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Chemistry, The Central Science, 11th edition
Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten
Chapter 17
Thermochemistry
Edited by Kavita Gupta
© 2009, Prentice-Hall, Inc.
Energy
• Energy is the ability to do work or transfer
heat.
– Energy used to cause an object that has mass to
move is called work.
– Energy used to cause the temperature of an object
to rise is called heat.
– There are different forms of energies- radiant,
thermal, chemical, potential etc. We will focus on
thermal (heat) energy in this unit.
• Units for energy:
• The SI unit of energy is the joule. Another
common unit of energy is calorie.
•
1 calorie = 4.184 joules
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Heat
• Energy can also be
transferred as heat.
• Heat flows from
warmer objects to
cooler objects.
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Temperature
TEMPERATURE: a measure of average kinetic
energy of the particles of matter (remember
that molecules of ALL gases at the same temp.
have the same average kinetic energy)
THERMOCHEMISTRY: Branch of Chemistry that
deals with the study of heat change in chemical
reactions.
Definitions:
System and Surroundings
• The system includes the
molecules we want to study
(here, the hydrogen and
oxygen molecules).
• The surroundings are
everything else (here, the
cylinder and piston).
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System
SYSTEM: that part of the universe on which attention is
focused. Usually reactants and products.
There are three types of systems:
• Open System: can exchange both mass and energy.
• Closed System: can exchange only energy.
• Isolated System: can neither exchange mass nor
energy.
• What kind of system is a heat/cold pack?
Surroundings
• SURROUNDINGS: what exchanges energy
with the system, the vessel in which the
reaction takes place and the air or other
material in contact with the reaction system.
• Label system and surrounding in the
following diagram.
Exchange of Heat between System
and Surroundings
• When heat is absorbed by the system from the
surroundings, the process is endothermic.
(usually the temp of the system increases and it
feels cold to touch!)
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Exchange of Heat between System
and Surroundings
• When heat is absorbed by the system from the
surroundings, the process is endothermic.
• When heat is released by the system into the
surroundings, the process is exothermic. (Usually
the temp. of the system lowers and it feels warm
to touch)
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State Functions
Usually we have no way of knowing the internal
energy of a system; finding that value is simply
too complex a problem.
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Enthalpy
•
•
•
•
•
•
Enthalpy (H) is defined as the heat content of a system at
constant pressure.
It is an extensive property.
It is not possible to measure the heat content of the system
only the changes in the heat content of the system can be
measured. Enthalpy change is defined as the amount of heat
absorbed or lost by a system during a process at constant
pressure.
Heat of reaction and change in enthalpy are used
interchangeably.
Units: kJ/mol
Qreaction at constant pressure= H(change in enthalpy) = H
(products) – H (reactants)
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Endothermicity and Exothermicity
• A process is
endothermic when
H is positive.
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Endothermicity and Exothermicity
• A process is
endothermic when
H is positive.
• A process is
exothermic when
H is negative.
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Energy Diagram for Endo- and
Exothermic Reactions
• Endothermic (dH +) • Exothermic ( dH -)
Enthalpy of Reaction
The change in enthalpy,
H, is the enthalpy of
the products minus the
enthalpy of the
reactants:
H = Hproducts − Hreactants
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Enthalpy of Reaction
This quantity, H, is called the enthalpy of
reaction, or the heat of reaction.
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The Truth about Enthalpy
1.
2.
3.
Enthalpy is an extensive property. hence it is
proportional to the amount of reactants and
products. e.g. for decomposition of two moles of
water twice as much energy is needed as for one
mole of water.
H for a reaction in the forward direction is equal in
size, but opposite in sign, to H for the reverse
reaction. Reversing a reaction changes the sign of
enthalpy. AB 10kJ, then BA –10kJ
H for a reaction depends on the state of the
products and the state of the reactants. (different
values of dH for H2O (l) and H2O (g).
