Chemistry 30 – Unit 1 Thermochemical Changes
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Transcript of Chemistry 30 – Unit 1 Thermochemical Changes
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Chemistry 30 – Unit 1Thermochemical Changes
To accompany Inquiry into Chemistry
PowerPoint Presentation prepared by Robert [email protected]
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Preparation Info
• Systems: Open, closed, and isolated - definitions
• First Law of Thermodynamics – Total energy of the universe is constant (energy can’t be created or destroyed)
• Second Law of Thermodynamics – In the absence of energy input, a system becomes more disordered (its entropy increases)
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Preparation
• Meaning?
• A system at lower temperature will be more ordered as the particles have less average kinetic energy
• Two systems in thermal contact will transfer energy such that the more ordered (cooler) one gains energy and becomes more disordered
• Consequence: heat always flows from hotter systems to cooler ones
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Preparation
Important Definitions:• Thermal Energy: the total kinetic
energy of all particles of a system
• Temperature: a measure of the average kinetic energy of the particles of a system
• Heat: a transfer of thermal energy between 2 systems
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Chapter 9, Section 9.1Questions:
• Which has more thermal energy, a hot cup of coffee or an iceberg?
• Which has a larger average thermal energy, a hot cup of coffee or an iceberg?
• If an iceberg and a hot cup of coffee come into contact, in which direction will heat flow?
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Preparation
• Heat energy transferred will be related to the temperature change of the system
• It takes different amounts of heat energy to change the temperature of
1 g of a substance by 1°C
• This number is called the specific heat capacity, c, and is measured in units of:
Jg C
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• Water has a c value of
• This means that it takes 4.19 J of heat to raise the temperature of 1 g of water by 1°C
• Water has a very large c compared to most other common substances
4.19 Jg C
Preparation
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• To determine the amount of heat transferred the formula used is
• Despite what your text says on page 337, I would always take ∆t as positive
• If heat is absorbed, temperature of surroundings will decrease; if heat is released temperature of surroundings will increase
• Examples: Practice Problems 1 and 4, page 337
Q mc t
Preparation
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• Practice Problem 1, page 337
• Since 1 J is such a small amount of heat energy I start my questions in kJ as shown above
• If necessary I move into MJ or GJ
0.100 2.44 25 6.1JQ mc t kg C kJgk
k C
Preparation
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• Practice Problem 4, page 337
• Putting kilo top and bottom cancels out and c stays the same
• The substance is granite
• Worksheet: WS 43 (Nelson) then BLM 9.1.1 (back only)
4.9370.790 0.790
0.25000 25.0
Q mc t
kJQ J Jc g Ck
kk g Cm t g C
Preparation
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Chapter 9, Section 9.1
• Energy changes in chemical reactions crucial to life
• Not just in photosynthesis, fuels, and batteries, but in the very way that your body metabolizes food and makes the energy available for life processes
• Thermodynamics: the study of energy and energy changes
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Chapter 9, Section 9.1
• Recall the first law of thermodynamics: ∆Euniverse= 0
• If a system loses energy, the surroundings gain energy (get warmer)
• If a system gains energy, the surroundings lose energy (get cooler)
∆Esystem = - ∆Esurroundings
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Chapter 9, Section 9.1• Energy types:
• Kinetic energy, Ek, energy of motion of particles of a system
• Temperature is a measure of the average Ek of the particles of a system
• Potential energy, Ep, stored energy, usually in chemical bonds
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Chapter 9, Section 9.1• Transfer of Ek: heat flows from hotter objects
to cooler ones (Preparation section of notes)
• Breaking bonds always requires energy (endothermic); forming bonds always releases energy (exothermic)
• Chemical reaction:breaking bonds + energy1 forming bonds + energy2
• If energy1 > energy2, reaction is endothermic
• If reverse is true, it is exothermic
• Worksheet BLM 9.1.3
input output
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Chapter 9, Section 9.1• New term: enthalpy (not entropy)
• Enthalpy (change), ∆H: the difference in potential energy between reactants and products, measured at constant pressure – measured in kJ (or MJ, etc)
• Molar Enthalpy (change), ∆rH: the enthalpy change for 1 mole of a specified substance – measured inkJ/mol (or MJ/mol etc)
• In common usage the word change gets left out
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Chapter 9, Section 9.1• Negative ∆H’s are exothermic
(think lose heat) and temperature of surroundings increases
• Positive ∆H’s are endothermic (think gain heat) and temperature of the surroundings decreases
• Note: this increase → negative, and decrease → positive is a stumbling block for many students
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Chapter 9, Section 9.1• Chemical reactions can be written using
∆H notation:
C6H12O6(s) + 6 O2(g) 6 CO2(g) + 6 H2O(l) ∆H=-2802.5 kJ
4 NO(g) + 6 H2O(g) 4 NH3(g) + 5 O2(g) ∆H=+906 kJ
• They can also be written with the heat as a term in the equation:
C6H12O6(s) + 6 O2(g) 6 CO2(g) + 6 H2O(l) + 2802.5 kJ
4 NO(g) + 6 H2O(g) + 906 kJ 4 NH3(g) + 5 O2(g)
Do ∆H Worksheet!
