about an _____________________________
through its center of mass.

_____________________________ Motion: molecules _____________________________ along the direction of a
bond.
Comparing types of motions among solids, liquids and gases:
Phase
Translational
Rotational
Vibrational
Gas
Liquid
Solid
Heat Capacities of gas molecules:

_____________________________ gases such as helium and neon have the _____________________________ molar
heat capacity
o

_____________________________ gases such as hydrogen and nitrogen have higher molar heat capacities
o

Have only _____________________________ motion since there is no bond about which to rotate or
on which to vibrate, therefore adding heat makes the atoms move faster, which increases the
temperature.
Have _____________________________ and _____________________________ motion as well as
_____________________________ motion, thus requiring more energy. Diatomic gases have similar
heat capacities.
_____________________________ gases such as water vapour and carbon dioxide have even higher heat
capacities
o These molecules can vibrate and rotate in more ways than diatomic gases, thus requiring
even _____________________________ energy.
5
Enthalpy:

Enthalpy is defined as the _____________________________ in a chemical reaction or physical process. It is
symbolized by the letter H.

The heat content if a substance cannot be measured _____________________________. Instead, we measure
enthalpy _____________________________ between the initial and final states. A chemical reaction
undergoes a change in enthalpy when the reaction either _____________________________ energy to the
surroundings or _____________________________ energy from the surroundings. This heat change is called
the enthalpy of reaction and is symbolized by _____________________________.

A thermochemical equation is written with the value of ΔH included:
 CuCl2(s)
Cu(s) + Cl2(g) 
ΔH = −220.1 kJ
or
Ex. The exothermic reaction of gaseous hydrogen and oxygen at constant pressure releases 241.8 kJ of
heat energy for every mole of water vapour formed. Write the thermochemical equation for the
production of one mole of water vapour.
Exothermic and Endothermic Reactions:

In a chemical reaction, the chemical
components involved in the process are
known as the _____________________________.
Everything outside the system then
becomes the _____________________________.

Since energy is always conserved, energy
lost by the system must be absorbed by
the _____________________________. Similarly,
energy absorbed by the system must
come from the surroundings.

