chapter 5 thermochemistry - philadelphia university · 2015. 3. 30. · thermochemistry . transfer...
TRANSCRIPT
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Chapter 5 Thermochemistry
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Topics
Energy and energy changes
Introduction to thermodynamics
Enthalpy
Calorimetry
Hess’s Law
Standard enthalpies of formation
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Copyright McGraw-Hill 2009 3
5.1 Energy and Energy Changes
Energy is involved in all types of physical and chemical changes
Energy: the capacity to do work or transfer heat
All forms of energy are either
Kinetic
Potential
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Kinetic energy - energy of motion
Defining equation
Where m is mass and u is velocity
Thermal energy - one form of kinetic energy associated with random motion
•Changes in thermal energy are monitored via changes in temperature
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Potential energy - energy of position or composition
Chemical energy is stored within structural units of chemical substances.
Electrostatic energy is energy resulting from the interaction of charged particles.
•Dependent on charges and distance between charges (Q = charge and d
= distance)
•Defining equation
•+ Eel: repulsive
•− Eel: attractive
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Law of conservation of energy
Energy can be converted from one form to another but not created or destroyed.
The total amount of energy in the universe is constant.
Example
•A chemical reaction (potential) gives off heat (thermal)
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System is the part of the universe of interest.
•Example
The reactants NaOH and HCl
Surroundings are the rest of the universe.
•Example
When heat is given off from the reaction of NaOH and HCl, the energy is transferred from the system to the surroundings.
system: what is inside the container. Reactants or products of a chemical reaction
surroundings: container ,room, etc.
Energy changes in chemical reactions
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The study of the transfer of heat (thermal energy) in chemical reactions.
Exothermic - transfer of heat from the system to the surroundings
2H2(g) + O2(g) 2H2O(l) + energy
Endothermic - the transfer of heat from the surroundings to the system
energy + 2HgO(s) 2Hg(l) + O2(g)
Thermochemistry
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Transfer of Energy
Exothermic process
energy is produced in reaction
energy flows out of system
container becomes hot when touched
CH4(g) + 2O2(g) CO2(g) + 2H2O(g) + heat
Endothermic process
energy is consumed by the reaction
energy flows into the system
container becomes cold when touched
N2(g) + O2(g) + Energy 2NO(g)
http://college.hmco.com/chemistry/shared/media/zumdahl/visual/visual.html?name=CH09&vis=_v9a
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Comparison of Endothermic and Exothermic Processes
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Joule (J) is the SI unit for energy.
•The amount of energy possessed by a 2 kg mass moving at a speed of 1 m/s
Units of Energy
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Calorie (cal) - commonly used on food labels
1 cal 4.184 J
1000 cal = 1 Cal = 1 kcal
Food calories (Cal) are really 1000 calories (cal).
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Calculate the kinetic energy of a neon atom
moving at a speed of 98 m/s.
226
k kg)(98m/s)10(3.3522
1 E
J101.6 22kE
kg103.352g 1000
kg 1
amu
g101.661amu 20.18 26
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Example
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5.2 Introduction to Thermodynamics
Types of systems:
open (exchange of mass and energy)
closed (exchange of energy)
isolated (no exchange)
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State functions depend only on initial and final states of the system and not on how the change was carried out.
Energy (E)
Pressure (P)
Volume (V)
Temperature (T)
State function
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State Function
energy change is independent of pathway thus it is a state function
(state property)
State function: a property of a system that depends only on its present state (it does not depend on system’s past or future, i.e, it does not depend on how the system arrived at the present state)
work and heat depend on pathway so they are not state functions
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First Law of Thermodynamics
Energy can be converted from one form to another but cannot be created or destroyed.
Based on the law of conservation of energy
Internal energy (U)
It is the sum of potential and kinetic energy of all particles in the system
U = PE + KE
Kinetic energy - molecular motion
Potential energy - attractive/repulsive interactions
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The change in internal energy of a system between final (f) and initial (i) states is defined as:
U = Uf Ui For a chemical system
Cannot calculate the total internal energy with any certainty
Can calculate the change in energy of the system experimentally
U = U(products) U(reactants)
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Consider:
S(s) + O2(g) SO2(g)
This reaction releases heat, therefore U is negative.
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Usystem + Usurroundings = 0
When a system releases heat, some of the chemical energy is released as thermal energy to the surroundings but this does not change the total energy of the universe.
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Usystem = Usurroundings
When a system undergoes a change in energy, the surroundings must undergo a change in energy equal in magnitude and opposite in sign.
