Chapter 6
Energy and Chemical Reactions
James P. Joule 1818-1889. Discovered
mechanical equivalent of heat, which led
to the First Law of Thermodynamics.
Germain Henri Hess
1802-1850.
Hess’s Law.
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The Nature of Energy
Kinetic and Potential Energy
Kinetic energy is the energy of motion.
Potential energy is energy which is available,
such as an object that can fall a distance,
potentially releasing kinetic energy when it
hits the ground.
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The Nature of Energy
Energy Units
SI Unit for energy is the joule, J:
For example, energy of motion (kinetic energy):
E k  mv
1
2
2
(in J)
We sometimes use the calorie instead of the joule:
1 cal = 4.184 J (exactly)
A nutritional Calorie:
1 Cal = 1000 cal = 1 kcal
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The Nature of Energy
All objects and substances have an internal energy, E
(this is a form of potential energy, and in this course is
commonly called “chemical energy”).
In chemistry, we’re interested in the internal energies of
reactants and products and how it changes (ΔE) in going
from:
reactants
products
We usually can only measure ΔE, not E
“Δ” means change in
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First Law of Thermodynamics
Total energy is conserved
Energy cannot be created or destroyed.
In chemistry, we are concerned with heat and
chemical energy.
The first law of thermodynamics tells us that the
sum of heat and chemical energy must be
constant (at constant pressure).
Heat at constant pressure is called enthalpy.
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Enthalpy
Enthalpy, H: Heat transferred between the system and
surroundings carried out under constant pressure.
Can only measure the change in enthalpy:
DH = Hfinal - Hinitial = qP
We will use enthalpy (H) and energy (E) interchangeably
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Endothermic and Exothermic Processes
Endothermic: system absorbs heat from surroundings.
(ΔH = +)
Exothermic: system transfers heat to the surroundings.
(ΔH = -)
An endothermic reaction feels cold.
An exothermic reaction feels hot.
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Enthalpies of Reaction
For a reaction
DHrxn = H(products) - H (reactants)
Enthalpy is an extensive property (magnitude DH is
directly proportional to amount):
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
DH = -802 kJ
2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(g) DH = -1604 kJ
When we reverse a reaction, we change the sign of DH:
CO2(g) + 2H2O(g)  CH4(g) + 2O2(g)
DH = +802 kJ
Change in enthalpy depends on state (gas, liquid, solid):
H2O(g)  H2O(l) DH = -44 kJ
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Calorimetry
Heat Capacity and Specific Heat
Calorimetry = measurement of heat flow.
Calorimeter = apparatus that measures heat flow.
Heat capacity = the amount of energy required to raise
the temperature of a substance (by 1o).
Molar heat capacity = heat capacity of 1 mol of a
substance (energy required to raise temp. of one mole
of a substance (by 1o)
Specific heat = specific heat capacity = heat capacity of
1 g (energy required to raise temp of one gram by 1o.
Heat absorbed by system, q, is given by:
q = (specific heat)  (grams of substance)  T
(Be careful of the sign of q)
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Calorimetry
Some specific heats (in J/g-K)…
N2(g):
1.04
Al(s):
0.90
Fe(s):
0.45
CH4(g):
2.20
NOTE the very high specific
heat for liquid water.
This has a moderating influence
on the weather and the
temperature of living organisms.
CaCO3(s): 0.82
H2O(l):
4.18
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Calorimetry
Constant-Pressure
Calorimetry
Atmospheric pressure is
constant!
DH = qP
Heat emitted by reaction (qrxn)
is absorbed by solution (qsoln)
So,
qrxn = -qsoln
qrxn =
-qsoln = -(specific heat of soln)  (grams of solution)  DT.
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Calorimetry
9.55g sample of solid NaOH dissolves in 100.0 g water in
coffee-cup calorimeter. Temp rises from 23.6 to 47.4oC.
Calculate ΔH (kJ/mol) for the soln process:
NaOH(s)
Na+ (aq) + OH- (aq)
Assume specific heat of soln is same as that of pure water
(which is 4.18 J/g-K).
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Calorimetry
Bomb Calorimetry (Constant-Volume Calorimetry)
Used to study
combustion processes
qrxn = -Ccal ΔT
where Ccal is the heat
capacity of the
calorimeter (to be
determined
experimentally)
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Hess’ Law
Hess’ law: if a reaction is carried out in a number of
steps, DH for the overall reaction is the sum of DH for
each individual step.
Note that:
DH1 = DH2 + DH3
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Hess’ Law
• A consequence of Hess’ Law is that reactions
can be added just like algebraic equations.
• For example:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
2H2O(g)  2H2O(l)
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
DH = -802 kJ
DH = -88 kJ
DH = -890 kJ
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Enthalpies of Formation
• If 1 mol of compound is formed from its constituent
elements, then the enthalpy change for the reaction is
called the enthalpy of formation, DHof .
This is also called the heat of formation.
Standard conditions (standard state): 1 atm
and 25 oC (298 K).
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Enthalpies of Formation
• Standard enthalpy of formation of the most stable form of an
element is zero. For example, it is zero for C(s), H2(g), Pb(s),
O2(g), etc
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Enthalpies of Formation
Using Enthalpies of Formation to Calculate
Enthalpies of Reaction.
We use Hess’ Law to calculate enthalpies of a reaction
from enthalpies of formation
Example:
Combustion of 1 mol of
propane gas, C3H8 (g)
Can construct alternate
path of three steps shown
as 1, 2 and 3
DHrxn =DH1 + DH2 +DH3
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Enthalpies of Formation
Using Enthalpies of Formation to Calculate
Enthalpies of Reaction
For a reaction:
DH rxn   nDH  f products    mDH  f reactants 
(a) 2 SO2(g) + O2(g)
2 SO3 (g) -- what is ΔH?
(c) 4 FeO (s) + O2 (g)
2 Fe2O3 (s) -- what is ΔH?
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Foods and Fuels
Foods
1 Nutritional Calorie = kcal
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Foods and Fuels
Fuels
U.S.: 1.0 x 106 kJ of fuel per day.
Most from petroleum and natural gas.
Remainder from coal, nuclear, and hydroelectric.
Fossil fuels are not renewable.
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