A reaction in which the system releases heat to the surroundings is known as an
_____________________________reaction. If the system absorbs heat from the surroundings, it is an
_____________________________ reaction.
Exothermic Reactions
Endothermic Reactions
Enthalpy of products is less than enthalpy of
reactants
Hproducts < Hreactants
Energy is released
∆H is _____________________________
Less energy is needed to break bonds than the
energy that is released when bonds form
Enthalpy of reactants is less than enthalpy of
products
Hreactants < Hproducts
Energy is absorbed
∆H is _____________________________
More energy is needed to break bonds than the
energy that is released when bonds form
6
Research into the Thermochemistry of Hot and Cold Packs
*Adapted from Exp 2 of Chen 115.3 Laboratory Manual 2005-2006*
Purpose:
To develop a prototype hot or cold pack using a dissolved salt that produces an exothermic or
endothermic process.
Background Information:
In a chemical reaction where there is an energy change, heat can either be released or absorbed by the
reaction. If energy (in the form of heat) is released into the surroundings, the surroundings get warmer.
This is known as an exothermic reaction. If heat is absorbed from the surroundings, the surroundings get
cooler. This is known as an endothermic reaction. The amount of heat transferred between a system and
its surroundings (at constant pressure) is represented by the symbol ∆H and called enthalpy. Exothermic
processes have a negative ∆H and endothermic have a positive ∆H.
In this lab you will be working with a research team to develop either a hot pack or a cold pack. If you
choose to develop a hot pack, you must create one that will reach a maximum temperature of 45oC. If you
choose a cold pack the minimum temperature it must reach is 5oC.
The way you will develop your hot or cold pack is by dissolving either CaCl2 or NH4NO3 in 100mL of water.
Since dissolving CaCl2 in water is an exothermic process, it would increase the temperature of the water.
Dissolving NH4NO3 in water is an endothermic process, so it would cool the water by absorbing the heat
from the water. Before you can start developing your hot/cold pack, you must calculate the appropriate
amount of salt to dissolve in the water, so you know where to begin.
1.00 mole of CaCl2 releases 82.9KJ of heat energy when dissolved in water (∆Hosoln= -82.9KJ/mole) and
1.00 mole of NH4NO3 absorbs 25.7KJ of heat energy when dissolved in water (∆Hosoln= +25.7KJ/mole).
In this lab, in order to determine the amount of salt needed to raise/lower the temperature you will need
to try several different amounts of salts. Since you will be changing the amount of salts, this is the variable
in the experiment and should be the only thing you change (in order to accurately evaluate the data).
Everything else in the experiment should remain constant.
Procedure:
1. Complete the beginning of your formal lab report (up to the data/observations section- don’t forget
your hypothesis!).
2. Gather a team of 3 or 4 people to perform this experiment with. Determine whether you will create a
hot or cold pack.
3. Calculate the amount of heat energy that the water needs to release/absorb to lower/raise the
temperature of 100mL of water to the appropriate temperature. Remember the specific heat of water
is 4.184J/goC. This means that is takes 4.184J of energy to raise the temperature of 1g of water by 1oC.
The density of water at 20oC is 0.998g/mL. The initial temperature we will use for these calculations is
20oC.
4. Calculate the amount of salt you will need to use to raise/lower the temperature of the water to the
desired value. To do this, first convert the molar enthalpy of solution to enthalpy in terms of heat
7
