Chemistry II
Chapter 5
Thermochemistry
1. The study of energy and its
transformations is known as
_thermodynamics__.
2. __Thermochemistry_____ considers the
relationships between chemical reactions
and the energy changes involving heat.
Nature of Energy
• The magnitude of the kinetic
energy, Ek, of an object depends
on its __mass__ (m) and
__velocity___(v)
•
Formula for kinetic energy:
•Ek = ½
2
mv
4. Potential energy is the energy an object
has by virtue of its ___position__ relative to
other objects.
Types of potential energy
• gravitational
•
• elastic
•
• chemical
•
• electrical
5.__Electrostatic Forces_________ results
from the interactions between charged
particles.
The formula for electrostatic force
is:
•Ee1 = kQ1Q2/ d
k is the constant of proportionality
(8.99 x 109 J m /C2)
- C is the _coulomb______, the unit for
electrical charge.
- Q1 and Q2 are the _charges__ on the two
particles involved
- d is the distance between the particles
8. When the charges are the
same, the electrostatic force, Ek,
is ______positive_______ and the
charges __repel____.
9. When the charges are opposite
the electrostic force, Ek, is
____negative_____ and the charges
__attract____.
•
• The energy of a substance is due to
• 1.) The thermal (kinetic) energy of
the molecules.
• 2.) The potential energy of the
bonds.
11. The SI unit for energy is the
__joule____ (J) named in honor of
___James Joule____ ( 1818 -1889) – a
British Scientist who investigated heat
and work.
• 12. A __calorie_______ was originally
defined as the amount of energy
required to raise the temperature of
1gram of water from 14.5 to 15.5
degrees Celsius
13. Today energy is defined in terms of the joule:
1 calorie = 4.184 joules
14. When we use thermodynamics to analyze
energy changes, we focus our attention on a
limited and well defined part of the universe.
a. The portion we single out to look at is called
the __system___.
b. Everything else is the _surroundings____.
c. A ___closed system_____ can exchange energy
but not material with its surroundings.
Transferring Energy: Work and Heat
Energy is transferred in two ways:
1.) To cause the motion of an object
against a force
2.) To cause a temperature change
15. The energy used to cause an object to
move against a force is called __work_____.
• W= F x D
16. The other way that energy is transferred is
__heat__- the energy transferred by a hotter object
to a colder one.
17. _Energy______ is the capacity to do
work or to transfer heat.
Section 2
The First Law of Thermodynamics
18. The First Law of Thermodynamics –
Energy is conserved
19. The _Internal_____ __energy___ of a system
is the sum of all of the kinetic and potential
energy in the components of the system.
• 20. We represent the internal energy with
the symbol _E__.
21. We can never really know the internal
energy of a system, E, but we can find the
____Change____ in E (ΔE)
• ΔE = Ef - Ei
22. A __Positive_____ value for ΔE
results when Efinal> Einitial which means
that energy has been put into the
system ( __Endothermic___)
23. A ___negative__ value for ΔE
results when Efinal< Einitial which means
that energy has been lost to the
surroundings ( _exothermic___)
Energy Diagrams
Exothermic
* Endothermic
24. When a system undergoes any chemical
or physical change, the change in internal
energy ΔE is given by the heat, q, added to
or released by the system plus the work
done on or by the system, w.
Formula:
ΔE = q + w
When heat, q, is added to a system q is positive
When work, w, is done on a system it has a
positive value
Sample exercise 5.3
The hydrogen and oxygen gases in the cylinderillustrated
in figure 5.3 are ignited. As the reaction occurs, the system
loses 1150 J of heat to the surroundings. The reaction also
causes the pistion to rise as the hot gases expand. The
expanding gas does 480 J of work on the surroundings as it
pushes against the atmosphere. What is the change in the
internal energy of the system ?
Endothermic –
heat flows into the system
Exothermicheat flows out of the system
State Functions
Although we cannot know exactly what the
internal energy, E, of a system, it does have
a definite value, Δ E can be determined.
26. Internal energy is a __state___ _function___.
A state function is a property of a
system that is determined by specifying
the conditions of state.
27. The value of a state function depends
only on its _present____ __conditon__ and
not on the particular ___history___ of the
sample.
Internal
energy, E, a
state function.
