INTRO TO THERMOCHEMISTRY

Chemical reactions involve changes in energy
–
–

These energy changes can be in the form of
heat
–


Breaking bonds requires energy
Forming bonds releases energy
Heat is the flow of chemical energy
The study of the changes in energy in chemical
reactions is called thermochemistry.
The energy involved in chemistry is real and
generally a measurable value
WHAT IS HEAT?
Hot & cold, are automatically associated
with the words heat and temperature
– Heat & temp are NOT synonyms
– The temperature of a substance is
directly related to the energy of its
particles, specifically its:
The Kinetic Energy defines the
temperature
– Particles vibrating fast = hot
– Particles vibrating slow = cold

Vibrational energy is transferred from one
particle to the next
–
One particle collides with the next particle
and so on; and so on – down the line
An Ice Cold Spoon
A Hot Spoon

Thermal energy is the total energy of all the particles that
make up a substance
Kinetic energy from vibration of particles
– Potential energy from molecular attraction (within or
between the particles)
–

Thermal energy is dependent upon the amount or mass of
material present
(KE =½mv2)
2 Hot Spoons

Thermal energy is also related to
the type of material

Different type of materials
May have the same temp, same mass, but different
connectivity
– Affected by the potential energy or the intermolecular forces
–
So it is possible to be at same temp (same KE) but have very
different thermal
energies
 The different abilities to hold
onto or release
energy is
referred to as the
substance’s heat capacity


Thermal energy can be transferred from object to object
through direct contact
–
Molecules collide, transferring energy from molecule to
molecule
DEFINITION
THE FLOW OF THERMAL ENERGY FROM SOMETHING
WITH A HIGHER TEMP TO SOMETHING WITH A LOWER
TEMP
UNITS
MEASURED IN JOULES OR CALORIES
THROUGH WATER OR AIR = CONVECTION
TYPES
THROUGH SOLIDS = CONDUCTION
TRANSFERRED ENERGY BY COLLISION WITH PHOTON
= RADIANT ENERGY
HEAT CAPACITY
 The measure of how well a material
absorbs or releases heat energy is its
heat capacity
– It
can be thought of as a reservoir to hold
heat, how much it holds before it overflows
is its capacity

Heat capacity is a physical property
unique to a particular material
– Water
takes 1 calorie of energy
to raise temp 1 °C
– Steel takes only 0.1 calorie of
energy to raise temp 1 °C
SPECIFIC HEAT CAPACITY
 The
amount of energy it takes to raise the
temp of a standard amount of an object
1°C is that object’s specific heat capacity
(Cp)
–
The standard amount =1 gram
 Specific
heats can be listed on data tables
Smaller the specific heat  the less
energy it takes the substance to feel hot
– Larger the specific heat  the more
energy it takes to heat a substance up
(bigger the heat reservoir)
–
SUBSTANCE
SPECIFIC HEAT CAPACITY, CP
WATER, H2O
4.18J/g°C OR 1cal/g°C
ALUMINUM, Al
.992J/g°C OR .237cal/g°C
TABLE SALT, NaCl
.865 J/g°C OR .207cal/g°C
SILVER, Ag
.235 J/g°C OR .056cal/g°C
MERCURY, Hg
.139 J/g°C OR .033cal/g°C
SPECIFIC HEAT CAPACITY
 Specific heats and heat capacities work for gains in heat and
in losses in heat
Smaller the specific heat  the less time it takes the
substance to cool off
– Larger the specific heat  the longer time it takes the
substance to cool off
–
 Specific heat capacity values are used to calculate changes in
energy for chemical reactions
–
It’s important for chemists to know how much energy is
needed or produced in chemical reactions
CHEMICAL REACTIONS
There are 2 types of chemical reactions
– Exothermic reactions  reactions in which heat energy is a
product
– Endothermic reactions  reactions in which heat energy is a
reactant
 Exothermic reactions typically feel warm as the reaction proceeds
– Gives off heat energy, sometimes quite allot
 Endothermic reactions typically feel cooler the longer the reaction
proceeds
– Absorbs heat energy, sometimes enough to get very cold