energy released or absorbed by 1.00g of each salt (divide the molar enthalpy value by the molar mass
of the salt). This value is the amount of heat absorbed/released by 1.00g of salt dissolving in water. Use
this value to determine the amount of salt needed to release/absorb the amount of energy you
calculated in step three.
5. Depending on the size of your group, choose 6-8 different amounts of solid to use for this experiment
(the amount you calculate, 2 or 3 smaller amounts and 3 or 4 larger amounts using 5g increments). The
largest amount used should be less than 60g. Prepare your observations section of your lab report as
follows:
State whether your group will be creating a hot or cold pack.
Salt used:
Appearance:
Mass of _____________
(g)
Initial
Temperature of
H2O
Ti (oC)
Max/Min
Temperature
of Solution
Tm (oC)
Person who made the
measurements
Change in
temperature
∆T (oC)
Prototype Data:
Initial Temperature of Water:
Temperature Change Required:
Mass of ___________ Required:
Final Temperature of Water:
6. Each member of your group will choose 2 of the 6-8 selected amounts of solid and complete the
following steps. Complete the observations section of your lab report using the data your team collects.
a. In a beaker, weigh the solid to the nearest 0.1g. Remember to try to keep as many variables
constant as possible.
b. Using a graduated cylinder, measure out 100mL of distilled water. Record the initial
temperature of the water (remember to always estimate one digit).
c. Add the water to the solid and stir (DO NOT USE THE THEMOMETER TO STIR THE SOLUTION).
Observe the temperature increase/decrease and record the maximum/minimum temperature,
Tm (Note: this will note be the final temperature, as the temperature of the room will warm/cool
the solution after it reaches the max or min).
7. On graph paper, plot ∆T (y-axis) vs. mass of solid (x-axis) in each of the 6-8 trials. Each team member
should make their own graph. The plot should give a straight line.
8. Create a prototype pack: Measure the initial temperature of the water to be used for the prototype
pack. Determine the ∆T required to change the temperature to 45/5oC. From the graph you created,
determine the weight of solid you need to make the hot/cold pack.
9. Place the experimentally determined amount of solid (from the graph) into a Ziploc bag. Pour the water
into a sandwich bag (new bag), close the bag with a small rubber band and cut off the excess bag. Place
8
the sealed packet of water into the Ziploc bag, remove as much air as possible and seal the Ziploc bag.
Be sure the bag is sealed completely.
10. Activate the hot/cold pack.
WARNING: The bag may break open. If the concentrated solution gets on your skin, be sure to
wash it off immediately with lots of water.
11. Determine the final temperature of the hot/cold pack by wrapping it around the end of the
thermometer.
Analysis:
Write a paragraph that explains why dissolving a salt in water changes the temperature of the water.
1. Describe how well the prototype worked. Did it perform as planned? If not explain why. If you think
further research could be completed, explain what that might be.
2. Explain why there was a difference observed between the calculated theoretical amount of chemical
required, and the amount determined experimentally.
3. Explain why there was a difference observed (if you saw a difference) between the expected final
temperature and the actual final temperature of the hot/cold pack.
4. What would you expect to observe if you used a finely powdered solid rather than pelletized solid?
Would the mass of solid required change?
5. What do you expect would happen if you used the same amount of solid but used 50mL of water in
your hot/cold pack instead of 100mL? Explain.
Conclusion:
Summarize your results from this lab.
9
Types of Enthalpy Equations:





Enthalpy of _____________________________ Equations
o Each equation shows the _____________________________ of one mole of product from its elements
and also contains the loss or gain of heat energy when the compound is formed.
Enthalpy of _____________________________ Equations
o Each equation shows the _____________________________ of one mole of substance in O2 and
includes a heat term. The enthalpy of combustion equation is always has a loss of heat so ∆H
is negative
Enthalpy of _____________________________ Equations
o Each equation shows the _____________________________ on one mole of compound in water and
includes a heat term.
Enthalpy of _____________________________ Equations
o Each equation shows the energy involved in changing one mole of liquid to a gas at its
_____________________________ point.
Enthalpy of _____________________________ Equations.
o Each equation shows the energy involved in changing one mole of solid to a liquid at its
_____________________________ point.
Bond Energy & Enthalpy of Formation:

Atoms in molecules or formula units are held together by the _____________________________ being
attracted to the nuclei. These bonds can be easily broken by adding energy. This amount of energy
is known as _____________________________ energy.
Ex. H-H + 436.4kJ  2H
When a bond between two hydrogen atoms break, it requires 436.4kJ of energy to occur.
Alternatively, if 2 hydrogen atoms where to bond, 436.4kJ of energy would be released.
2H  H-H + 436.4kJ
Bond Energies at 298K
Bond


Bond energies can be used to determine
whether heat will be _________________________
or _____________________________and how much
when a mole of a compound forms from its
elements.
Bond energies are given in units kJ and
this represents the amount of energy
released/absorbed per _____________________
of substance.
H-N
H-O
H-S
H-P
H-H
H-F
H-Cl
H-Br
H-I
C-H
C-C
C=C
C≡C
C-N
Bond
Energy
(KJ∙mol-1)
393
460
368
326
436.4
568.2
431.9
366.1
298.3
414
347
620
812
276
Bond
C=N
C≡N
C-O
C=O
C-P
C-S
C=S
N-N
N=N
N≡N
N-O
N-P
O-O
Bond
Energy
(KJ∙mol-1)
615
891
327
804
263
255
477
393
418
941.4
176
209
142
Bond
O=O
O-P
O=S
P-P
P=P
S-S
S=S
F-F
Cl-Cl
Br-Br
I-I
C-Cl
N-F
Bond
Energy
(KJ∙mol-1)
498.7
502
469
197
489
268
352
150.4
242.7
192.5
151.0
326
275
10
Ex. Assume the following reaction takes place in a series of steps and energy is absorbed when bonds
break and released when bonds form. Determine whether the reaction is endothermic or exothermic, and
the amount of energy released or absorbed.
H2 + Cl2  2HCl
To determine final reaction we add these three steps together and cancel anything that occurs in reactants
and products:
Final Reaction:
∆H = ________________________ for the formation of one mole of hydrochloric acid.
∴ _______othermic and
We have now used bond energies to determine the amount of energy released when 1 mole of a
compound is formed from its elements. This is known as the enthalpy of formation.
See Bond Energies Assignment
Enthalpy of Formation and Hess’s Law:



The standard enthalpy of formation of a substance is the loss or gain in heat energy when one mole
of the substance is _____________________________ from its elements in their _____________________________
states, ΔH°f.
o The superscript ° indicates that the value of ΔH is measured when all substances are in their
_____________________________ states.
o The subscript f indicates that the energy is associated with the _____________________________ of
one mole of a compound from its elements in their standard states.
o The enthalpy of formation for different compounds is found using the bond energy
breaking/formation, just like we did in the last assignment.
The more negative ΔH°f, the more _____________________________ the compounds, because these
compounds require a lot of energy to break the bonds.
Many chemical reactions occur in _____________________________. For example, when carbon is burned in
the presence of excess oxygen, carbon dioxide is produced. In this reaction 393.5 kJ of heat energy
is liberated for every mole of carbon used. This means that the enthalpy change for the reaction is
−393.5 kJ/mol.
 CO2(g)
C(s) + O2(g) 
ΔH = −393.5 kJ
OR
(1)
(2)
 CO(g)
C(s) + ½ O2(g) 
 CO2(g)
CO(g) + ½ O2(g) 
 CO2(g)
C(s) + O2(g) 
ΔH = −110.5 kJ
ΔH = −283.0 kJ
ΔH = −393.5 kJ
The end result is the same whether it is formed in a one-step reaction or in a multi-step process. This is an
illustration of Hess’s Law of Constant Heat _____________________________: The change in enthalpy for a
chemical reaction is constant, whether the reaction occurs in one step or several.
11
*When elements are formed, there is no change in enthalpy so ∆H is 0 kJ/mol*
Substance
ΔHf (kJ/mol)
Substance
ΔHf (kJ/mol)
Al(s)
Al2O3(s)
Br2(l)
HBr(g)
Ca(s)
CaCO3(s)
CaCl2(s)
C(s) (graphite)
C(s) (diamond)
CCl4(l)
CCl4(g)
CHCl3(l)
CH4(g)
C2H2(g)
C2H4(g)
C2H6(g)
C3H8(g)
C6H6(l)
CH3OH(l)
C2H5OH(l)
CH3CO2H(l)
CO(g)
CO2(g)
COCl2(g)
CS2(g)
CS2(l)
Cl2(g)
HCl(g)
Cr(s)
CrCl3(s)
Cu(s)
CuO(s)
CuCl(s)
CuCl2(s)
F2(g)
HF(g)
He(g)
H2(g)
H2O(l)
H2O(g)
H2O2(l)
Fe(s)
FeO(s)
Fe2O3(s)
Fe3O4(s)
FeCl2(s)
FeCl3(s)
FeS2(s) (pyrite)
Pb(s)
0
−1675.7
0
−36.4
0
−1206.9
−795.8
0
+1.90
−135.4
−96.0
−134.5
−74.8
+226.7
+52.3
−84.7
−103.8
+49.0
−238.7
−277.7
−484.5
−110.5
−393.5
−218.8
+117.4
+89.70 kJ
0
−92.3
0
−556.5
0
−157.3
−137.2
−220.1
0
−271.1
0
0
−285.8
−241.8
−187.8
0
−272.0
−824.2
−1118.4
−341.8
−399.5
−178.2
0
PbCl2(s)
Mg(s)
MgCl2(s)
MgO(s)
Hg(l)
HgS(s)
Ne(g)
N2(g)
NH3(g)
N2H4(l)
NH4Cl(s)
NH4NO3(s)
NO(g)
NO2(g)
N2O(g)
N2O4(g)
HNO3(l)
O(g)
O2(g)
O3(g)
P4(s) (white)
P4(s) (red)
PH3(g)
PCl3(g)
P4O6(s)
P4O10(s)
H3PO4(s)
K(s)
KCl(s)
KClO3(s)
KOH(s)
Ag(s)
AgCl(s)
AgNO3(s)
Na(s)
NaCl(s)
NaOH(s)
Na2CO3(s)
S(s) (rhombic)
S(s)
SF6(g)
H2S(g)
SO2(g)
SO3(g)
H2SO4(l)
Sn(s) (white)
Sn(s) (grey)
SnCl2(s)
SnCl4(l)
−359.4
0
−641.3
−601.7
0
−58.2
0
0
−46.1
+50.6
−314.4
−365.6
+90.3
+33.2
+82.1
+9.2
−174.1
+249.2
0
+142.7
0
−70.4
+5.4
−287.0
−2144.3
−2984.0
−1279.0
0
−436.7
−397.7
−424.8
0
−127.1
−124.4
0
−411.2
−425.6
−1130.7
0
+278.8
−1209.0
−20.6
−296.8
−395.7
−814.0
0
−2.1
−325.1
−511.3
12
Finding ΔH for an Equation (Hess’s Law cont.):

If two or more thermochemical equations are added to give a final equation, then the
_____________________________ can be added to give the enthalpy for the final equation.
Ex. Use the given intermediate steps (1 and 2) for the production of one mole of tetraphosphorus
decaoxide, to determine the ΔH for the overall reaction from its elements.
Intermediate steps:
(1)
 P4O6(s)
4 P(s) + 3 O2(g) 
ΔH = −1640 kJ
(2)
P4O6(s) + 2 O2(g) 
 P4O10(s)
ΔH = −1344 kJ
Cancelling out species that appear on both sides of the reaction, we are left with the equation for the
overall reaction:
The enthalpy of the overall reaction is the sum of the enthalpies of the intermediate steps.
ΔH = ________________________ + __________________________ = ___________________________
Since ΔH° < 0, the reaction is __________othermic.

