SCH4U1
UNIT 3  Energy Changes and Rates of Reaction
STSE Focus:
D1.1 analyse some conventional and alternative energy technologies (e.g., fossil
fuel–burning power plants, hydro-powered generators, solar panels, wind turbines,
fuel cells), and evaluate them in terms of their efficiency and impact on the
environment
Sample issue: The cooling of homes and commercial buildings in summer requires more
energy than heating in the winter at peak times. Brownouts are more likely in summer
than in winter. However, new technologies use deep lake water cooling as an alternative
to conventional air conditioning systems in office towers. This significantly reduces
energy use and its environmental impact.
Sample questions: What proportion of Ontario’s energy needs is served by solar and
wind technologies? What are the pros and cons of expanding the availability of these
technologies? What types of chemical reactions occur in different types of fuel cells?
What are the advantages and disadvantages, in terms of efficiency and environmental
impact, of using corn to produce ethanol fuel?
D1.2 analyse the conditions (e.g., temperature, pressure, presence of a catalyst)
required to maximize the efficiency of some common natural or industrial chemical
reactions (e.g., decomposition, combustion, neutralization), and explain how the
improved efficiency of the reaction contributes to environmental sustainability.
Sample issue: Bleaches such as hydrogen peroxide and chlorine are used when fibres are
processed into paper or textiles. Concentrations of these substances can harm the
environment, but if enzymes are added to these processes as biocatalysts, fewer
chemicals are needed, less energy is consumed, and there is less environmental impact.
Sample questions: How can you increase the rate of decomposition in a home
composter? What can be done to improve the efficiency of an automobile that runs
entirely on fossil fuels? Why is just a very small quantity of catalyst required in industrial
processes? Why is the ozone layer still deteriorating despite the banning of
chlorofluorocarbons (CFCs)?
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SCH4U1
UNIT 3 – Thermochemistry (Chapter 5)
*Opening Reflection Questions:
1. Consider the following changes: ice melting, water evaporating, water vapour condensing,
photosynthesis, respiration, and combustion of gasoline. Classify these changes as absorbing or releasing
thermal energy.
2. Based on your current understanding of energy, how is electrical power produced in Ontario? What are
the sources of energy that produce this power?
3. How is nuclear power different from hydroelectric power? How is it similar?
*Activity: Burning Food
*Mass of nut = _________
a) How much energy was absorbed by the water?
heat = (mass of water) x (specific heat capacity of water) x (temperature change of water)
q = mc∆T
= (50 g) x (4.18 J/g°C) x (_______)
=
b) Where did this energy come from?________________________________________________
c) Calculate the amount of heat produced per gram.
q=
qTotal
=
m Nut
d) What is the relationship between a calorie, Calorie, and Joule?
*1 calorie (gram calorie) = __________ Joules (J)
*1 J = the energy required to increase the temperature of ____ g of water by____°C (This value, 4.18, is
referred to as the _______________heat capacity of water.)
*1000 calories = 1 kilocalorie = 1 Calorie (kg calorie) = 1 “food Calorie” = 4.18 kJ
*4.18 kJ = energy required to increase the temperature of 1 ______ of water _____°C
*There are ___________ calories in 1 pound of fat.
Some Energy Densities:
*Fat
 9 Cal / gram
*Protein  4 Cal / gram
*Carbs  4 Cal / gram
(9 x 4.18 = 37 kJ / gram)
(4 x 4.18 = 17 kJ / gram)
(4 x 4.18 = 17 kJ / gram)
5.1 Changes in Matter and Energy
Terms:
 Thermodynamics - the study of energy and energy transfer
 Thermochemistry - the study of energy involved in ______________ energy
 Law of Conservation of Energy - total energy of the universe is ___________; energy can neither be
_____________ nor ______________
 Chemical System - the _____________ and ____________ of a chemical reaction; represented by a
chemical equation
 Surroundings - the system’s _________________
 Universe = System + Surroundings
Q. What is the difference between heat, thermal energy, and temperature?
 Heat (q) - refers to the amount of ___________ energy ______________ between substances; heat is
transferred, not possessed
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Thermal energy - an object possesses thermal energy [type of ____________ energy (energy associated
with the moving objects/entities)] *heating = giving substance thermal energy and thereby increasing the
kinetic energy of the molecules)
 Temperature - a measure of the ___________________ kinetic energy of the particles that make up a
substance or system [measured in °C or Kelvin (K)]
*Explain the following: A glass of water at 100°C has a higher temperature than a swimming pool full of
water at 20°C, but the pool has more thermal energy than the glass of water.

