2. UNIT 3: Energy Changes and Rates of Reaction
Chapter 5: Energy Changes
Chapter 6: Rates of Reaction

3. Chapter 5: Energy Changes
UNIT 3
Chapter 5: Energy Changes
Chapter 5: Energy Changes
Energy changes are involved in every chemical reaction—even those that occur in our bodies. While many processes involve the release of energy, others are accompanied by the absorb energy.
Image source: MHR, Chemistry 12 © ISBN ; page 276
Sarah Reinertsen is a competitive runner. She relies a great deal on energy produced by the chemical reactions in her cells.
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4. 5.1 The Nature of Energy and Heat
UNIT 3
Chapter 5: Energy Changes
Section 5.1
5.1 The Nature of Energy and Heat
Chemical and physical changes in matter involve changes in energy content. These changes in energy involve the release or absorption of energy.
Thermochemistry is a branch of chemistry that studies heat released during chemical and physical processes.
Image source: MHR, Chemistry 12 © ISBN ; page 278
Image source: MHR, Chemistry 12 © ISBN ; page 208
Burning wood involves changes to the matter that make up the wood. A change in energy content also occurs. Energy is released, mostly as heat and light.
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5. Some Foundational Concepts for Thermochemistry
UNIT 3
Chapter 5: Energy Changes
Section 5.1
Some Foundational Concepts for Thermochemistry
Two categories of energy:
Kinetic Energy
energy of motion
anything moving has kinetic energy
Potential Energy
stored energy due to the condition or position of the object
SI unit of energy is the joule, J (1 kJ = 1000 J)
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6. System and Surroundings
UNIT 3
Chapter 5: Energy Changes
Section 5.1
System and Surroundings
universe = system + surroundings
the system is the sample being observed
the surroundings is everything else
interactions between a system and its surroundings involve exchange of energy and matter
Three types of systems based on this type of exchange are:
Image source: MHR, Chemistry 12 © ISBN ; page 279
can exchange energy and matter with surroundings
can exchange energy, not matter, with surroundings
cannot exchange matter or energy with surroundings
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7. Calculating the Amount of Heat Entering and Leaving a System
UNIT 3
Chapter 5: Energy Changes
Section 5.1
Calculating the Amount of Heat Entering and Leaving a System
Measurable properties of a system include:
volume, mass, pressure, temperature, and specific heat capacity.
A calculated property of a system includes heat (Q) that enters or leaves an object.
Image source: MHR, Chemistry 12 © ISBN ; page 280
ΔT = Tfinal – Tinitial
When heat enters a system, ΔT is positive and so is Q.
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8. Answer on the next slide
UNIT 3
Chapter 5: Energy Changes
Section 5.1
Learning Check
A 1.0 g sample of copper is heated from 25.0°C to 31.0°C. How much heat did the sample absorb?
Answer on the next slide
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9. Learning Check UNIT 3 Q = mc∆T Q = (1.0 g)(0.385 g/J°C)(6.0°C)
Chapter 5: Energy Changes
Section 5.1
Learning Check
Q = mc∆T
Q = (1.0 g)(0.385 g/J°C)(6.0°C)
Q = 2.3 J
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10. The First Law of Thermodynamics: Energy is Conserved
UNIT 3
Chapter 5: Energy Changes
Section 5.1
The First Law of Thermodynamics: Energy is Conserved
The first law of thermodynamics states that:
energy can be converted from one form to another but cannot be created or destroyed.
Since any change in energy of the universe must be zero,
ΔEuniverse = ΔEsystem + ΔEsurroundings = 0
ΔEsystem = –ΔEsurroundings
if a system gains energy, that energy comes from the surroundings
if a system loses energy, that energy enters the surroundings
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11. UNIT 3
Chapter 5: Energy Changes
Section 5.1
Enthalpy
One way chemists express thermochemical changes is by a variable called enthalpy, H. The change in enthalpy, ΔH, of a system can be measured.
It depends only on the initial and final states of the system, and is represented by
ΔH = ΔE + Δ(PV)
For reactions of solids and liquids in solutions, Δ(PV) = 0
If heat enters a system
ΔH is positive
the process is endothermic
If heat leaves a system
ΔH is negative
the process is exothermic
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12. The Second Law of Thermodynamics
UNIT 3
Chapter 5: Energy Changes
Section 5.1
The Second Law of Thermodynamics
The second law of thermodynamics states that:
when two objects are in thermal contact, heat is transferred from the object at a higher temperature to the object at the lower temperature until both objects are the same temperature (in thermal equilibrium)
Image source: MHR, Chemistry 12 © ISBN ; page 283
When in thermal contact, energy from hot particles will transfer to cold particles until the energy is equally distributed and thermal equilibrium is reached.