Exothermic reaction
C3H8 + 5O2  3CO2 + 4H2O + 2043kJ
– To a cold camper, the important product here is the heat
energy

Endothermic reaction
NH4NO3+H2O+ 752kJ NH4OH+HNO3
– Similar system as what is found in cold packs
CHANGE IN HEAT ENERGY (ENTHALPY)
The energy used or produced in a chemical reaction is called the
enthalpy of the reaction
– Burning a 15 gram piece of paper produces a particular amount
of heat energy or a particular amount of enthalpy
 Enthalpy is a value that also contains a component of direction
(energy in or energy out)
– Heat gained is the out-of
direction; ie exo
CHANGE IN HEAT ENERGY (ENTHALPY)
The energy used or produced in a chem rxn is called the enthalpy of
the reaction
– Burning a 15 gram piece of paper produces a particular amount
of heat energy or a particular amount of enthalpy
 Enthalpy is a value that also contains a component of direction
(energy in or energy out)
– Heat gained is the out-of
direction; ie exo
–
Heat lost is the into
direction; ie endo-
HEAT
HEAT
HEAT
HEAT
CHANGE IN ENTHALPY
 Most common version of enthalpy is when we have a change
in enthalpy (H)
 The enthalpy absorbed or gained (changed) in a reaction is
dependent on the amount of material reacting
Amount is usually in the form of moles
– We can use the coefficient ratios to energy ratios to
calculate how much energy a reaction used or produced
–
USING H IN CALCULATIONS
 Chemical reaction equations are very powerful tools.
–
Given a reaction equation with an energy value, We can
calculate the amount of energy produced or used for any
given amount of reactants.
(For Example)
How much heat will be released if 1.0g of (H2O2)
decomposes in a bombardier beetle to produce a
defensive spray of steam
2H2O2 2H2O + O2 Hº = -190kJ
2H2O2 2H2O + O2 Hº = -190kJ
Analyze: we know that if we had 2 mols of H2O2 decomposing we
would produce 190kJ of heat, but how much would it
be if only 1.0 g of H2O2
Therefore: we have to convert our given
moles of H2O2
1.0g H2O2
1molH2O2
34gH2O2
Molar mass
1.0 g of H2O2 to
= .02941 mol
2H2O2 2H2O + O2 Hº = -190kJ
Therefore: with 2 moles of H2O2 we would produce 190 kJ of
energy, but we don’t have 2 moles we only have
.02941 mols of H2O2, so how much energy would
the bug produce?
-190kJ
.02941 mol
= -2.8kJ
2molH2O2
Reaction equation
Example #2
How much heat will be released when 4.77 g of ethanol (C2H5OH)
react with excess O2 according to the following equation:
C2H5OH + 3O2 2CO2 + 3H2O Hº=-1366.7kJ
analyze: we know that if we had 1 mol of ethanol(assuming
coefficient of 1 in rxn equation) burning we would
produce 1366.7kJ of heat, but how much would it be if
only we only had 4.77 g of ethanol?
C2H5OH + 3O2 2CO2 + 3H2O Hº=-1366.7kJ
4.77g C2H5OH
1mol C2H5OH
46g C2H5OH
= .1037 mol
Therefore: with 1 mole of C2H5OH we would produce 1366.7 kJ of
energy, but we don’t have 1 mole we only have .1037
mols of C2H5OH, so how much energy would the
reaction produce?
.1037 mol
-1366.7kJ
1mol C2H5OH
= -142 kJ

We can also track energy changes due to temp changes, using
H=mCT:
H =
MASS
SPECIFIC
HEAT
FINAL TEMP –
INITIAL TEMP