Sometimes you may be required to _____________________________ an intermediate step in order for
the reactants and products you want to be on the appropriate side. If you reverse an
intermediate step, don’t forget to reverse the _____________________________ of the ΔH.

You may also need to multiply or divide the intermediate equations so that the quantities of
reactants/products match the quantities you want in your final reactions. If you multiply of
divide an intermediate step, remember to multiply or divide your ___________________ value as well.
Ex. The enthalpy changes for the following reactions are
(1)
 2 CO2(g) + H2O(l)
C2H2(g) + 5/2 O2(g) 
ΔH = −1299 kJ
(2)
 6 CO2(g) + 3 H2O(l)
C6H6(l) + 15/2 O2(g) 
ΔH = −3267 kJ
Find ΔH° for the following reaction:
3 C2H2(g) 
 C6H6(l)
Is the reaction endothermic or exothermic?
See Hess’s Law Assign
13
Enthalpy Change in a Reaction

From Hess’s law, we now know the enthalpy change for a reaction can be determined by
subtracting the sum of the heat of formation of the reactants from the sum of heat of formations of
the products.
ΔH°reaction =
Ex. Using the ΔH°f values from your tables, determine the enthalpy change for the combustion reaction of
benzene, C6H6(l).
See Enthalpy Change in Rxn Assign
Enthalpy and Phase Change:

In all solids the particles vibrate back and forth weakly. They have low kinetic energy. Heating a
solid results in ______________________ vibration of particles so their kinetic energy increases. The solid
particles move further away from one another, which causes their potential energy to _______________.

These energy changes can be observed by an increase in _____________________________ and an increase
in _____________________________.
o Note that the physical appearance of the solid will not change.

The heat energy (Q) involved in a solid changing temperature can be calculated using _______________.

A _____________________________ change occurs when the physical state of a substance changes. Melting,
freezing, evaporation and condensation involve phase changes.

Both melting and vaporization are _____________________________ processes: heat is absorbed from the
surroundings.

During the reverse processes, in
which steam condenses or water
freezes, heat energy is released to
the surroundings. These are
_____________________________ processes.
These enthalpy values are negative.

During a phase change, the
_____________________________ of the
substance does not change even
though you are adding/removing
energy. This is because the energy
being added or removed is working
to change the state of the substance,
not change the temperature.
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
The heat energy (Q) involved in changing the state of one mole of a substance at its boiling/melting
point is known as the _____________________________ enthalpy of phase change.
Molar enthalpies of phase change (kJ•mol−1)
Substance
ΔHvap ΔHcond ΔHmelt ΔHfre
Water
+40.7
-40.7
+6.02
-6.02
Methane
+10.4
-10.4
+0.94
-0.94
Mercury
+59.3
-59.3
+2.3
-2.3
Sodium chloride
+207
-207
+27.2
-27.2
*Note: values for melting/freezing or vaporization/condensation are the same, but have opposite signs *

When you have molar enthalpies of phase change you can use these values to calculate the heat
energy. The formula is as follows:
Q = ∆Hvap, cond, melt, fre ∙ n
Where Q = heat energy
∆Hvap, cond, melt, fre = molar enthalpy of phase change
n = # moles
Ex. Calculate the enthalpy change involved when 17.2 g of liquid mercury is vaporized.
First calculate moles of Hg:
See Energy Needed to Melt Ice Lab &
Enthalpy and Phase Changed Assignment
Calorimetry

The branch of thermochemistry
concerned with measurement of
such _____________________________ is
known as calorimetry. Any device
used for measuring a heat change
is called a calorimeter.