Caption:
E depends only on the present state of the system and not on the path by which
it arrived at that state. The internal energy of 50 g of water at 25 °C is the same
whether the water is cooled from a higher temperature to 25 °C or warmed from
a lower temperature to 25 °C.
28. B/c internal energy, E, is a state
function Δ E depends only on the
__initial_____ and _final___ states and not
the change that occurred.
29.)Internal energy is a state function,
but q+w is not a state function.
Section 3: Enthalpy
• 30.)Most commonly the only kind
of work produced by chemical
change is _____mechanical_ work.
• 31.)Usually reactions are carried
out at __constant___atmospheric
pressure.
For example: Consider the following equation
Zn + 2 HCl → ZnCl2 + H2
Net Ionic
Zn + 2H+ → Zn+2 + H2
32.) As the hydrogen gas is produced it
does positive__ against the atmosphere.
We can better see this work if the reaction takes place in a reaction vessel like the one below.
33.) This kind of work is called _Pressure__ __volume__ work (P-V Work).
• Formula:
• W= -PΔV
• 34. When the change in volume is positive
the work done is ___negative__.
• 35. The thermodynamic function called
___enthalpy___ accounts for the heat
flow in chemical changes occurring at
constant pressure when no forms of
work are performed other than P-V work.
Enthalpy equals the internal energy plus
the product of the pressure and the change
in volume of the system.
• 36.) The symbol for enthalpy is
_____H_____.
Formula for enthalpy:
H = E + PV
•Enthalpy (H) equals the total
internal energy ( E) and pressure
volume work.
37.) Enthalpy is a __state____ function because
internal energy, E, pressure and volume are all
state functions.
ΔH = ΔE + ΔPV
W = - PΔV
-W = P ΔV
ΔH = ΔE + ΔPV= q + w-w=q ( at constant pressure)
ΔH = qp
Change in enthalpy equals the heat gained or lost
at constant pressure.
It is easier to measure the change in temperature
than the change in internal energy.
• 38.) When ΔH is ____positive____ (qp
is positive) the system has gained heat
from the surroundings.
39.) When ΔH is positive the reaction is
_endothermic____.
40.) When ΔH is negative heat is
___lost___to the surroundings and the
reaction is __exothermic___.
41. Because H is a state function, ΔH
depends only on the __initial__ and the
__final____ states of the system.
• In section two we learned that q is not a
state function. This is not a contradiction
of ΔH = qp . Enthalpy change, ΔH, only
is equal to the heat released or absorbed
at __constant__ pressure.
Section 4: Enthalpies of Reaction
ΔH= ΔHfinal + ΔHinitial
ΔH= ΔHproducts + ΔHreactants
42. The enthalpy change that accompanies
a reaction is called the __enthalpy of
reaction_ or heat of reaction (ΔHrxn)
(rxn is often used as an abbreviation for
reaction)
Consider the following reaction:
2H2 (g) + O2 (g) →
2H2O (g) ΔH= -483.6 KJ
• 43. When the above reaction is controlled so
that exactly 2 moles of hydrogen gas react with
exactly one mole of oxygen gas, 483.6
kilojoules of heat is released. Change in
enthalpy (ΔH) is negative so the reaction is
___exothermic___.
44. Balanced chemical equations that show
the associated enthalpy changes are called
__thermochemical____ equations.
45. The enthalpy change accompanying a
reaction can also be represented in an
enthalpy _diagram_____.
46. In an ___exothermic_____ reaction the
enthalpy of the products will be lower than
the reactants.
47. In an ___endothermic____ reaction the
enthalpy of the products will be higher
than the reactants
The following guidelines are helpful when
using thermochemical equations and
enthalpy diagrams.
• 48.)Enthalpy is an extensive property.
• a.)_Extensive___ properties depend on the amount
of substance present (i.e., mass, volume, amount of
heat)
• b.)_Intensive____ properties do not depend on the
amount of substance present ( i.e., density, melting
point, boiling point
49.) ΔH is directly dependent on the
__amount___ of reactant consumed.
For example,
If when 2 moles of hydrogen gas react with
1 mole of oxygen gas 493.6 KJ of heat is
released.