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Thermochemical Equation
• Thermochemical equation: A chemical
equation that shows the enthalpy relation
between products and reactants is called a
thermochemical equation.
• Ex. H2 (g) + Cl2 (g)2 HCl (g)
H=-185 kJ
Heat Capacity and Specific Heat
The amount of energy required to raise the
temperature of a substance by 1 K (1C) is its heat
capacity.
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Heat Capacity and Specific Heat
We define specific heat capacity (or simply
specific heat) as the amount of energy required to
raise the temperature of 1 g of a substance by 1 K.
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Heat Capacity and Specific Heat
Formula for Specific Heat
q(amount of heat)= m (g of substance) x c(specific heat) x T (
temperature change)
Heat Capacity (C)
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Calorimetry
Since we cannot know
the exact enthalpy of
the reactants and
products, we measure
H through
calorimetry, the
measurement of heat
flow.
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Calorimeter
• Calorimeter is used to measure heat flow in a
reaction.. Only heat flow is between the reaction
system and the calorimeter, the walls of the
calorimeter is insulated, so no heat exchange
occurs between calorimeter and surroundings.
• q reaction = - q calormeter
• if reaction is exothermic, heat flows into
calorimeter, if reaction is endothermic heat flows
into reaction mixture..
Types of Calorimeters
• There are two types of calorimeters:
• Coffee-Cup Calorimeter (Constant Pressure
Calorimetry): It consists of foam cup filled with
water. Cup has a tight fitting lid through which a
thermometer is inserted. Because Styrofoam is a
good insulator, very little heat is lost to the
surroundings. This means that the heat capacity
of calorimeter is that of the water, which can be
calculated by
Qrxn=-mwater x 4.18 j/g-0c x t
Coffee cup calorimeter can be used to measure
heat of reaction of solutions only
Constant Pressure Calorimetry
By carrying out a reaction
in aqueous solution in a
simple calorimeter such as
this one, one can indirectly
measure the heat change
for the system by
measuring the heat
change for the water in
the calorimeter.
© 2009, Prentice-Hall, Inc.
Constant Pressure Calorimetry
Because the specific heat
for water is well known
(4.184 J/g-K), we can
measure H for the
reaction with this
equation:
q = m  s  T
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Bomb Calorimeter
• Bomb Calorimeter: Bomb calorimeter is used
to measure heat of reactions for gases and
also for reactions that reach high
temperatures. To use it, a weighed sample of a
reactant is put in a heavy metal vessel called
as a “bomb”. Initial temperature is recorded
and reaction is started by an ignition. Final
temperature is recorded.
• Qrxn= -q calorimeter
Bomb Calorimetry
• Reactions can be
carried out in a sealed
“bomb” such as this
one.
• The heat absorbed (or
released) by the water
is a very good
approximation of the
enthalpy change for the
reaction.
© 2009, Prentice-Hall, Inc.
Bomb Calorimetry
• Because the volume in
the bomb calorimeter is
constant, what is
measured is really the
change in internal
energy, E, not H.
• For most reactions, the
difference is very small.
© 2009, Prentice-Hall, Inc.
Heat Transfer (Heating and Cooling
Curves)
• Heat Transfer :Look at the following heating/cooling curve
for water (Remember this from Chap 12)
• Label on this graph for: temperatures, fusion, freezing,
vaporization, condensation
• http://www.wwnorton.com/college/chemistry/gilbert/tuto
rials/interface.asp?chapter=chapter_11&folder=heating_cu
rves (Heating/Cooling Curve Animation)
For Calculating Energy at Phase
Change
• During phase changes all the heat is used for
breaking the bonds. To calculate amount of heat
absorbed/released during phase changes (q), the
following formula is used: Q = m. H ( or n. H)
• Value of H changes with phase change. H of
water has following values:
• Hfusion = 334 J/g (6.02 kJ/mol), Hfreezing = - 334 J/g
(-6.02 kJ/mol),
• Hvaporization = 2260 J/g (40.7 kJ/mol), Hcondensation =
-2260 J/g (-40.7 kJ/mol)
For Calculating Energy other than at
phase change
• Use the formula Q=m x c x t
• Total energy change can be calculated by
adding all the energy changes.