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13. Comparing Categories of Enthalpy Changes: Enthalpy of Solution
UNIT 3
Chapter 5: Energy Changes
Section 5.1
Comparing Categories of Enthalpy Changes: Enthalpy of Solution
Three processes occur when a substance dissolves, each with a ΔH value.
bonds between solute molecules or ions break
bonds between solvent molecules break
bonds between solvent molecules and solute molecules or ions form
Sum of the enthalpy changes: enthalpy of solution, ΔHsolution
The three fundamental types of processes in which enthalpy change is considered are: physical, chemical, and nuclear.
Recall: physical changes are changes to the condition of a substance that do not change the chemical properties of the substance.
Image source: MHR, Chemistry 12 © ISBN ; page 284
The orange arrow shows the overall ΔH.
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14. Comparing Categories of Enthalpy Changes: Enthalpy of Phase Changes
UNIT 3
Chapter 5: Energy Changes
Section 5.1
Comparing Categories of Enthalpy Changes: Enthalpy of Phase Changes
Heat must be added to or removed from a substance in order for the phase of the substance to change.
The ΔH for each phase change has a particular symbol. For example, ΔHmelt is called the enthalpy of melting.
Image source: MHR, Chemistry 12 © ISBN ; page 285
The ΔH for one phase change is the negative of the ΔH for the opposite phase change.
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15. Comparing Categories of Enthalpy Changes: Chemical Changes
UNIT 3
Chapter 5: Energy Changes
Section 5.1
Comparing Categories of Enthalpy Changes: Chemical Changes
Every chemical reaction is associated with an enthalpy of reaction, ΔHr.
For example:
The combustion of methane is
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
For every mole of methane combusted ΔHr = –890 kJ
Enthalpy of reaction is discussed in much more detail in sections 5.2 and 5.3.
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16. Comparing Categories of Enthalpy Changes: Nuclear Changes
UNIT 3
Chapter 5: Energy Changes
Section 5.1
Comparing Categories of Enthalpy Changes: Nuclear Changes
Alpha and Beta Decay:
Atoms of some elements emit alpha (α) or beta (β) particles, transforming them into other elements with smaller masses.
Nuclear Fission:
a heavier nucleus is split into smaller nuclei
energy is released and the products are radioactive
Alpha particles are identical to helium nuclei, with two protons and two electrons. Beta particles are identical to electrons.
Examples of elements that spontaneously emit these particles are shown on the slide.
An example of nuclear fission is shown on the slide. When uranium-235 is bombarded with neutrons, it splits into krypton-92, barium-141, and three neutrons.
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17. Comparing Categories of Enthalpy Changes: Nuclear Changes
UNIT 3
Chapter 5: Energy Changes
Section 5.1
Comparing Categories of Enthalpy Changes: Nuclear Changes
Nuclear Fusion:
two small nuclei combine to form a larger nucleus
the most common reaction in the Sun is the fusion of deuterium and tritium
Image source: MHR, Chemistry 12 © ISBN ; page 289
In this nuclear fusion reaction the highly unstable helium-5 breaks down to helium-4 and a neutron.
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18. Comparing Enthalpy Changes
UNIT 3
Chapter 5: Energy Changes
Section 5.1
Comparing Enthalpy Changes
ΔH for physical changes, such as phase changes and dissolving of a solute in solvent, are the smallest
ΔHr are much larger than ΔH for physical changes, but smaller than ΔH for nuclear changes.
Image source: MHR, Chemistry 12 © ISBN ; page 290 Table 5.3
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19. Section 5.1 Review UNIT 3 Chapter 5: Energy Changes Section 5.1
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20. 5.2 Thermochemical Equations and Calorimetry
UNIT 3
Chapter 5: Energy Changes
Section 5.2
5.2 Thermochemical Equations and Calorimetry
Chemical reactions involve
initial breaking of chemical bonds (endothermic)
then formation of new bonds (exothermic)
ΔHr is the difference between the total energy required to break bonds and the total energy released when bonds form.
Image source: MHR, Chemistry 12 © ISBN ; page 292
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21. Thermochemical Equations
UNIT 3
Chapter 5: Energy Changes
Section 5.2
Thermochemical Equations
Thermochemical equations include the enthalpy change.
