Download Chapter 6 - Saint Leo University Faculty

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Chapter 6
Thermochemistry
I) Why do chemical reactions occur?
1) Relationship between energy & stability. Substances that are highly energetic
are less stable and more reactive than substances that are less energetic.
2) For a reaction to take place spontaneously, the energy of the products must be
lower than the energy of the reactants.
3) Thermochemistry investigates the absorption or release of heat energy that
accompanies chemical reactions.
II) Definition of Energy
1) Energy is defined as the capacity to do work or supply heat.
Work is the result of a force acting through a distance.
Heat is the flow of energy caused by a temperature difference.
Energy = Heat + Work
2) Energy is broken down into several broad classes
A) Kinetic Energy
- Energy of motion (depends on object’s mass and velocity)
- KE = 1/2 mv2
B) Potential Energy
- Stored energy due to position or composition
- PE = mgh
C) Thermal Energy
- Type of kinetic energy associated with temperature of objects
(recall that objects are in constant motion; things move faster as
temperature rises).
D) Chemical Energy
- Type of potential energy associated with bond making and bond
breaking in chemical reactions. More on this later.
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3) Units for Energy
A) Joule = SI unit for energy (1 J = 1 kg m2 / sec2 )
B) Common units in the lab are kiljoules.
C) Another unit for energy is the calorie which is the amount of energy
required to raise the temperature of 1 g of a substance by 1 degree Celsius.
-
1 calorie = 4.184 J (exactly)
D) Do not confuse the calorie with a food calorie (Cal). A food calorie is
the same as 1000 calories (1 kilocalorie).
-
1 Cal = 1000 cal = 1 kcal
III) Law of Conservation of Matter and Energy
1) Law of Conservation of Matter also applies to energy.
A) Energy can be converted from one form into another but can not be
created nor destroyed
2) The total energy of the universe is constant (First Law of Thermodynamcis).
3) Types of Energy
A) Thermal energy studies the kinetic energy of molecular motion.
- Temperature: property that reflects random motion of particles in
a substance.
- Heat: transfer of energy between 2 particles as a result of a T
B) Chemical Energy
- Stored energy found within the chemical bonds of molecules.
4) The study of energy and its interconversions is called thermodynamics.
IV) Thermodynamics Basics:
1) Essential Terminology
A) System – the substances we focus on in an experiment.
B) Surroundings – everything outside the system.
C) Universe – sum of the sytem and surroundings.
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2) Thermodynamic Quantities
- Thermodynamic quantities described by three things.
i) magnitude (amount)
ii) unit (description of type of value)
iii) sign (direction of flow)
- Direction of flow always reflects system’s point of view
3) Internal Energy: A Thermodynamic Quantity
A) Let E = Internal Energy (state function) =  PE +  KE
- E is the sum of KE and PE of all “particles” in system. Energy
can be changed by flow of heat, work, or both.
- First Law of Thermodynamics
E=q+w
q = heat
w = work
- Reaction Pathway
i) conditions
ii) change in Energy is independent of pathway
iii) work and heat are not independent
B) E is a state function
- State functions depend ONLY on their present state, not
on how it arrived there
- change in property in going from one state to another is
independent of particular pathway taken.
- work & heat are not state functions.
IV) Work
1) Work is defined as the force that produces movement of an object over a
certain distance.
2) Work (w)
- w = (-)
system does work on surroundings
(E flows out of system)
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- w = (+)
surroundings do work on system
(E flows into system)
-Types of Work
Expansion (Compression of a Gas)
Work = -PV
V) Heat
1) Exothermic vs. Endothermic Reactions
A) Exothermic Rxns evolve heat (E flows out of the system)
CH4 (g) + O2 (g) -----> CO2 (g) + 2 H2O (g) + heat energy
q = (-) EXOTHERMIC
Esys decreases
B) Endothermic Rxns absorb heat (E flows into system)
N2 (g) + O2 (g) + heat energy ------> 2 NO (g)
q = (+) ENDOTHERMIC
Esys increases
2) Where does Energy come from?
A) Energy is always conserved (Law of Conservation)
Esys = Esurr
- Energy required to break bonds
- Energy released when bonds are formed
- Endothermic Rxn: Reactants more stable (lower E) than
products.
- Exothermic Rxn:
reactants.
Products more stable (lower E) than
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VI) Enthalpy (H)
1) Enthalpy (H) is a state function
(E, P, and V are also state functions)
2)Definition of Enthalpy (Mathematical Derivation)
H = E + PV
E = internal energy
E = qp + w
qp is heat associated with constant P
w = - PV
P is constant external pressure
E = qp - PV
qp = E + PV = H
H = qp
At constant pressure, H = flow of energy as heat.
H = +
H = -
endothermic process
exothermic process
Hrxn =  Hproducts -  Hreactants
3) Standard Enthalpies of Formation Hf0
A) Change in enthalpy that accompanies the formation of 1 mole of a
compound from its elements with all substances in their standard states.