When we measure heat changes
with this calorimeter, we assume
that all the heat is absorbed by the
__________________________, and that the
polystyrene cup does not absorb
heat and that no heat is lost to the surroundings. If the quantity of heat emitted during the reaction is
not large, such a “coffee cup” calorimeter will give reasonably accurate results. If large quantities of
heat are involved, then heat losses to the surroundings will lead to ___________________________.
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Ex. A mass of 100 g of water is placed in a coffee cup calorimeter. The solution temperature is measured to
be 14.4°C. A mass of 0.412 g of calcium metal is placed in the calorimeter. When the reaction is complete
the temperature is recorded as 24.6°C. Calculate the standard molar enthalpy change for this reaction:
Start by finding the heat amount that the water changed by (Q), then determine the amount of
heat release by the reaction. Next we will determine molar enthalpy by dividing the heat that
the reaction released by the the mole s of calcium used in the reaction.
 Ca(OH)2(s) + H2(g)
Ca(s) + 2 H2O(l) 
The Flame Calorimeter:
Some calorimeters may not be fully insulated. This means that the water in the calorimeter and the
calorimeter itself _____________________________ absorb heat.
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Ex. A flame calorimeter composed of steel and with a mass of 322 g is filled with 225 g of water. The
original temperature of the water and calorimeter is 10.6°C. When 1.02 g of ethanol, C2H5OH, is burned,
the temperature rises to 38.4°C. Calculate the molar enthalpy of combustion of ethanol.
Specific heat of steel = 0.44 J/g•C
The heat released by the combustion process will be absorbed (mainly) by the water and the calorimeter
as the temperature rises by 27.8°C (the change in temperature).
Heat absorbed by water and calorimeter: Q = mc∆t for water + mc∆t for steel
See Calorimetry Assignment
17
18
Name:_____________________
Heat Capacity Assignment
1. Calculate the heat capacity of an aluminum block that must absorb 629 J of heat from its surroundings
in order for its temperature to rise from 22°C to 145°C.
2. Calculate the heat capacity of a sample of brake fluid if the sample must absorb 911 J of heat in order for
its temperature to rise from 15°C to 100°C.
3. A burner on an electric range has a heat capacity of 345 J/K. What is the value of q, in kilojoules, as the
burner cools from 467°C to 23°C?
4. How much heat, in joules and in kilojoules, does it take to raise the temperature of 225 g of water from
25.0 to 100.0°C?
5. How much heat in kilojoules, does it take to raise the temperature of 814 g of water from 18.0°C to
100.0°C?
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6. What mass of water, in kilograms, can be heated from 5.5°C to 55.0°C by 9.09 × 1010 J of heat?
7. What will be the final temperature if a 5.00-g silver ring at 37.0°C gives off 25.0 J of heat to its
surroundings?
8. A 454-g block of lead is at an initial temperature of 22.5°C. What will be the temperature of the lead
after it absorbs 4.22 kJ of heat from its surroundings?
9. How many grams of copper can be heated from 22.5 to 35.0°C by the same quantity of heat that can
raise the temperature of 145 g H2O from 22.5 to 35.0°C?
10. Calculate the molar heat capacity of ethanol, C2H5OH(l). The specific heat of ethanol is 2.46 J/g·°C.
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11. When 5.16 kJ of heat is added to 167 g of gaseous ammonia at 45.0°C the temperature of the gas rises to
60.0°C. From this data determine the following:
a) the specific heat of NH3(g)
b) the molar heat capacity of NH3(g)
12. The molar heat capacity of liquid sodium is 28.4 J/K•mol. How much heat is required to raise the
temperature of 5.67 g of liquid sodium by 3.75 kelvins?
(Find sp. heat, then Q)
13. If, at 25°C, 15.7 g of carbon dioxide absorbs 1.2 kJ of heat, calculate the final temperature of the gas. The
molar heat capacity of CO2 is 37.11 J/K•mol.
14. A 27.7 g sample of ethylene glycol, a car radiator coolant, loses 688 J of heat. What was the initial
temperature of the ethylene glycol if the final temperature is 32.5°C (c of ethylene glycol = 2.42 J/g•K)?
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Name:______________________
Bond Energy & Enthalpy of Formation
Determine the bond energy released/absorbed in the following reactions. Show your work by writing out
the bond breaking and formations. Be sure to balance the equations:
1. NH3 + Cl2  N2 + HCl
2. CH4 + O2  CO2 + H2O
3. CH3OH + O2  CO2 + H2O
4. Cl2 + CH4  CH3Cl + HCl
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Name:________________________
Hess’s Law Assignment
1. The enthalpy changes for the formation of two tungsten bromides are shown below.
W(s) + 2 Br2(l) → WBr4(s)
ΔH° = −146.7 kJ
W(s) + 3 Br2(l) → WBr6(s)
ΔH° = −184.4 kJ
Calculate the standard enthalpy change for the following reaction.
Br2(l) + WBr4(s) → WBr6(s)
2. Given:
N2O4(g) → 2 NO2(g)
NO(g) + ½ O2(g) → NO2(g)
ΔH° = +58 kJ
ΔH° = −56 kJ
Calculate the standard enthalpy change for the following reaction.
2 NO(g) + O2(g) → N2O4(g)
3. Use the following reactions and enthalpy changes to predict the standard enthalpy change for:
2 NO2(g) + 2 H2O(g) → 3 O2(g) + N2H4(g)
½ N2(g) + O2(g) → NO2(g)
H2(g) + ½ O2(g) → H2O(g)
N2(g) + 2 H2(g) → N2H4(g)
ΔH° = +33.2 kJ
ΔH° = −241.8 kJ
ΔH° = +47.6 kJ
4. Use the following formation reaction evidence to calculate the standard enthalpy change for the
complete combustion of cycloheptane.
C(s) + O2(g) → CO2(g)
ΔH° = −393.5 kJ
H2(g) + ½ O2(g) → H2O(g)
ΔH° = −241.8 kJ
7 C(s) + 7 H2(g) → C7H14(l)
ΔH° = +115.0 kJ
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5. Given Equation (a), calculate ΔH for Equation (b).
a)
H2(g) + I2(s)  2 HI(g)
ΔH = +52.96 kJ
b)
HI(g)  ½ H2(g) + ½ I2(s)
6. From the following two thermochemical equations,
 Fe2O3(s)
2 Fe(s) + 3/2 O2(g) 
ΔH = −824 kJ