Exothermic reaction:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) kJ
Endothermic reaction:
N2(g) + 2O2(g) kJ → 2NO2(g)
The enthalpy term can also be written beside the equation.
NOTE: It is important to understand the meaning of the energy term that is given in thermochemical equations. The equations are written with the assumption that the coefficients of reactants and products represent the amount, in moles, of each substance that is involved in the reaction. The proper way to read the combustion of methane reaction is
“1 mol of methane and 2 mol of oxygen react to produce 1 mol of carbon dioxide, 2 mol of water, and kJ of heat”
Another way to read the equation is
“ the enthalpy change of the reaction between methane and oxygen, as written, is –890.8 kJ
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) ΔHr = –890.8 kJ
N2(g) + 2O2(g) → 2NO2(g) ΔHr = kJ
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22. Enthalpy Diagrams UNIT 3
Chapter 5: Energy Changes
Section 5.2
Enthalpy Diagrams
Enthalpy diagrams clearly show the relative enthalpies of reactants and products.
For exothermic reactions, reactants have a larger enthalpy than products and are drawn above the products.
For endothermic reactions, the products have a larger enthalpy and are drawn above the reactants.
Image source: MHR, Chemistry 12 © ISBN ; page 294
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23. Molar Enthalpy of Combustion
UNIT 3
Chapter 5: Energy Changes
Section 5.2
Molar Enthalpy of Combustion
Combustion reactions have their own symbol: ΔHcomb
For standard molar enthalpy of combustion data in reference tables:
reactants and products are at standard conditions
refer only to the compound undergoing combustion
are for combustion of 1 mol of the compound being combusted, not for equations “as written.” For some compounds to be written with a coefficient of 1, equations are written with fractions as coefficients
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24. Reactant Amounts and Enthalpy of Reaction
UNIT 3
Chapter 5: Energy Changes
Section 5.2
Reactant Amounts and Enthalpy of Reaction
The enthalpy of a reaction is directly proportional to the amount of substance that reacts.
The combustion of 2 mol of propane releases twice as much energy as the combustion of 1 mol of propane.
Knowing the thermochemical equation, the enthalpy change associated with a certain amount of reactants or products can be calculated.
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25. Answer on the next slide
UNIT 3
Chapter 5: Energy Changes
Section 5.2
Learning Check
Determine ΔHcomb for 15.0 g of propane.
Answer on the next slide
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26. Learning Check UNIT 3 The molar mass of propane: 44.1 g/mol
Chapter 5: Energy Changes
Section 5.2
Learning Check
The molar mass of propane: 44.1 g/mol
Therefore, 15.0 g is mol
ΔHcomb = ΔH°comb
ΔHcomb = (0.340 mol) (– kJ/mol)
ΔHcomb = 755 kJ/mol
The value for the standard molar enthalpy of combustion of propane is from Table 5.4 on page 295.
Image source: MHR, Chemistry 12 © ISBN ; page 98
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27. Using Calorimetry To Study Energy Changes
UNIT 3
Chapter 5: Energy Changes
Section 5.2
Using Calorimetry To Study Energy Changes
Calorimeters are used to measure heat released or absorbed by a chemical or physical process.
The water of a calorimeter is considered one system and the container in which the process occurs another system. These two systems are in thermal contact.
Image source: MHR, Chemistry 12 © ISBN ; page 300
For an endothermic process, heat will be transferred from the water to the process system. The temperature of the water will decrease.
For an exothermic process, heat will be transferred from the process to the water. The temperature of the water will increase.
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28. Using a Simple Calorimeter
UNIT 3
Chapter 5: Energy Changes
Section 5.2
Using a Simple Calorimeter
nested polystyrene cups can be used (good insulators)
a known mass of water is in the inner cup, where the process occurs
the process often involves compounds dissolved in water
the change in temperature of the water is measured; the solution absorbs or releases energy
the solution is dilute enough so that the specific heat capacity of water is used
Image source: MHR, Chemistry 12 © ISBN ; page 301
A simple calorimeter.
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29. Using a Simple Calorimeter
UNIT 3
Chapter 5: Energy Changes
Section 5.2
Using a Simple Calorimeter
If the “system” is the process being studied and the “surroundings” is the water in the calorimeter:
The change in thermal energy caused by the process can be calculated:
Q = mcΔT
m is the mass of the water, c is the specific heat capacity of water, and ΔT is the change in temperature of the water
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30. How to Use a Simple Calorimeter
UNIT 3
Chapter 5: Energy Changes
Section 5.2
How to Use a Simple Calorimeter
Note: Material on this slide is taken directly from page 302 of the textbook. It can be a reference when using a simple calorimeter.