If the temp difference is positive
– The reaction is exothermic because the final temp is greater than
the initial temp
– So the enthalpy is positive

if the temp change is negative
– makes the enthalpy negative
– the reaction absorbed heat into the system, so it’s endothermic
If you drink 4 glasses of ice water at 0°C, how much heat energy is
transferred as this water is brought to body temperature?
Each glass contains 250 g of water & body temperature is 37°C.
 mass of 4 glasses of water:
– m = 4 x 250g = 1000g H2O
 change in water temp:
– Tf – Ti = 37°C - 0°C
 specific heat of water:
– C = 4.18 J/g•C° (from previous slide)
H=mCT
H=(1000g)(4.18J/g•°C)(37°C)
H= 160,000J
Enthalpy is dependent on the conditions of the reaction
– It’s important to have a standard set of conditions
– This allow us to compare the affect of temperatures, pressures,
etc., has on different substances
 Chemists have defined a standard set of conditions
– Standard Temperature = 298K or 25°C
– Standard Pressure = 1atm or 760mmHg
 Enthalpy produced in a reaction under standard conditions is the
standard enthalpy (H°)

Standard enthalpies can be found on tables of data measured as
standard enthalpies of formations
 Standard enthalpies of formations are measured values for the
energy to form chemical compounds (Hf°)
– H2 gas & O2 gas can be ignited to produce H2O and a bunch of
energy
– The amount of energy produced by the reaction is 285kJ for every
mol of water produced

H2 + ½02  H2O
Hf°=-285.8kJ/mol
STANDARD ENTHALPIES OF FORMATION
SYMBOL
FORMULAS
Hf°kJ/mol
AlCl3(s)
Al + 3/2Cl2  AlCl3
-705.6
Al2O3(s)
2Al + 3/2O2  Al2O3
-1676.0
CO2(g)
C + O2  CO2
-393.5
H2O(g)
H2 + 1/2O2  H2O
-241.8
C3H8(g)
3C + 4H2  C3H8
-104.7
CALORIMETRY
Calorimetry is the process of measuring heat energy
– Measured using a device called a calorimeter
– Uses the heat absorbed by H2O to measure the heat given off by a
reaction or an object
 The amount of heat soaked up by the water is equal to the amount of
heat released by the reaction

Hsys is the system or what is taking place in the main chamber
(reaction etc.) And Hsur is the surroundings which is generally
water.
HSYS=-HSUR
A COFFEE CUP
CALORIMETER
USED FOR A REACTION
IN WATER,
OR JUST A TRANSFER
OF HEAT.
A BOMB
CALORIMETER
USED WHEN TRYING
TO FIND THE AMOUNT
OF HEAT PRODUCED BY
BURNING SOMETHING.

CALORIMETRY
With calorimetry we use the sign of what happens to the
water
– When the water loses heat into the system it obtains a (-)
sign
HSYS
+ SIGN MEANS
HEAT WAS
ABSORBED BY
THE REACTION
=
-HSUR
- SIGN MEANS
HEAT WAS
RELEASED BY
WATER
HEAT
HEAT
HEAT
HEAT
CALORIMETRY
You calculate the amount of heat absorbed by the water (using H=
mCT)
 Which leads to the amount of heat given off by the reaction
– you know the mass of the water (by weighing it)
– you know the specific heat for water (found on a table)
– and you can measure the change in the temp of water (using a
thermometer)

A block of Al that weighs 72.0g is heated to 100°C is dropped in a
calorimeter containing 120. mL of water at 16.6°C.
the H2O’s temp rises to 27°C.
-
mass of Al = 72g
Tinitial of Al = 100°C
Tfinal of Al = 27°C
CAl = .992J/g°C (from table)
H =
72g
.992J/g°C
H
=
-5214J
HSYS
27°C-100°C