2H2(g) + O2(g)  2H2O(g) ΔH= -493.6 KJ
Then when 4 moles of hydrogen gas react
with 2 moles of oxygen gas 985.2 KJ are
released.
Sample Example 5.5
How much heat is released when 4.5 g of methane gas is
burned in a constant pressure system?
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) ΔH= - 890 KJ
Practice Exercise
Hydrogen Peroxide can decompose to water and
oxygen by the following reaction:
2H2O2  2H2O(l) + O2(g) Δ H = -196 KJ
Calculate the value of q when 5.00 grams of
H2O2(l) decomposes at constant temperature.
50.) The enthalpy change for a reaction is
equal in magnitude but opposite in sign to
ΔH for the reverse reaction.
For example:
CH4(g)
+
2O2(g) → CO2(g) + 2H2O(l)
CO2(g) + 2H2O(l) → CH4(g)
+
ΔH=
-890 KJ
2O2(g) ΔH= __890_ KJ
51.)When we reverse a reaction, the roles of
reactants and products are reversed.
52.)The enthalpy change for a reaction depends
on the state of the reactants and products.
i.e., if the product of combustion of methane
were gaseous H2O instead of liquid H2O, the
ΔH would be= -802 KJ instead of -890 KJ.
53.)It is very important to specify the
___states____ of reactants and products in
thermochemical equations.
Definitions
• 54.) ___Enthalpy______ (H) A quantity
defined by the relationship H = E + PV ;
the enthalpy change, ΔH, for a reaction
that occurs at constant pressure is the
heat evolved or absorbed in the reaction;
ΔH=qp
• 55)Enthalpy of formation(Hf)• the enthalpy change that accompanies the
formation of a substance from the most stable
forms of its component elements.
• 56.)Enthalpy of reaction (Hrxn)
• The enthalpy change associated with a
chemical reaction.
Using Enthalpy as Guide
• 57.) A process that is thermodynamically
favored to happen is called a
___spontaneous______ process.
• 58.) By “spontaneous” we don’t mean that
the reaction will form products without
intervention. Most often some amount of
energy must be imparted to get the process
started.
•
• 59.)The enthalpy change of a
reaction does give an indication as
to whether the reaction is likely to
be __spontaneous__.
• 60.)As a general rule, when the enthalpy
change is _large_ it is the dominant factor
in determining spontaneity.
• 61.)When ΔH is large and negative
the reaction tends to be
__spontaneous_____.
• 62.)Reactions for which ΔH is large
and positive tend to be spontaneous in
the __opposite___ direction
Section 5 : Calorimetry
• 63.)The measurement of heat flow is
_calorimetry_______
• 64.)An apparatus used to measure heat flow is
a _________________.
Heat Capacity and Specific Heat
• 65.)The ___heat_ _capacity__ of an object
is the amount of heat required to raise the
temperature by 1K or 1⁰C. ( one degree
Kelvin is equal to one degree _celsius___)
1K = 1⁰C
• 66.)The greater the heat capacity the
_greater___ the heat required to
produce a given rise in temperature.
• 67.)The heat capacity of 1 mole of a
substance is called its molar heat
capacity.
• 68.)The heat capacity of 1 gram of a substance
is called its specific heat capacity.
• 69.)Formula for specific heat capacity:
• C =q/mΔT
•
•
•
•
C – specific heat
q-heat
m-mass
ΔT – change in temp
• 70.)The specific heat capacity for water is
much __higher__ than for other
substances. Water can _absorb/release___
much heat with little change in its
temperature__.
• 71.)The high specific heat of water is
essential for the maintenance of our
____climate______
Sample exercise 5.6
• How much heat is needed to warm 250 g of water
(about 1 cup) from 22ᵒC (about room temperature) to
near its boiling point, 98 ⁰C? The specific heat of
water is 4.18 J/g-K.
•
• What is the molar heat capacity of water?
Practice Exercise
• Large beds of rock are used in solar heated homes to store heat.
Assume that the specific heat of the rocks is 0.82J/g-K. Calculate the
quantity of heat absorbed by 50 kg of rocks if their temperature
increases by 12 ᵒC.
•
• What temperature change would these rocks undergo if they emitted
450 KJ of heat?
Constant Pressure Calorimetry
• 72.)In calorimetry problems we will
assume that the heat produced by the
reaction, qrxn is entirely
____absorbed_____ by the solution; it
does not escape the calorimeter.