Ex. How much heat is needed to vaporize 5.0 g
of ice at -5 degree C to steam at 110 degree C?
• Enthalpy of Formation(Hf), of a substance is the
enthalpy change for the reaction in which the
substance is formed from its constituent elements.
Write an equation for heat of formation of Carbon
dioxide gas.
Write a thermochemical equation for the formation of
AgCl, if Hf0 for AgCl is –127.1 kJ/mol.
• Standard enthalpy of formation, Hf0,of a substance is
the change in enthalpy for the reaction that forms 1
mole of a substance from its elements with all the
reactants and products1 atm pressure and 298 K.
• For a pure element in its most stable
state,Hf0=0
• Compounds with high negative enthalpy of
formation are very stable while the ones with
high positive enthalpy of formation are
unstable. Why?
• The importance of standard enthalpies of
formation is that once we know their values,
we can calculate standard enthalpy of
reaction, H0. How?
Hess’s Law
• H is well known for many reactions, and
it is inconvenient to measure H for
every reaction in which we are
interested.
• However, we can estimate H using
published H values and the properties
of enthalpy.
© 2009, Prentice-Hall, Inc.
Hess’s Law
Hess’s law states that
“[i]f a reaction is carried
out in a series of steps,
H for the overall
reaction will be equal to
the sum of the enthalpy
changes for the
individual steps.”
© 2009, Prentice-Hall, Inc.
Hess’s Law
Because H is a state
function, the total
enthalpy change depends
only on the initial state of
the reactants and the final
state of the products.
© 2009, Prentice-Hall, Inc.
Enthalpies of Formation
An enthalpy of formation, Hf, is defined as
the enthalpy change for the reaction in which
a compound is made from its constituent
elements in their elemental forms.
© 2009, Prentice-Hall, Inc.
Standard Enthalpies of Formation
Standard enthalpies of formation, Hf°, are
measured under standard conditions (25 °C and
1.00 atm pressure).
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Calculation of H
We can use Hess’s law in this way:
H =  nHf°products –  mHf° reactants
where n and m are the stoichiometric
coefficients.
© 2009, Prentice-Hall, Inc.
Calculation of H
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l)
H = [3(-393.5 kJ) + 4(-285.8 kJ)] – [1(-103.85 kJ) + 5(0 kJ)]
= [(-1180.5 kJ) + (-1143.2 kJ)] – [(-103.85 kJ) + (0 kJ)]
= (-2323.7 kJ) – (-103.85 kJ) = -2219.9 kJ
© 2009, Prentice-Hall, Inc.
Hess’s Law Contd.
• Many compounds can not be synthesized directly from their elements. In
such cases, Hess’s law is used to calculate H of the reaction.
• Hess’s Law:
• According to Hess’s law, if a reaction is carried out in a series of steps, H
for the reaction will be equal to the sum of the enthalpy changes for the
steps.
Example:
Use Hess's law to calculate H for the reaction:
MgO(s) + CO2(g) MgCO3(s)
from the following enthalpy of formation data.
Mg(s) + 1/2 O2(g)MgO(s)
Hf0 = -601.70 kJ/mol MgO
C(s) + O2(g)CO2(g)
Hf0= -393.51 kJ/mol CO2
Mg(s) + C(s) + 3/2 O2(g)MgCO3(s) Hf0 = -1095.8 kJ/mol MgCO3
• Do practice problems1,2,3 on page 524
Enthalpy
• Enthalpy of Combustion is the heat released
by the complete combustion of one mole of a
substance. Hc
•
• Example: Write a thermochemical equation
for combustion of C6H6, if Hc for C6H6 is –
3169.0 kJ/mol.