2 Al(s) + 3/2 O2(g)
Al2O3(s)
ΔH = −1676 kJ
calculate the enthalpy change for the reaction
 2 Fe(s) + Al2O3(s)
2 Al(s) + Fe2O3(s) 
7. Calculate the enthalpy change for reaction:
2 C(graphite) + 2 H2(g) → C2H4(g)
Given the data in equations:
 CO2(g)
C(graphite) + O2(g) 
 2 CO2(g) + 2 H2O(l)
C2H4(g) + 3 O2(g) 
 H2O(l)
H2(g) + ½ O2(g) 
8. Calculate the enthalpy change for the reaction
2 CH4(g) + 3 O2(g) → 2 CO(g) + 4 H2O(l)
from the following data:
CO(g) + ½ O2(g) → CO2(g)
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l)
9. Calculate the enthalpy change for the reaction
C2H4(g) + H2(g) → C2H6(g)
given the following data:
C2H4(g) + 3 O2(g) → 2 CO2(g) + 2 H2O(l)
2 C2H6(g) + 7 O2(g) → 4 CO2(g) + 6 H2O(l)
2 H2(g) + O2(g) → 2 H2O(l)
ΔH = ?
ΔH = -393.5 kJ
ΔH = -1410.9 kJ
ΔH = -285.8 kJ
ΔH = ?
ΔH = -283.0 kJ
ΔH = -890.3 kJ
ΔH = ?
ΔH = -1410.9 kJ
ΔH = -3119.4 kJ
ΔH = -571.6 kJ
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Name:___________________
Enthalpy Change in a Reaction Assignment
1. Synthesis gas is a mixture of carbon monoxide and hydrogen that is used to synthesize a variety of
organic compounds. One reaction for producing synthesis gas is
3 CH4(g) + 2 H2O(l) + CO2(g) → 4 CO(g) + 8 H2(g)
Use standard enthalpies of formation from your data tables to calculate the standard enthalpy change
for this reaction.
2. Ethylene, derived from petroleum, is used to make ethanol for use as a fuel or solvent. The reaction is
C2H4(g) + H2O(l) → CH3CH2OH(l)
Use your data tables to calculate ΔH o for this reaction.
3. Using the ΔH°f values from your tables, determine the enthalpy change for the reaction:
2C(graphite) + 2 H2(g) → C2H4(g)
Compare this result with question 7 from the Hess`s Law Assignment.
4. Calculate the enthalpy change for the reaction below using your data tables. Compare with question 8
from the Hess`s Law Assignment.
2 CH4(g) + 3 O2(g) → 2 CO(g) + 4 H2O(l)
5. Calculate the enthalpy change for the reaction below using your data tables.
10 N 2O g   C3 H 8 g  
 10 N 2 g   3 CO2 g   4 H 2O g 
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Name:_________________
Energy Needed to Melt Ice: Calculating Molar Heat of Melting
*There is no formal lab report required for this lab*
Molar Heat of Melting is the energy needed to change one mole of a solid at its melting point into a liquid.
Procedure
Observations
o
1. Measure 100mL of warm water (50 C) into a calorimeter. Initial Volume of Warm Water:
Accurately record the temperature of the water.
________________mL
2. Add a couple of dry ice cubes to the water. Stir the
mixture until the temperature is about 0oC. Add more ice Initial Temperature of Warm Water:
_____________oC
if needed to cool the water. Record the lowest
temperature reached
Coldest Temperature of Water:
3. Remove any unmelted ice draining back as much water as _____________oC
possible into the cup
Volume of water remaining in the
4. Measure the volume of water remaining in the
cup: ____________mL
calorimeter.
Calculations:
1. Calculate the temperature change of the water:
2. Calculate the mass of water that supplied the ice with heat energy:
3. Calculate the heat energy lost by the water (Q=):
4. How much energy was gained by the ice (Q=)?
5. Calculate the mass of ice melted:
6. Calculate the number of moles of ice melted:
7. Calulate the molar heat of melting of ice (kJ/mole)
8. Calculate your % error:
9. Write an equation for the melting of ice and include the molar heat of melting in the equation:
26
Name:_________________
Enthalpy and Phase Changes Assignment
1. Calculate the energy needed to change 36g of ice at 0oC to liquid at 0oC.
2. Calculate the amount of energy you need to remove to change 3.0 moles of steam at 110oC to liquid
water at 100oC.
3. Calculate the energy needed to change 9.0g of ice at 0oC to steam at 135oC.
4. What is the heat change in J associated with 20.3 g of liquid water at 5.00 ° C changing to solid water at
-5.00 °C?
27
5. Calculate q when 0.10 g of ice is cooled from 10.0°C to −75°C.
6. How many grams of ice at -14.4oC can be completely converted to liquid at 7.4oC if the available heat
for this process is 5.58×103kJ ?
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Name:________________________
Calorimetry Assignment
1. A 24.6 g sample of nickel is heated to 110.0°C and then placed in a coffee cup calorimeter containing
125 g of water at a temperature of 23.00°C. After the nickel cools, the final temperature of the metal
and water is 24.83°C. Calculate the specific heat of nickel.
2. A flame calorimeter made of copper has a mass of 587 g and is filled with 314 g of water. The original
temperature of the water and calorimeter is 12.2°C. When 0.920 g of 1-propanol, C3H7OH, is burned,
the temperature rises to 32.3°C. Calculate the molar enthalpy of combustion of 1-propanol, given the
following additional information: specific heat of water = 4.18 J/g•°C, specific heat of copper = 0.385
J/g•°C
3. In a purity check for industrial diamonds, a 10.25-carat (1 carat = 0.2000 g) diamond is heated to
74.21°C and immersed in 26.05 g of water in a constant-pressure calorimeter. The initial temperature
of the water is 27.20°C. Calculate ΔT of the water and of the diamond (cdiamond = 0.519 J/g•K).
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4. A 15.5-g sample of a metal alloy is heated to 98.9°C and then dropped into 25.0 g of water in a
calorimeter. The temperature of the water rises from 22.5 to 25.7°C. Calculate the specific heat of the
alloy.
5. Find the heat transferred (in kJ) when 5.50 L of ethylene glycol (d = 1.11 g/mL and c = 2.42 J/g•K) in a
car radiator cools from 37.0°C to 25.0°C.
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