CH 6: Thermochemistry
Renee Y. Becker
Valencia Community College
CHM 1045
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Energy
• Energy: is the capacity to do work, or supply heat.
Energy = Work + Heat
• Kinetic Energy: is the energy of motion.
EK = 1/2 mv2 (1 Joule = 1 kgm2/s2)
(1 calorie = 4.184 J)
• Potential Energy: is stored energy.
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Ek & Ep
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Example 1: KE
Which of the following has the greatest kinetic energy?
1. A 12 kg toy car moving at 5 mph?
2. A 12 kg toy car standing at the top of a large hill?
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Energy
• Thermal Energy is the kinetic energy of
molecular motion
• Thermal energy is proportional to the
temperature in degrees Kelvin. Ethermal T(K)
• Heat is the amount of thermal energy
transferred between two objects at different
temperatures.
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• In an experiment: Reactants and products are
the system; everything else is the surroundings.
• Energy flow from the system to the surroundings has a negative sign (loss of energy). (-E or - H)
• Energy flow from the surroundings to the system has a positive sign (gain of energy). (+E or +H)
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• The law of the conservation of energy: Energy cannot be created or destroyed.
• The energy of an isolated system must be constant.
• The energy change in a system equals the work done on the system + the heat added.
E = Efinal – Einitial = E2 – E1 = q + w
q = heat, w = work8
• Pressure is the force per unit area.
(1 N/m2 = 1 Pa)
(1 atm = 101,325
Pa)
• Work is a force (F) that produces an object’s movement, times the distance moved (d):
Work = Force x Distance
A
F
Area
Force=Pressure
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The expansion in volume that occurs during a reaction forces the piston outward against atmospheric pressure, P.
Work = -atmospheric pressure * area of piston * distance piston moves
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Example 2: Work
How much work is done (in kilojoules), and in which direction, as a result of the following reaction?
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• The amount of heat exchanged between the system and the surroundings is given the symbol q.
q = E + PV
At constant volume (V = 0): qv = E
At constant pressure: qp = E + PV = H
Enthalpy change: H = Hproducts – Hreactants
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Example 3: Work
The explosion of 2.00 mol of solid TNT with a volume of approximately 0.274 L produces gases with a volume of 489 L at room temperature. How much PV (in kilojoules) work is done during the explosion? Assume P = 1 atm, T = 25°C.
2 C7H5N3O6(s) 12 CO(g) + 5 H2(g) + 3 N2(g) + 2 C(s)
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• Enthalpies of Physical Change:
Enthalpy is a state function, the enthalpy change from solid to vapor does not depend on the path taken between the two states.
Hsubl = Hfusion + Hvap 14
• Enthalpies of Chemical Change: Often called heats of reaction (Hreaction).
Endothermic: Heat flows into the system from the surroundings and H has a positive sign.
Exothermic: Heat flows out of the system into the surroundings and H has a negative sign.
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Bromination vs. Chlorination
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• Reversing a reaction changes the sign of H for a reaction.
C3H8(g) + 5 O2(g) 3 CO2(g) + 4 H2O(l) H = –2219 kJ
3 CO2(g) + 4 H2O(l) C3H8(g) + 5 O2(g) H = +2219 kJ
• Multiplying a reaction increases H by the same factor.
3 [C3H8(g) + 15 O2(g) 9 CO2(g) + 12 H2O(l)] H = 3(-2219) kJ
H = -6657 kJ17
Example 4: Heat
• How much heat (in kilojoules) is evolved or absorbed in each of the following reactions?
a) Burning of 15.5 g of propane:
C3H8(g) + 5 O2(g) 3 CO2(g) + 4 H2O(l)
H = –2219 kJ/mole
b) Reaction of 4.88 g of barium hydroxide octahydrate with
ammonium chloride:
Ba(OH)2·8 H2O(s) + 2 NH4Cl(s) BaCl2(aq) + 2 NH3(aq) + 10 H2O(l)
H = +80.3 kJ/mole18
• Thermodynamic Standard State: Most
stable form of a substance at 1 atm pressure
and 25°C; 1 M concentration for all
substances in solution.
• These are indicated by a superscript ° to the
symbol of the quantity reported.
• Standard enthalpy change is indicated by the
symbol H°.
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Example 5:
Is an endothermic reaction a favorable process thermodynamically speaking?