•
Enthalpy of Vaporization
• Enthalpy of Vaporization: Amount of heat
required to boil (or condense) 1 mol of a
substance at its boiling point. Hvap for
water= 40.7 kJ/mol, what would Hcondensation
be for water?
Example:
• How much heat is required to vaporize 5
moles of water at its B.P.?
Enthalpy of Fusion
• Enthalpy of Fusion: Amount of heat required
to melt (or freeze) 1.00 mol of a substance at
its freezing point. Hfus for water=6.02kJ/mol
Spontaneity of A Chemical Reaction:
• Spontaneity of a reaction refers to if the reaction occurs
spontaneously. Spontaneous reactions may not always be
fast. Some examples of spontaneous reactions are rusting,
reaction of baking soda with vinegar etc.
• Non-Spontaneous Reactions: need something to get them
started like heat, pressure or catalyst etc.
• There are two factors that affect the spontaneity of a
reaction:
• Enthalpy (H), exothermic reactions (-H) are usually
spontaneous while endothermic reactions (+H) are
usually non spontaneous. However, some endothermic
reactions are spontaneous such as melting.
• Entropy: Reaction is spontaneous, if value of S is positive
and non spontaneous, if value of S is negative.
Entropy (S)
• It is a measure of randomness or disorder of the particles in a
system. Units: kJ/mol-K
• General rules for predicting entropy changes.
1. Look at the states first. (gases > liquids > solids)
2. If both states are the same then look at the number of moles of
reactants and products and decide if there has been an increase in
the number of moles or a decrease.
3. If the volume of a container of gas is increased, it results in
increased entropy (more space for random movement)
4. If the pressure on a container of gas is increased, then the entropy
decreases (due to restricted movement of particles)
• Arrange ice, water and water vapor in order of decreasing entropy.
Gibbs Free Energy
• Gibbs Free Energy(or Free Energy) G, is a
thermodynamic quantity that combines the values of
enthalpy and entropy. For a process that occurs at
constant temperature,
• G=H-TS
• The sign of G relates to the spontaneity of the
reaction. When it is negative process is spontaneous.
Relating Enthalpy, Entropy, Gibbs free energy to the spontaneity of the reactions
H
S
G
Negative value(exothermic)
Positive value(disorder)
Always negative
( spontaneous)
Negative value
Negative value
Negative at lower
temperatures
Positive value
Positive value
Negative at higher
temperatures
Positive value
Negative value
Never negative (non
spontaneous)
Kinetics
• Kinetics
• What is Chemical Kinetics?
•
•
Chemical kinetics is the study of rates of reaction.
Rate: measure of the speed of any change that occurs within an
interval of time.
• EX: A sprinter covers 100 meters in 11.5 seconds. Rate is 8.7 m/s.
• 1. Reaction X  O / min, then, 1 min. XXXXXO 2 min. XXXXOO 3
min. XXXOOO 4 min. XXOOOO 5 min. XOOOO
Reaction Rates in Chemistry
• Reaction Rates in Chemistry are usually described in terms of change in
concentration per unit time. It can be written either as appearance of
product per unit time or disappearance of reactants per unit time. Units:
M/s. Reaction rates can be determined experimentally by changing
concentration of reactants. For example, rate of reaction for the following
reaction can be written in terms of disappearance of HI or appearance of
H2 or I2.
• 2 HI (g)  H2(g) + I2 (g)
• Also, note that the rate of change of HI is twice as compared to formation
of H2 or I2. So, we can equate these rates as follows:
• Units for rate: change in concentration/ change in time (Molarity/second
or Ms-1)
Rate Laws
• Rate Laws of Reactions:
• Rate law is an equation that shows quantitative
relationship between the reaction rate and
concentrations of reactants. This relationship can only
be measured experimentally.
• r [A], r k[A], where k is specific rate constant. The
value of k changes only with temperature.
• To calculate rate law, the concentration of reactants are
changed and changes in rate is observed
experimentally.