Hreaction0 =  nf Hf0 (products) -  nr Hr0 (reactants)
B) Standard states (Conventional Definitions)
Hf0 for elements in their standard states is defined as zero.
Hf0 for substances
(T = 25 C)
Gases: pressure is 1.0 atm
Solutions: concentration is 1.0 M
Pure liquids & solids: standard state is pure liquid or solid.
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4) Enthalpies of Physical Changes
A) Enthalpy of Fusion (Melting)
solid  liquid
B) Enthalpy of Vaporization (Boiling)
liquid  gas
C) Enthalpy of Sublimation
solid  gas
D) Expressing Physical States in Reaction becomes very important
VII) Calorimetry
1) Science of measuring heat absorbed or discharged by a body.
2) Calorimeter : device used to measure heat associated with a chemical reaction.
3) Heat Capacity (Table 6.4)
A)
Heat Capacity  C = heat absorbed / increase in Temperature
- Specific heat capacity = Energy required to raise temperature of
1 g of any substance by 1o C (or K).
[Units are J / g C or J / g K]
- Molar heat capacity = Energy required to raise temperature of
1 mole of any substance by 1o C (or K).
[Units are J / mol C or J / mol K ]
4) Constant Pressure Calorimetry
- “coffee cup calorimeter”
used to measure solid/liquid heat transfer
used to determine Hrxn for solutions.
- qmetal = qwater
qp = H
q = m c T
m = mass (g)
c = specific heat (J / g oC)
T = change in temperature (oC)
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5) Constant Volume Calorimetry
V = 0
w = - PV = 0
E = q + w
E = qv
VIII) Hess’s Law
1) Hess’s Law states that the overall enthalpy change for a reaction is equal to
the sum of the enthalpy changes for the individual steps in the reaction. H is
NOT dependent upon the pathway taken, since H is a state function.
2) Characteristics of H changes
A) If a reaction is reversed, the sign of H is also reversed.
B) The magnitude of H is directly proportional to the quantities of
reactants and products in a reaction. If the coefficients in a balanced
chemical reaction are multiplied by an integer, the value of H is
multiplied by the same integer.
3) Review Example Problems & Remember the 2 hints given above.
4) Standard Enthalpies of Formation Hf0
A) Change in enthalpy that accompanies the formation of 1 mole of a
compound from its elements with all substances in their standard states.
Hreaction0 =  nf Hf0 (products) -  nr Hr0 (reactants)
B) Standard states (Conventional Definitions)
Hf0 for elements in their standard states is defined as zero.
Hf0 for substances
(T = 25 C)
Gases: pressure is 1.0 atm
Solutions: concentration is 1.0 M
Pure liquids & solids: standard state is pure liquid or solid.
IX) Bond Dissociation Energies
1) Bond dissociation energies (BDEs) are just standard enthalpy changes (H0),
and can be used when one does not have sufficient H0 data for a process.
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2) H0 =  BDE (bonds broken) -  BDE (bonds formed)
bond breaking
bond formation
endothermic
exothermic
H +
H -
3) BDE values are always positive when taken from Tables
X) Energy Consumption & The Impact on Our Environment
1) Fossil Fuels
A) Fuels derived from decaying plant matter. Examples include coal,
petroleum,& natural gas.
2) Petroleum / Natural Gas
- transportation (cars, trucks, planes, etc.) which leads to air
pollutants like CO, CO2 , NO, CO2 .
- natural gas used to heat homes, cook food, etc.
3) Coal
- used to generate electricity for heating & cooling homes.
- plentiful (20 - 30 % of our energy needs satisfied by coal)
- expensive to mine & is highly dangerous
- harmful to our environment (photochemical smog & acid rain)
ACID RAIN..
S (coal) + O2 ----> SO2
2 SO2 + O2 -----> 2 SO3
SO3 + H2O -----> H2SO4
B) Greenhouse Effect
1. Result of increase in CO2 content in atmosphere over last 150 years
from combustion of fossil fuels. [ Industrial Revolution ]
2. Combustion reactions always produce carbon dioxide & water.
C) Alternate Energy Sources
1. Achieve balance between economic, climatic, & supply factors.
2. Possible candidates include
Solar (Sun)
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Nuclear (Fusion/Fission)
Biomass (Plants)
Synthetic Fuels (Syngas)
Hydrogen gas
XI) Spontaneity, Free Energy, & Entropy
1) A spontaneous process is one that proceeds on its own without any continuous
external influence.
2) A nonspontaneous process takes place only in the presence of an outside
influence.
3) Entropy (S) is a thermodynamic function that measures randomness or
disorder.
A) Second Law of Thermodynamics – entropy of universe is increasing.
B) Be familiar with predicting whether entropy increases or decreases for
a given reaction.
4) Free Energy (G) is a thermodynamic quantity that measures spontaneity for a
given process.
5) Mathematical Relationship between G, H, & S
G = H - TS
Gibbs Free Energy Equation
6) Predicting spontaneity using Gibbs Free Energy Equation
G is negative for a spontaneous process
G is positive for a nonspontaneous process
G is zero for a process at equilibrium