We can do the same calc with the water information:
– Mass of H2O= 120g
– Tinitial of H2O= 16.6°C
HSUR
– Tfinal of H2O = 27°C
– CH2O= 4.18J/g°C (from table)
H = 120g
4.18J/g°C
H = 5216J
27°C-16.6°C
Equal but opposite, means that since the Al decreased in
temperature, it released heat causing the H2O to increase in
temperature.
Thermochemistry
System - the part of the universe that
is being studied
Surroundings - the rest of the universe
Boundary - the separation between the
the system and surroundings
Systems may be: (1) open; (2) closed; or
(3) isolated (adiabatic)
Thermochemistry
Laws of Thermodynamics
Zero Law - Two objects in contact will have the
same temperature
1st Law - The energy of the universe is constant
2nd Law - The entropy of the universe is expanding
3rd Law - The entropy of a perfect crystalline substance
is 0 at absolute zero (0 K)
The First Law of Thermo dynamics
The total energy and mass of a system
plus its surroundings remains constant.
∆E = q + w
Energy is measured in units of:
Calorie (cal) or Joule* (j) where
1 calorie = 4.184 joules
*The SI unit is the Joule
Heat of reaction is measured in a calorimeter.
A bomb calorimeter
constant volume
E (internal
energy)
∆E
An ordinary calorimeter
constant pressure
H (enthalpy)
∆H
The enthalpy change of a reaction (∆H):
the heat absorbed or released
during a reaction at constant
pressure.
A reaction that releases heat is exothermic,
∆H is negative.
A reaction that absorbs heat is endothermic,
∆H is positive.
The first law does not
predict the direction
of a process!
The Second Law
of Thermodynamics
The Second Law of Thermodynamics
Chemical and physical changes will take
place in a direction to produce maximum
disorder in the system + surroundings.
Order
Disorder or
randomness
Entropy (S) is a quantitative measure
of disorder.
The more positive is the value of S, the
greater is the disorder.
(∆Ssystem + ∆Ssurroundings) > 0 for a spontaneous process
(∆Ssystem + ∆Ssurroundings) > 0 for a spontaneous process
Hard to measure ∆Ssystem + surroundings
Free Energy (G)
∆G° = ∆H° - T∆S°
∆G°: the change in free energy of a system
∆H°: the change in enthalpy of the system
∆S°: the change in entropy of the system
T: absolute temperature (in Kelvin units, K, oC+273)
For a spontaneous process, ∆G < 0.
More free energy
(great work capacity)
Spontaneous change
∆G < 0
Less free energy
(less work capacity)
battery
AB
∆G = GB - GA
(1) ∆G < 0: exothermic (exergonic) (there is a net loss of
energy).
The “A to B” reaction is favorable and spontaneous.
(2) ∆G > 0: endothermic (endergonic) (there is a net gain
of energy)
The “A to B” reaction is unfavorable and not spontaneous.
The battery is dead!
(3) ∆G = 0: the reaction is at equilibrium.
∆G tells us how far the system is from equilibrium
AB
∆G = GB-GA
A
B
∆G << 0
∆G < 0
∆G = 0
∆G > 0
The relationship of ∆G to Keq and
the concentration of reactants and products
A+BC+D
R is the gas constant (1.987 cal/mol.degree).
T is the absolute temperature (oK)
P=1M, R=1M (standard condition)
∆Go = - RT ln Keq = - 2.303 RT log Keq
∆Go is the standard free energy change
At the standard temperature of 298oK (25oC):
∆Go = -1.36 log Keq
(1) Keq >1, then ∆Go < 0
(2) Keq< 1, then ∆Go > 0
(3) Keq= 1, then ∆Go = 0
At pH 7, [H+] is 10-7M, the symbol ∆Go´ is used.
∆Go´ = -1.36 log K´eq
K´eq
10-5
10-4
10-3
10-2
10-1
1
10
102
103
104
∆Go´
-------Kcal/mol
6.82
5.46
4.09
2.73
1.36
0
-1.36
-2.73
-4.09
-5.46
kJ/mol
28.53
22.84
17.11
11.42
5.69
0
-5.69
-11.42
-17.11
-22.84
∆G = ∆Go + RT ln
P
R
1. ∆G is predictive .
2. ∆G depends on Keq, and the concentrations of P and R.
Sometimes, ∆Go > 0, but if P/R is very small, then ∆G < 0.
Summary:
RP
∆H, ∆S and ∆G
∆G = ∆H - T∆S
∆G < 0, R
P
RP
∆G0
∆G0 = -RTlnKeq
∆G =
∆Go
Keq>1, ∆G0<0
+ RT ln
∆G < 0, R
P
P
R