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31. Using Calorimetry Data To Determine the Enthalpy of Reaction
UNIT 3
Chapter 5: Energy Changes
Section 5.2
Using Calorimetry Data To Determine the Enthalpy of Reaction
Q = mwcwΔTw is the thermal energy absorbed or released by the water.
The opposite sign of that value is the thermal energy released or absorbed by the system (compounds in solution).
Since pressure is constant, Q exchanged by the system and the water is the same as the enthalpy change of the system ΔH. This is used to determine molar enthalpy change, ΔHr.
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32. Using Flame Calorimetry To Determine the Enthalpy of Combustion
UNIT 3
Chapter 5: Energy Changes
Section 5.2
Using Flame Calorimetry To Determine the Enthalpy of Combustion
Flame calorimeters are flame-resistant and often made of metals cans
A flame calorimeter
is used for determining ΔHcomb
absorbs a great deal of energy, which must be included in energy calculations
is used for burning impure materials like food; ΔHcomb is reported in kJ/g
Image source: MHR, Chemistry 12 © ISBN ; page 306
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33. Using Bomb Calorimetry To Measure Enthalpy Changes during Combustion
UNIT 3
Chapter 5: Energy Changes
Section 5.2
Using Bomb Calorimetry To Measure Enthalpy Changes during Combustion
A bomb calorimeter
is used for more accurately determining ΔHcomb
determines ΔHcomb at constant volume
has a particular heat capacity, C
Image source: MHR, Chemistry 12 © ISBN ; page 307
Q = CΔT is used for bomb calorimetry calculations
Bomb calorimeters are much more sophisticated than flame calorimeters or simple calorimeters.
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34. Answer on the next slide
UNIT 3
Chapter 5: Energy Changes
Section 5.2
Learning Check
A chemical reaction is carried out in a dilute aqueous solution using a simple calorimeter. Calculate the enthalpy change of the reaction using the data below.
mass of solution in calorimeter: 2.00 х 102 g
specific heat capacity of water: 4.19 J/g°C
Initial temperature: 15.0°C
Final temperature: 19.0°C
Answer on the next slide
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35. Learning Check UNIT 3 Q = msolutioncsolution∆Tsolution
Chapter 5: Energy Changes
Section 5.2
Learning Check
Q = msolutioncsolution∆Tsolution
Q = (2.00 х 102 g)(4.19 J/g°C)(4.0°C)
Q = 3.35 х 103 J
This is the thermal change of the “surroundings.”
Image source: MHR, Chemistry 12 © ISBN ; page 98
∆Hsystem = –Q
∆Hsystem = –3.35 kJ
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36. Section 5.2 Review UNIT 3 Chapter 5: Energy Changes Section 5.2
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37. UNIT 3
Chapter 5: Energy Changes
Section 5.3
5.3 Hess’s Law
The enthalpy change of nearly any reaction can be determined using collected data and Hess’s law.
The enthalpy change of any reaction can be determined if:
the enthalpy changes of a set of reactions “add up to” the overall reaction of interest
standard enthalpy change, ΔH°, values are used
Image source: MHR, Chemistry 12 © ISBN ; page
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38. Combining Sets of Chemical Equations
UNIT 3
Chapter 5: Energy Changes
Section 5.3
Combining Sets of Chemical Equations
To find the enthalpy change for formation of SO3 from O2 and S8, you can use
Because only 1 mol of sulfur trioxide is in the final step:
Divide the first equation by 8 so 1 mol of sulfur dioxide is in the first step
Divide the second equation by 2
Image source: MHR, Chemistry 12 © ISBN ; page
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39. Techniques for Manipulating Equations
UNIT 3
Chapter 5: Energy Changes
Section 5.3
Techniques for Manipulating Equations
Reverse an equation
the products become the reactants, and reactants become the products
the sign of the ΔH value must be changed
Multiply each coefficient
all coefficients in an equation are multiplied by the same integer or fraction
the value of ΔH must also be multiplied by the same number
Image source: MHR, Chemistry 12 © ISBN ; page
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40. Standard Molar Enthalpies of Formation
UNIT 3
Chapter 5: Energy Changes
Section 5.3
Standard Molar Enthalpies of Formation
Data that are especially useful for calculating standard enthalpy changes:
standard molar enthalpy of formation, ΔH˚f
the change in enthalpy when 1 mol of a compound is synthesized from its elements in their most stable form at SATP conditions
enthalpies of formation for elements in their most stable state under SATP conditions are set at zero
since formation equations are for 1 mol of compound, many equations include fractions
standard molar enthalpies of formation are in Appendix B
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41. Formation Reactions and Thermal Stability
UNIT 3
Chapter 5: Energy Changes
Section 5.3
Formation Reactions and Thermal Stability
The thermal stability of a substance is the ability of the substance to resist decomposition when heated.