To calculate rate law, the concentration of reactants are changed and changes in
rate is observed experimentally.
Experiment #
[A]
[B]
Rate (M/s)
.100
.100
4.0 X 10^-5
1
2
.100
.200
8.0 X 10^-5
3
.200
.100
1.6
• Ex: 2H2(g) + 2NO(g)  N2(g) + 2 H2O(g)
• In the above example, 4 moles of reactant
produces 3 moles of product gases, pressure
goes down as reaction proceeds. Rate
determined by measuring change in pressure
over time.
• In the above example, 4 moles of reactant produces 3 moles of
product gases, pressure goes down as reaction proceeds. Rate
determined by measuring change in pressure over time.
• 1. If start with same initial conc of NO but different H2 . Initial
reaction rate is found to vary directly with the hydrogen conc. If
double H conc doubles rate, if triple H2 triples rate, IF r= react rate
and (H2) = conc of H2 in moles/L. then r is proportional, , to (H2).
R (H2)x 2 (2)x r=1
• if use same concentration of H2 but vary NO . Initial reaction rate is
found to increase 4X when NO conc is doubled and nine fold when
conc is tripled.. Reaction rate varies directly with the square of the
nitrogen monoxide conc. R (NO)2. R (NO)x 4 (2)x r=2
• r (H2)(NO) 2
(To remove the proportionality sign, a constant
called as specific rate constant (k) is
introduced.)
k is proportionality constant. So r= k(H2)(NO) 2
k is specific for a given reaction at a given temp,
k usually increases with increase in temp.
Order of Reaction
• Order of Reaction:
• The power to which a reactant conc is raised is called the
ORDER in that reactant..
• An order of 1, rate is directly proportional.
•
2, rate is directly proportional to the square of the
reactant
•
0, rate does not depend on the concentration of
the reactant as long as some of the reactant is present.
• Overall order of reaction: sum of order of reactants.
• Order with respect to H2 is first order, with respect to NO is
second order and to overall reaction is 3rd.
Units of Rate Constants
• In zero order reaction,the rate of reaction is
independent of the concentration of reactants.
• Rate= k
units of k= M/s
• In first order reaction, the rate is proportional to
the concentration of only one of the reactants
raised to the first power.
• Rate= k[A]
units of k= s-1
• In second order reaction, the overall order of
reaction is 2,
• Rate= k[A][B] or Rate= k [A]2 units: M-1s-1
Reaction Pathways for Forward and Reverse Reactions:
An Energy Diagram
For Forward reaction:
w =reactants
a =activation energy, Ea
y = activated complex
c= energy change in the
reaction, Delta E
z = products
b = energy needed to
activate the reverse
reaction ( endothermic) Ea’
Forward reaction Reverse reaction
 exothermic  endothermic
Reaction Mechanism: A reaction mechanism details the individual steps
that occur in the course of a reaction.Each of these steps are called as
elementary steps. An elementary step may produce an intermediate, a
product that is consumed in a later elementary step and therefore does
not appear in the overall stoichiometry of the reaction.
If a mechanism has several elementary steps, then the overall rate is
determined by t he slowest step, called the rate-determining step.
Collision Theory: Atoms, ions and molecules
can react to form products when they collide with
one another, if they have
1. enough KE and 2. the right orientation.
Activation Energy: Ea, The minimum amount of
energy that particles must have in order to react.
Activated complex: an unstable arrangement of
atoms that forms momentarily at the peak of the
activation energy.
Factors Affecting Reaction Rate:
1.Temperature: increase temp, increase collisions, &
increase # collisions with enough kinetic energy to get
over Ea barrier.
2.Concentration: increase concentration increases
frequency of collisions leads to higher reaction rate.
3.Particle Size: increase in surface area increases
amount of reactant exposed, increases collision
frequency and reaction rate.
4.Catalysts: increases rate of reaction without being
used up during the reaction. Lower activation energy
barrier so more reactants have the energy to form
products within a given time.