value for the reaction as written
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Chapter 9, Section 9.1• Potential energy diagrams for the
same 2 reactions are shown below:
∆H = -2802.5 kJ
H (
kJ)
C6H12O6(s) + 6 O2(g)
6 CO2(g) + 6 H2O(l)
reactants
products
H (
kJ)
4 NO(g) + 6 H2O(g)
4 NH3(g) + 5 O2(g)reactant
s
products
∆H = +906 kJ
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Chapter 9, Section 9.2• Recalling that breaking bonds always
endothermic and forming new bonds is always exothermic, more complete Ep diagrams might be shown as follows:
Endothermic Exothermic
reactants
intermediate
products
ΔH
Ep (
kJ)
reactants
intermediate
products
ΔH
Ep (
kJ)
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Chapter 9, Section 9.1• Alternate forms of potential energy diagram
(from Chemistry 30 Diploma Exam Bulletin)
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Chapter 9, Section 9.1• Example: Practice Problem 3, page 346a) C(s) + 2 H2(g) CH4(g) + 74.6 kJ
b) C(s) + 2 H2(g) CH4(g) ∆H = -74.6 kJ
c)
H (
kJ)
C(s) + 2 H2(g)
CH4(g)
products
reactants
∆H = -74.6 kJ
Do Ep diagrams for formation of Cr2O3(s), simple decomp* of AgI(s), and formation of SO2(g)
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Chapter 9, Section 9.2
Formation of Cr2O3(s)
Ep (kJ)
reaction coordinate
2 Cr(s) + 3/2 O2(g)
Cr2O3(s)
ΔH=ˉ1139.7 kJ
Ep (kJ)
reaction coordinate
ΔH=+61.8 kJ
simple decomposition of AgI(s)
AgI(s)
Ag(s) + ½ I2(s) Ep (kJ)
reaction coordinate
ΔH=ˉ296.8 kJ
1/8 S8(s) + O2(g)
SO2(g)
formation of SO2(g)
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Chapter 9, Section 9.1• Molar enthalpy of combustion: the
enthalpy change for the complete combustion of 1 mol of a substance
• Complete combustions of fossil fuels always yields CO2(g) and H2O
• Open systems – constant pressure – gases escape – H2O(g)
• Isolated systems – H2O(l)
• Human body – cellular respiration - H2O(l)
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Chapter 9, Section 9.1• Table of Molar Enthalpies of Combustions
of alkanes, page 347
• Practice Problem 5b, page 347 (open system)
OR: note change in units!