• 73.)For an exothermic reaction, heat is
_produced/released___by the reaction and
absorbed___ by the solution.
• 74.)For an endothermic reaction opposite
is true- heat is __absorbed______ by the
reactants and is _lost____ by the solution.
• qsoln = -q rxn
• qsoln = (specific heat of soln) x ( gram of
soln)x ΔT)= -q rxn
• 75.) For dilute aqueous solutions,
the specific heat of the solution
will be the same as water
( 4.18 J/g-k)
Sample Exercise 5.7
• When a student mixes 50 mL of a 1.0 M HCl and 50mL
of a 1.0M NaOH in a coffee cup calorimeter, the
temperature of the resultant solution increases from 21.0
⁰C to 27.5⁰C. Calculate the enthalpy change for the
reaction, assuming that the calorimeter loses only a
negligible quantity of heat, that the total volume of the
solution is 100 mL , that its density is 1.0 g/mL and that
its specific heat is 4.18 J/g-K.
Practice Exercise
• When 50 mL of a 0.100M AgNO3 and 50mL of a 0.100M HCl are
mixed in a constant –pressure calorimeter, the temperature of the
mixture increases from 22.30⁰C to 23.11⁰C. The temperature increase is
caused by the following reaction:
• AgNO3(aq) + HCl(aq) → AgCl(s) + HNO3(aq)Calculate ΔH
for this reaction, assuming that the combined solution has a mass of
100 g and a specific heat of 4.18 J/g-K
Bomb Calorimetry: Constant –Volume
Calorimetry
• 76.)One of the most important types of
reactions studied using calorimetry is
__combustion_____ in which a compound
reacts completely with excess oxygen.
77.)Combustion reactions are most conveniently
studied using a bomb calorimeter.
• 78.)The substance to be studied is
placed in a small cup within a sealed
vessel called a _bomb__ , which is
designed to withstand high pressures.
• 79.)__Heat___ is released when
combustion occurs.
• 80.)The heat is absorbed by the
calorimeter contents causing a rise
in the _temperature______ of the
water.
• 81.)The temperature of the water is
carefully measured __before____ and
_after__ reaction.
• 82.)To calculate the heat of combustion
from the measured temperature increase
in the bomb calorimeter, we must know
the heat capacity of the calorimeter ,Ccal.
• Example Problem
• The combustion of 1 gram of benzoic acid is
known to produce 26.38 KJ of heat. If exactly
1 gram of benzoic acid is combusted in a
bomb calorimeter and it increases the
temperature by 4.857ᵒC, what is the heat
capacity for that calorimeter?
•
• Ccal = 26.38 KJ / 4.857 ᵒC = 5.431 KJ/ᵒC
• 83.)To find the heat of a reaction, qrxn
using a specific calorimeter:
• Qrxn = -Ccal x ΔT
Sample exercise 5.8
• Methylhydrazine (CH6N2) is commonly used as a liquid rocket fuel. The
combustion of methylhydrazine with oxygen produces N2(g), CO2(g) and
H2O(l):
• 2 CH6N2(l) + 5O2(g) → 2N2(g) + 2 CO2(g) +6 H2O(l)
• When 4.00 g of methylhydrazine is combusted in a bomb calorimeter, the
temperature of the calorimeter increases from 25.00ᵒC to 39.50 ᵒC. In a
separate experiment the heat capacity of the calorimeter is measured to be
7.794KJ/g-ᵒC. What is the heat of reaction for the combustion of one mole
of CH6N2 in this calorimeter?
Practice exercise
• A 0.58665g sample of lactic acid (HC3H5O3)
is burned in a calorimeter whose heat
capacity is 4.812 KJ/ᵒC . The temperature
increases from 23.10ᵒC to 24.95ᵒC. Calculate
the heat of combustion of (a) lactic acid per
gram (b) per mole.
• 84.Because reactions in a bomb calorimeter are
carried out under constant volume conditions, the
heat transferred corresponds to a change in internal
energy ΔE, rather than __enthalpy____ ΔH.
However both ΔE and ΔH are usually very close in
value.
Section 6 Hess’s Law
• 85.)It is possible to calculate ΔH for a
reaction from the __tabulated___ ΔH
values of other reactions.