*Measuring Energy Changes: CALORIMETRY
Calorimetry – the technological process of measuring energy changes in a chemical system
*The Heat Transfer Formula:
q = _______________
i.e., the quantity of heat transferred depends on:
1. the _______________ of substance present
2. the _________________change
3. the ___________ of substance involved [i.e. the specific heat capacity *(c) of the substance]
*c = the quantity of heat required to raise the temperature of a unit mass (1 gram or 1 kg) of a substance
1°C (e.g. The specific heat capacity, c, of water is 4.18 J/(g • °C). Explain what this means in one
sentence.) 
Q. When 600 g of water in an electric kettle is heated from 20°C to 95°C to make a cup of tea, how much
heat flows into the water?
*Try p. 302 #8 – 13
Heat Transfer and Enthalpy Change
 Enthalpy (H) – the total internal energy of a substance at constant pressure (the sum of many different
forms of energy, both kinetic and potential, present in a chemical system)
*These include the energies of:
i) moving _______________ within atoms
ii) the vibration of _______________ connected by chemical bonds; and
iii) the rotation and translation of molecules that are made up of these atoms.
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iv) the _________________ potential energy of protons and neutrons in atomic nuclei; and
v) the electronic potential energy of atoms connected by chemical bonds.
**Impossible to measure enthalpy of a chemical system
** It is possible, however, to determine enthalpy _____________.
 Enthalpy change (∆H) – the energy _____________ or ___________ to the surroundings when a
system changes from ___________ to ____________ (Can be determined using calorimetry, for example.)
*Enthalpy changes result from chemical bonds (sources of stored chemical potential energy) being
__________and __________. The formation of a bond ___________ energy; breaking a bond
_____________energy.
 Endothermic change – net absorption of energy by the chemical system (i.e. the energy absorbed in
the breaking of bonds is ____________ than the energy released in the formation of product bonds). For
an endothermic change the temperature of surroundings ______________.
 Exothermic change – net release of energy by the chemical system (i.e. the energy absorbed in the
breaking of bonds is __________ than the energy released in the formation of product bonds). For an
exothermic change the temperature of surroundings _____________.
Blue line (Ek) = change to surroundings
Red line (∆H) = change to system
*When a change occurs in a system, the chemical
potential energy change, ∆H, is numerically _______
to the heat (q) transferred to the ______________.
(e.g. If 50 kJ of energy is ______ by the system, 50 kJ
of energy is _________ by the surroundings and vice
versa.) This is represented mathematically by the
following formula:
∆Hsystem =  q surroundings
Sign Conventions:
 Enthalpy changes for exothermic reactions are given a ____________ sign. (The negative sign refers to
the fact that energy has __________ the system; the energy of the system is decreasing..)
 Enthalpy changes for endothermic reactions are given a ____________ sign. (i.e. Energy has been
______________ by the system; the energy of the system is increasing.)
Exothermic:
Endothermic:
Types of Enthalpy Changes That Matter Can Undergo:
1. _______________ change  no chemical bonds are broken (no new substance produced)
e.g. change of state: H2O(l) + heat  H2O(g)
dissolving:
CaCl2(s)
2O
H

 CaCl2(aq) + heat
2. _______________ change  chemical bonds between atoms are rearranged resulting in new substances
e.g. combustion of propane: C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g) + heat
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3. _____________ change – a change in the protons or neutrons in an atom, resulting in the formation of
new atoms
e.g. nuclear fusion/fission; nuclear decay:
238
92
U  24 He +
234
90
Th + heat
*Enthalpy changes associated with each of these changes:
Change
Enthalpy Change Range (kJ/mol)
Physical
100  _________
Chemical
102  _________
Nuclear
*1010  __________
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*E = mc ; where m = mass & c = 3 x 108 (i.e. the speed of light)
 Therefore, a small amount of _______ contains a massive amount of __________. Mass is really just
“frozen” energy. This equation describes the amounts of energy evolved via nuclear _________
reactions in the ________.
5.2 Molar Enthalpies
*Molar Enthalpy, (∆Hx)  the enthalpy change associated with a physical, chemical, or nuclear change
involving _______ mole of a substance (The variable,x , indicates the type of change occurring .)