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where q is heat, w is work
Work and heat
U can be changed (- or + ∆U) by a flow of work, heat, or both
Usys = q + w
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Sign Conventions of q and w
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Calculate the overall change in internal
energy for a system that absorbs 125 J of
heat and does 141 J of work on the
surroundings.
q is + (heat absorbed)
w is (work done)
Usys = q + w = (+125 J) + (141J) = 16 J
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Work
common types of work Expansion- work
done by the system
Compression- work done on the system
A
FP
hAPhFw
VPw
expansion +∆V -w
compression -∆V +w
P is external pressure
Work is force applied over distance
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Pressure-volume work, w, done by a system is
w = PV
Usys = q + w
Usys = q P V
Constant volume, V = 0
qv = Usys qv means a constant-volume process
5.3 Reactions carried out at constant volume
or at constant pressure
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U = q + w
U = qp PV
qp = U + PV
Reactions carried out at constant pressure
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Enthalpy and enthalpy changes
Enthalpy is a property of a system definition: H = U + PV
since U, P and V are all state functions, then H is also a state function
The term enthalpy is composed of the prefix en-, meaning "to put into" and the Greek word -thalpein, meaning "to heat", that is “to put heat into”
http://en.wikipedia.org/wiki/Greek_language
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Enthalpy change
H = U + PV
The change in enthalpy
H = U + (PV)
When pressure is held
constant
H = U + PV
Since
qp = U + PV
qp = H
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At constant pressure
qp = H
Enthalpy of reaction
H is + for endothermic changes.
H is − for exothermic changes.
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Copyright McGraw-Hill 2009
Equations that represent both mass and enthalpy changes
Endothermic reaction H2O(s) H2O(l) H = + 6.01 kJ/mol
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Exothermic reaction
Thermochemical Equations
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Copyright McGraw-Hill 2009 32
Comparison of Endothermic and Exothermic Changes
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Always specify state of reactants and products.
When multiplying an equation by a factor (n), multiply the H value by same factor.
Reversing an equation changes the sign but not the magnitude of the H.
Thermochemical equation guidelines
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Copyright McGraw-Hill 2009 34
Given the following equation,
C6H12O6(s) + 6O2(g) → 6CO2(g) + 6H2O(l)
ΔH = 2803 kJ/mol
calculate the energy released when 45.00 g of glucose is burned in oxygen.
kJ 700.0OHC mol 1
kJ 2803
OHC g 180.2
OHC mol 1OHC g 45.00
61266126
61266126
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5.4 Calorimetry
The science of measuring
heat changes
calorimeter- device used to experimentally find the heat associated with a chemical reaction
substances respond differently when heated ( to raise T for two substances by 1 degree, they require different amount of heat)
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Specific heat (s)
Specific heat (s) - the amount of heat required to raise the temp of 1 g of a substance by 1C.
Units: J/g C
Relation to amount of heat (q)
where q is heat, m is mass, s is specific heat
and T = change in temp
(T = Tfinal Tinitial)
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Copyright McGraw-Hill 2009 37
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Copyright McGraw-Hill 2009 38
Heat capacity (C) - the amount of heat required to raise the temp of an object by 1C.
Units: J/C
Relation to amount of heat (q)
where q is heat, C is heat capacity
and T = change in temp
(T = Tfinal Tinitial)
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Calculate the amount of energy required
to heat 95.0 grams of water from 22.5C
to 95.5C.
T = Tfinal – Tinitial = 95.5 oC − 22.5oC
T = 73.0 oC
q = (95.0 g) (4.184 J/gC) (73.0C)
q = 2.90 x 104 J or 29.0 kJ
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Constant-Pressure Calorimetry
uses simplest calorimeter (like coffee-cup calorimeter) since it is open to air
used to find heat exchange
between the system and
surroundings of a chemical
reaction
i.e., changes in enthalpy
for reactions occurring
in a solution
since qP = ∆H
System: Reactants & Products
Surroundings: the water
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System: reactants and products
Surroundings: water
Assuming that the calorimeter
does not leak or absorb heat
coffee cup calorimeter
Constant-pressure calorimetry
For an exothermic reaction:
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Copyright McGraw-Hill 2009 42
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Copyright McGraw-Hill 2009 43
A metal pellet with a mass of 85.00 grams at an original
temperature of 92.5C is dropped into a calorimeter\
with 150.00grams of water at an original temperature of
23.1C. The final temperature of the water and the pellet
Is 26.8C. Calculate the heat capacity and the specific
heat for the metal.