decomposition is the reverse of formation
the opposite sign of an enthalpy change of formation for a compound is the enthalpy change for its decomposition
the greater the enthalpy change for the decomposition of a substance, the greater the thermal stability of the substance
Image source: MHR, Chemistry 12 © ISBN ; page
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42. Using Enthalpies of Formation and Hess’s Law
UNIT 3
Chapter 5: Energy Changes
Section 5.3
Using Enthalpies of Formation and Hess’s Law
For example:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
Image source: MHR, Chemistry 12 © ISBN ; page
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43. UNIT 3
Chapter 5: Energy Changes
Section 5.3
Learning Check
Determine ∆H˚r for the following reaction using the enthalpies of formation that are provided.
C2H5OH(l) + 3O2(g) → 2CO2(g) + 3H2O(l)
∆H˚f of C2H5OH(l): –277.6 kJ/mol
∆H˚f of CO2(g): –393.5 kJ/mol
∆H˚f of H2O(l): –285.8 kJ/mol
Answer on the next slide
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44. UNIT 3
Chapter 5: Energy Changes
Section 5.3
Learning Check
∆H˚r = [(2 mol)(∆H˚f CO2(g)) + (3 mol)(∆H˚f H2O(l))] –
[(1 mol)(∆H˚f C2H5OH(l)) + (3 mol)(∆H˚fO2(g)]
∆H˚r = [(2 mol)(–393.5 kJ/mol) + (3 mol)(–285.8 kJ/mol)] –
[(1 mol)(–277.6 kJ/mol) + (3 mol)(0 kJ/mol)]
Image source: MHR, Chemistry 12 © ISBN ; page 98
∆H˚r = (– kJ) – (–277.6 kJ)
∆H˚r = – kJ
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45. Section 5.3 Review UNIT 3 Chapter 5: Energy Changes Section 5.3
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46. 5.4 Energy Efficiency and Energy Resources
UNIT 3
Chapter 5: Energy Changes
Section 5.4
5.4 Energy Efficiency and Energy Resources
Energy efficiency can be calculated using the equation:
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47. Using Energy Efficiently
UNIT 3
Chapter 5: Energy Changes
Section 5.4
Using Energy Efficiently
A challenge in the development of energy efficient technology is to find ways to best convert energy input into useful forms.
For example, efficiency of appliances:
conversion of input of electrical energy versus output of energy usually all that is considered
but should also consider efficiency of the source of the electricity
Image source: MHR, Chemistry 12 © ISBN ; page 328
Energy use distribution in Canadian homes
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48. Conventional Energy Sources in Ontario
UNIT 3
Chapter 5: Energy Changes
Section 5.4
Conventional Energy Sources in Ontario
The three main sources of electrical energy in Ontario:
nuclear power plants
power plants that burn fossil fuels (natural gas and coal)
hydroelectric generating stations
Image source: MHR, Chemistry 12 © ISBN ; page
The distribution of energy sources in Ontario
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49. Alternative Renewable Energy Sources in Ontario
UNIT 3
Chapter 5: Energy Changes
Section 5.4
Alternative Renewable Energy Sources in Ontario
Renewable energy sources in Ontario:
account for about 25% of energy production
are projected to increase to as high as 40% by 2025
include hydroelectric power (major source), wind energy (currently ~ 1% and projected to 15% in 2025), and solar energy (currently low but may be as high as 5% in 2025). Much lower contributors are biomass, wave power, and geothermal energy.
Image source: MHR, Chemistry 12 © ISBN ; page
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50. What Is a “Clean” Fuel? UNIT 3
Chapter 5: Energy Changes
Section 5.4
What Is a “Clean” Fuel?
Different fuels have differing impacts on the environment. One way this impact is measured is through emissions.
For example, CO2(g) emissions per kJ of energy produced.
Fuel
kg CO2/ kJ energy
Anthracite coal
108.83
Oil
78.48
Natural gas
56.03
Nuclear
0.00
Renewables
Image source: MHR, Chemistry 12 © ISBN ; page 333 Table 5.6
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51. Section 5.4 Review UNIT 3 Chapter 5: Energy Changes Section 5.4
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