Example:
* Describe the type of molar enthalpy that would be associated with each of the following
changes/reactions.
a) Br2(l)  Br2(g) ------------------------------------------
b) CO2(g)  CO2(s) -----------------------------------------
c) LiBr(s)  Li+(aq) + Br-(aq)-------------------------------
d) C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(l) --------------
e) NaOH(aq) + HCl(aq)  2NaCl(aq) + H2O(l)---------
Ex. 1: Combustion of hydrogen
H2(g) + 1/2O2(g)  H2O(g) + 285.8 kJ
**∆Hcombustion = -285.8 kJ/mol
i) How much energy would be released to the surroundings if:
- 2 moles of hydrogen were combusted? ________________
- 6.2 moles? ___________________
*Equation used:
(*For this example, the subscript x = __________________)
Ex. 2: Vapourization of water:
H2O(l) + 40.8 kJ  H2O(g)
**∆Hvapourization = +40.8 kJ/mol
i) How much energy does the water absorb from the surroundings if:
- 3 moles of water are vapourized? ____________
*Equation used:
*Try the following:
1. When 4.0 g of methane burns in oxygen, 200 kJ of heat is evolved. Calculate the amount of heat (in kJ)
evolved when 1 mol of methane burns.
2. When 0.726 g of carbon reacts with sulfur to give carbon disulfide, 5.40 kJ is absorbed. Calculate the
heat absorbed (in kJ) when 1 mole of carbon disulfide is formed from carbon and sulfur.
3. Hydrazine, N2H4(l), is used in rocket fuel. The thermochemical equation for the combustion of
hydrazine is:
N2H4(l) + O2(g)  N2(g) + 2H2O(l) + 622.4 kJ
What quantity of heat is liberated by the combustion of 1.00 g of N 2H4(l)?
4. Glucose, C6H12O6(s), is converted into ethanol, C2H5OH(l), in the fermentation of fruit juice to
produce wine:
C6H12O6(s)  2C2H5OH(l) + 2CO2(g) + 67.0 kJ
What quantity of heat is liberated when a 0.500 L of wine containing 47.5 g of C 2H5OH(l) is produced?
5. Determine the quantity of heat required to completely vaporize 5.4 g of water.
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*More Practice  p. 308 #1-3
*Using Calorimetry to Find Molar Enthalpies
*There are three simplifying assumptions often used in calorimetry:
1. no heat is transferred between the _____________ and the outside __________________;
2. any heat absorbed or released by the calorimeter materials, such as the container, is negligible; and,
3. a dilute aqueous solution is assumed to have a density and specific heat capacity equal to that of pure
_______________(1.00 g/mL and 4.18 J/(g•°C) or 4.18 kJ/(kg•°C).
Example:
*In a chemistry experiment to investigate the properties of a fertilizer, 10.0 g of urea, NH 2CONH2(s), is
dissolved in 150 mL of water in a simple calorimeter. A temperature change from 20.4°C to 16.7°C is
measured. Calculate the molar enthalpy of solution for the fertilizer urea.
*Formula:
ΔH = q
*Therefore:
=
n∆Hsolution
=
SYSTEM (urea)
mc∆T
SURROUNDINGS (H2O)
m=
n=
=
=
∆Hsolution = ?
mcT
n
(
)(
=
(
cH2O = _________ J/(g°C)
∆T =
=
∆Hsolution =
)(
)
)
= _____________ J/mol
= _____________ kJ/mol
*Is this value (i.e. ∆Hsolution) negative or positive?
If:
-
the temperature of the surroundings INCREASES, ∆Hx is _______________ - energy has
__________the system and entered the surroundings;
-
the temperature of the surroundings DECREASES, ∆Hx is _______________ - energy has been
_______________ by the system from the surroundings.
*In this example the temperature of the surroundings _______________, therefore the change is
__________________ and ∆Hsolution = __________ kJ/mol.
*Translation  For every mole of urea that is dissolved in a sufficient amount of water, 13.9 kJ of energy
is _____________ by the _______________.