Example
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qwater = msT
= (150.00 g) (4.184 J/gC) (3.7C)
= 2300 J (water gained energy)
= -2300 J (pellet released energy)
Heat capacity of pellet: q = CT
C = q/T
= 2300 J/65.7C = 35 J/C
Specific heat of pellet: J/goC
s =35 J/oC
85.00g= 0.41J/goC
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bomb calorimeter
Constant-volume calorimetry
Isolated system
Here, the ∆V = 0 so -P∆V = w = 0
∆U = q + w = qV for constant volume
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Typical procedure used in a bomb calorimeter
•Known amount of sample placed in steel
container and then filled with oxygen gas
•Steel chamber submerged in known amount of water
•Sample ignited electrically
•Temperature increase of water is determined
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Copyright McGraw-Hill 2009 47
A snack chip with a mass of 2.36 g was burned in a bomb calorimeter. The heat capacity of the calorimeter 38.57 kJ/C. During the combustion the water temp rose by 2.70C. Calculate the energy in kJ/g for the chip
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Copyright McGraw-Hill 2009 48
qrxn = − Ccal T
= −(38.57 kJ/C) (2.70C)
= − 104 kJ
Energy content is a positive quantity.
= 104 kJ/2.36 g
= 44.1 kJ/g
Food Calories: 10.5 Cal/g
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Example 3
Compare the energy released in the combustion of H2 and CH4 carried out in a bomb calorimeter with a heat capacity of 11.3 kJ/°C. The combustion of 1.50 g of methane produced a T change of 7.3°C while the combustion of 1.15 g of hydrogen produced a T change of 14.3°C. Find the energy of combustion per gram for each.
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Example 3 methane: CH4
hydrogen: H2
The energy released by H2 is about 2.5 times the energy released by CH4
gkJgkJH
kJCC
kJH
TCH rcalorimete
/555.1/83
83)3.7()3.11(
gkJgkJH
kJCC
kJH
TCH rcalorimete
/14115.1/162
162)3.14()3.11(
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5.5 Hess’s Law
Hess’s Law: The change in enthalpy that occurs when reactants are converted to products is the same whether the reaction occurs in one step or a series of steps.
It is sed for calculating enthalpy for a reaction that cannot be determined directly.
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Example 1
N2(g) + 2O2(g) 2NO2(g) ∆H = 68 kJ
OR
N2(g) + O2(g) 2NO(g) ∆H = 180 kJ
2NO(g) + O2(g) 2NO2(g)∆H = -112 kJ
N2(g) + 2O2(g) 2NO2(g) ∆H = 68 kJ
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Rules
1. If a reaction is reversed, the sign of ∆H must be reversed as well.
because the sign tells us the direction of heat flow at constant P
2. The magnitude of ∆H is directly proportional to quantities of reactants and products in reaction.
If coefficients are multiplied by an integer, the ∆H must be multiplied in the same way.
because ∆H is an extensive property
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Copyright McGraw-Hill 2009 54
Given the following equations:
H3BO3(aq) HBO2(aq) + H2O(l) Hrxn = 0.02 kJ
H2B4O7(aq) + H2O(l) 4 HBO2(aq) Hrxn = 11.3 kJ
H2B4O7(aq) 2 B2O3(s) + H2O(l) Hrxn = 17.5 kJ
Find the H for this overall reaction.
2H3BO3(aq) B2O3(s) + 3H2O(l)
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2H3BO3(aq) 2HBO2(aq) + 2H2O(l) x 2
Hrxn = 2(−0.02 kJ) = −0.04 kJ
2HBO2(aq) 1/2H2B4O7(aq) + 1/2H2O(l) reverse, ÷2
Hrxn = +11.3 kJ/2 = 5.65 kJ
1/2H2B4O7(aq) B2O3(s)+ 1/2H2O(l) ÷ 2
Hrxn = 17.5 kJ/2 = 8.75 kJ
2H3BO3(aq) B2O3(s) + 3H2O(l)
Hrxn = 14.36 kJ
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Copyright McGraw-Hill 2009 56
5.6 Standard Enthalpies of Formation
Symbol: Hf
The enthalpy change that results when 1 mole of a compound is formed from its elements in their standard states.
Hf for an element in its standard
state is defined as zero.
Standard state: 1 atm, 25C
Values found in reference tables
Used to calculate the Hrxn
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Standard States
For a compound:
for gas: P = 1 atm
For pure substances, it is a pure liquid or pure solid state
in solution: concentration is 1 M
For an element:
form that exists in at 1 atm and 25°C
O: O2(g) K: K(s) Br: Br2(l)
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Copyright McGraw-Hill 2009 58
Hrxn = nHf
(products) mHf
(reactants)
where denotes summation
n and m are the coefficients in the
balanced equation
Defining equation for enthalpy of reaction:
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Calculate the Hrxn for the following reaction
from the table of standard values.
CH4(g) + 2O2(g) CO2(g) + 2H2O(l)
Hrxn = nHf
(products) - mHf
(reactants)
= [1(393.5) + 2(285.8)] [1(74.8) + 2(0)]
= 890.3 kJ/mol (exothermic)
Example:
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Copyright McGraw-Hill 2009 60
Key Points
First law of thermodynamics
Enthalpy (heat of formation; heat of reaction)
State function
Calorimetry
Specific heat
Hess’s law
Calculations involving enthalpy, specific heat, energy