• In thermodynamics it is acceptable to write equations with fractional coefficients – don’t do this elsewhere
• Try question 5a, page 347
C4H10(g) + 13/2 O2(g) 4 CO2(g) + 5 H2O(g) ∆H = -2657.3 kJ
2 C4H10(g) + 13 O2(g) 8 CO2(g) + 10 H2O(g) ∆H = -5314.6 kJ
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Chapter 9, Section 9.1• Question 5a page 347
• Note that the value of ∆H varies directly as the number of moles of reacting substances
• This formula gets used to calculate enthalpy changes for ∆Ep like phase changes, chemical reactions, and nuclear reactions
C5H12(l) + 8 O2(g) 5 CO2(g) + 6 H2O(g) ∆H = -3244.8 kJ
rH n H
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Chapter 9, Section 9.1• Example Practice Problem 3a, page
349
Find for 56.78 g of pentane
56.783244.8 2553
72.17 /
r
r
H n H
H
g kJH n H kJmolg mol
Note: from table, page 347 - comment
mol of pentane
5 123244.8 kJ
molr C HH
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Chapter 9, Section 9.1• Example Practice Problem 6, page
349
• molar enthalpy change for?• a) ammonia
• b) oxygen
• c) nitrogen monoxide
• d) water
4 NH3(g) + 5 O2(g) 4 NO(g) + 6 H2O(g) ΔH = -906 kJ
3
906227
4r NH
kJ kJH molmol
2
906181
5r O
kJ kJH molmol
906227
4r NO
kJ kJH molmol
2
906151
6r H O
kJ kJH molmol
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Chapter 9, Section 9.1• Do Worksheet BLM 9.1.6
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Chapter 9, Section 9.2• Finding the value of energy changes
experimentally: calorimetry
• Device: calorimeter
• The following diagrams show the principle behind calorimetry – note arrow directions
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Chapter 9, Section 9.2• A simple calorimeter like the one you
will use
2 nested styrofoam cups containing a measured volume of water
sitting in a beaker so that it doesn’t fall over
3rd styrofoam cup inverted on top with hole for thermometer (stirrer)
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Chapter 9, Section 9.2• Assumptions in styrofoam cup
calorimetry:• Amount of energy transferred to
cups and thermometer is small and can be ignored
• The system is isolated• The solution produced has the
same density and specific heat capacity as water
• The process occurs at constant pressure
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Chapter 9, Section 9.2• The enthalpy change of a chemical
reaction = energy lost or gained, and is indicated by the symbol ΔH
• Energy gained or lost by the water causes a temperature change and is indicated by the symbol Q
• In an ideal calorimeter ΔH = Q• But recall:
• Therefore
rH n H Q mc t and
rn H mc t calorimetry equation
system calorimeter “water”
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Chapter 9, Section 9.2• I will redo the example on page
354 using this formula
r
mc t mc tH
n cv
limiting reagent, if not stated, or substance question asks about
• remember m c Δt is for the “water” and n (c v) for the CuSO4(aq) 2 0.05000 4.19 24.60 21.40
89.40.300 0.05000
kJkg C kJ
molr molL
kg CH
L
Since the temperature has gone up the process is exothermic
Correct answer: 89.4kJmol
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Chapter 9, Section 9.2• Practice Problem 9, page 355• Note that question asks for molar
enthalpy of reaction for sodium• n will be moles of sodium (question asks)
20.175 4.19 25.70 19.302.9 10
0.3722.99
r
kJkg C kJ
molr
gmol
n H mc t
kg Cmc tH
gn
• Since temperature increases, answer is correctly expressed as
22.9 10 or 0.29kJ MJmol mol
Do Practice Problems 7, 10, 12, page 355
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Chapter 9, Section 9.2• Investigation 9.A page 356 (goes
with the questions you’ve been doing)
• Molar enthalpy of combustion: Investigation 9.B, page 357
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Chapter 9, Section 9.2• Bomb Calorimetry: a bomb
calorimeter is used to make accurate and precise measurements
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Chapter 9, Section 9.2• Reaction takes place inside an
inner container called the “bomb” that contains pure oxygen
• Chemicals are electrically ignited and heat is released to or absorbed from calorimeter water
• Calorimeter materials: stirrer, thermometer, containers are not ignored
• With calorimeter filled to a set level with water, all of their heat capacities are combined as shown:
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Chapter 9, Section 9.2
• Note that C contains the mass and specific heat capacity of each component of the calorimeter
• How do you know when to use
2 2
2 2
r H O H O ther ther stir stir contains contains
r H O H O ther ther stir stir contains contains
r
n H m c t m c t m c t m c t
n H m c m c m c m c t
n H C t
bomb calorimeter equation
versus ?r rn H C t n H mc t
Heat capacity of calorimeter
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Chapter 9, Section 9.2• Look for:
- words “bomb calorimeter”- no mention of the mass or volume of water- words “heat capacity” rather than “specific heat capacity”- units J/°C rather than J/g°C
• Question 2, Worksheet 46
• Since temperature increases, answer is -286 kJ/mol
• Do rest of Worksheet 46
40.00 3.54286
1.002.02
r
kJC kJ
molr
gmol
n H C t
CC tH
gn
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Chapter 9, Section 9.2• More practice with
• WS 9.1.5
Q mc t
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Chapter 9, Section 9.2• Review: page 366-7 good
questions: 1, 3, 4 (no actual calculation needed), 5c (data page 347), 6a (data page 347), 8, 10, 13, 15, 16, 17, 18, 19, 21
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Chapter 9, Section 9.2