*Try pp. 310- 311 #5, 10; p. 312 # 1-5
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5.3 Representing Enthalpy Changes
The Four Methods:
1. Thermochemical Equations with Energy Terms
2. Chemical Equations with ∆H values separate from the equation
3. Molar Enthalpies of Reaction
4. Potential Energy Diagrams:
Ex. 1  An exothermic process (Respiration i.e. combustion of glucose)
*Method 1:
*Method 2:
*Method 3:
*Method 4:
Ex. 2  An endothermic process (Photosynthesis i.e. reverse of respiration reaction)
*Method 1:
*Method 2:
*Method 3:
*Method 4:
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Try pp. 319-320 #1-5; p. 320 (bottom) #1-3 (4 – challenge)
5.4 Hess’s Law
*Hess’s law provides another method by which to calculate enthalpy changes. (Method 1 =____________)
*When is Hess’s Law used?
 when chemical systems cannot by analyzed using________________
Ex. 1 Rusting of iron
 extremely __________ reaction
 resulting temperature change would be too ________to be measured using convention calorimetry
Ex.2 Combustion of carbon
 impossible to measure with a calorimeter because the combustion of carbon produces CO 2 and CO
__________________
Ex.3 Combustion of Magnesium
 very rapid and exothermic
Example:
*Target Equation: CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
*i.e. the reaction for which the enthalpy change needs to be determined
Pathway #1 (1 step)
∆H1 =
∆Htarget = ∆H1 =
*Pathway #2 (2 steps)
Step 1:
∆H2 =
Step 2:
∆H3 =
SUM:
_________________________________________________
∆Htarget = ∆H2 + ∆H3 =
=
*Notice that Steps 1 and 2 from Pathway #2 “add up” to the target equation.
*Hess’s Law represented mathematically: ∆Htarget = ∑∆Hknown
Example 2:
1. Determine:
i) pathway #1 (1 step)
ii) pathway #2 (2 steps)
2. Show how pathway #2 adds up to pathway #1.
*Show: sum of equation and enthalpies
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Practice Questions:
*Use Hess’s Law to determine the enthalpy change associated with the indicated reaction
1. Show how the molar enthalpy of formation for gaseous ammonia, NH 3(g), can be calculated from the
following reactions.
*First, indicate the Target Equation: 3H2(g) + N2(g) 2NH3(g) or [3/2H2(g) + 1/2N2(g) NH3(g) ∆Hf = ?]
Known equation #1:
NH3(g) +
Known equation #2:
H2(g) +
3
3
1
O2(g)  H2O(g) + N2(g)
4
2
2
1
O2(g)  H2O(g)
2
∆H1 = -316.4 kJ
∆H2 = -241.6 kJ
*Changes to equation #1 
*Changes to equation #2 
2. Nitromethane is a rapid-burning fuel often used in dragsters where rate, not energy yield is important.
4CH3NO2(g) + 3O2(g)  4CO2(g) + 2N2(g) + 6H2O(g)
ΔH˚ = ?
(a) Use Hess’s law to determine the enthalpy change for the reaction above using the following
thermochemical equations. (b) Determine the molar enthalpy of combustion of nitromethane.
C(s) + O2(g)  CO2(g)
2H2(g) + O2(g)  2H2O(g)
2C(s) + 3H2(g) + 2O2(g) + N2(g)  2CH3NO2(g)
ΔH˚ = -393 kJ
ΔH˚ = -484 kJ
ΔH˚ = -226 kJ
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Try  pp. 326-327 #1-3 ; p. 329 #4, 5
5.5 Using Standard Enthalpies of Formation to Determine ΔH
*standard enthalpy of formation  the amount of energy associated with the formation of ____ mole of a
substance from its ____________in their standard states (i.e. their most stable state at SATP)
 SATP = Standard Ambient Temperature (298 K / 25°C) and Pressure (100 kPa)
 STP = Standard Temperature(273 K / 0°C) and Pressure (101.3 kPa)
e.g.
(i) C(s) + O2(g)  CO2(g)
*Must include states.
ΔHf˚= -393.5 kJ/mol
(ii) The formation equation for liquid ethanol:
2C(s) + 3H2(g) +
1
O2(g)  C2H5OH(l)
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*Practice  Write the formation equations for the following compounds:
a) benzene
b) potassium bromate
c) glucose(C6H12O6)
d) magnesium hydroxide
e) acetylene gas (C2H2) f) creatine (C4H9N3O2)
g) potassium iodide h) iron(II) sulfate
Using Enthalpies of Formation to Find ΔH
(A third method) *What are the other two? ______________________ & ______________________
Important:
The ΔHf˚ of an element in its standard state is _____________!
Examples:
The Formula:
ΔH =  n ΔHf˚(Products) -  n ΔHf˚(Reactants)
*Use Table Appendix C6 (pp. 799-800) to answer the following.
Examples:
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