1) Yes
2) No
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Hess’s Law
• Hess’s Law: The overall enthalpy change for a reaction is equal to the sum of the enthalpy changes for the individual steps in the reaction.(not a physical change, chemical change)
3 H2(g) + N2(g) 2 NH3(g) H° = –92.2 kJ
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• Reactants and products in individual steps can be added and subtracted to determine the overall equation.
(1) 2 H2(g) + N2(g) N2H4(g) H°1 = ?
(2) N2H4(g) + H2(g) 2 NH3(g) H°2 = –187.6 kJ
(3) 3 H2(g) + N2(g) 2 NH3(g) H°3 = –92.2 kJ
H°1 + H°2 = H°reaction
Then H°1 = H°reaction - H°2
H°1 = H°3 – H°2 = (–92.2 kJ) – (–187.6 kJ) = +95.4 kJ
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Example 6: Hess’s Law
• The industrial degreasing solvent methylene chloride (CH2Cl2, dichloromethane) is prepared from methane by reaction with chlorine:
CH4(g) + 2 Cl2(g) CH2Cl2(g) + 2 HCl(g)
Use the following data to calculate H° (in kilojoules) for the above reaction:
CH4(g) + Cl2(g) CH3Cl(g) + HCl(g) H° = –98.3 kJ
CH3Cl(g) + Cl2(g) CH2Cl2(g) + HCl(g) H° = –104 kJ
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• Standard Heats of Formation (H°f): The
enthalpy change for the formation of 1 mole of substance in its standard state from its constituent elements in their standard states.
• The standard heat of formation for any element in its standard state is defined as being ZERO.
H°f = 0 for an element in its standard
state24
Standard Heats of Formation
• Calculating H° for a reaction:
H° = H°f (Products) – H°f (Reactants)
• For a balanced equation, each heat of formation must be multiplied by the stoichiometric coefficient.
aA + bB cC + dD
H° = [cH°f (C) + dH°f (D)] – [aH°f (A) + bH°f (B)]25
-1131Na2CO3(s)49C6H6(l)-92HCl(g)
-127AgCl(s)-235C2H5OH(g)95.4N2H4(g)
-167Cl-(aq)-201CH3OH(g)-46NH3(g)
-207NO3-(aq)-85C2H6(g)-286H2O(l)
-240Na+(aq)52C2H4(g)-394CO2(g)
106Ag+(aq)227C2H2(g)-111CO(g)
Some Heats of Formation, Some Heats of Formation, HHff° ° (kJ/mol)(kJ/mol)
Standard Heats of Formation
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Example 7: Standard heat of formation
Calculate H° (in kilojoules) for the
reaction of ammonia with O2 to yield
nitric oxide (NO) and H2O(g), a step in
the Ostwald process for the commercial
production of nitric acid.
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Example 8: Standard heat of formation
Calculate H° (in kilojoules) for the
photosynthesis of glucose and O2 from
CO2 and liquid water, a reaction carried
out by all green plants.
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Example 9:
Which of the following would indicate an endothermic reaction? Why?
1. -H
2. + H
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Heat of Phase Transitions from Hf
Calculate the heat of vaporization, Hvap of water, using standard enthalpies of formation
HfH2O(g) -241.8 kJ/mol
H2O(l) -285.8 kJ/mol
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Calorimetry and Heat Capacity
• Calorimetry is the science of measuring heat changes (q) for chemical reactions. There are two types of calorimeters:
• Bomb Calorimetry: A bomb calorimeter measures
the heat change at constant volume such that q =
E.
• Constant Pressure Calorimetry: A constant
pressure calorimeter measures the heat change
at constant pressure such that q = H. 31
Constant PressureBomb
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Calorimetry and Heat Capacity
• Heat capacity (C) is the amount of heat required to raise the temperature of an object or substance a given amount.
Specific Heat: The amount of heat required to raise
the temperature of 1.00 g of substance by 1.00°C.
q = s x m x t
q = heat required (energy)
s = specific heat
m = mass in grams
t = Tf - Ti
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Calorimetry and Heat Capacity
• Molar Heat: The amount of heat required to raise the temperature of 1.00 mole of substance by 1.00°C.
q = MH x n x t
q = heat required (energy)
MH = molar heat
n = moles
t = Tf - Ti
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Example 10: Specific Heat
What is the specific heat of lead if it takes 96 J to raise the temperature of a 75 g block by 10.0°C?
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Example 11: Specific Heat
How much energy (in J) does it take to increase the temperature of 12.8 g of Gold from 56C to 85C?
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Example 12: Molar Heat
• How much energy (in J) does it take to increase the temperature of 1.45 x104
moles of water from 69C to 94C?
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