If a particular reaction can be carried out in one step or in
a series of steps the sum of the enthalpy changes
associated with the individual steps must be the same as
the enthalpy change associated with the one step process.
CH4(g) + 2O2(g) → CO2(g) +2 H2O(g)
2H2O(g) → 2H2O(l)
CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)
ΔH= -802 KJ
ΔH= -88KJ
ΔH= -890 KJ
To obtain the net equation
1.)The sum of the reactants of the two equations is
placed on the left side of the arrow.
2.)The sum of all products is placed on the product
side
3.)Any substance that is on both sides of the equation
in the same phase can be canceled like in algebraic
equations.
Hess’s Law:
• If a reaction is carried out in a series of
steps, ΔH for the reaction will equal the
sum of the enthalpy changes for the
individual steps.
• The overall enthalpy change for the process is
independent of the number of steps or path in
which the reaction occurs.
• If a reaction is reversed, change the sign of Δ
H.
• If a reaction is multiplied by a factor, multiply
the enthalpy change by that same factor.
• 85.)Hess’s Law provides a useful
means of calculating _energy changes
that are difficult to measure directly.
The enthalpy of combustion of C to CO2 is
-393.5kJ/mol carbon, and the enthalpy of
combustion of CO to CO2 is -283.0 kJ/mol CO:
C(s) + O2(g) → CO2(g)
ΔH= -393.5kJ
CO2(g) + ½ O2(g) → CO2(g) ΔH = -283.0 kJ
Using the above data , calculate the enthalpy of combustion of C to CO:
C(s) + ½ O2 → CO(g)
Practice Exercise
• Carbon occurs in two forms, graphite and diamond. The enthalpy of
combustion of graphite is -393.5KJ /mol and that of diamond is 395.4kJ/mol:
•
•
•
•
C(graphite) + O2(g) → CO2(g)
ΔH= -393.5 kJ
C(diamond) + O2(g) → CO2(g)
ΔH= -395.4 kJ
Calculate ΔH for the conversion of graphite to diamond:
C(graphite) → C(diamond)
Sample Exercise 5.10
• Calculate ΔH for the reaction
• 2C(s) + H2(g) → C2H2(g)
• given the following reactions and their respective enthalpy changes:
C2H2(g) + 5/2 O2(g) →2CO2(g) + H2O(l)
ΔH=-1299.6 kJ
C(s) + O2(g) → CO2(g)
ΔH = -393.5 kJ
H2(g) + ½ O2(g) → H2O(l)
ΔH = -285.8 kJ
Calculate ΔH for the reaction
NO(g)
+ O(g) →
NO2 (g)
Given the following information:
NO(g) + O3(g) → NO2(g) + O2(g)
ΔH= -198.9 kJ
O3(g) → 3/2 O2(g)
ΔH= -142.3 kJ
O2(g) → 2 O(g)
ΔH= 495.0 kJ
We always get the same value of ΔH for an
overall reaction regardless of how many
steps we use to get to the final product.
We can use Hess’s Law to find the enthalpy
change for a large number of reactions from
the tabulated ∆H values.
86.)Tables exist for the following:
a.) Enthalpy of vaporization- ΔH for
converting liquids to gases
b.) Enthalpy of fusion- ΔH for melting
solids
86.)Tables exist for the following:
c.) Enthalpy of combustion- ΔH for
combusting a substance with oxygen
d.) Enthalpy of formation- ΔH for
formations of a substance from its
elements in standard state.
∆H is dependent on temperature and pressure;
so, we define the conditions at standard state.
87.)What is standard state for most
enthalpies?
• 1 atmosphere of pressure and 25 ⁰ C
88.)What is the standard enthalpy of
formation?
• ΔHf ⁰ the change in enthalpy for the
reaction that forms 1 mole of the
compound from its elements with all
substances in their standard state
• The stoichiometry is always arranged so that 1
mole of the product is formed.
•
Sample exercise 5.11
For which of the following reactions at 25◦C would
the enthalpy change represent a standard enthalpy of
formation? For those where it does not, what changes
would need to be made in the reaction conditions?
a.)2Na(s) + ½ O2(g) → Na2O(s)
b.)2K(l) + Cl2(g) → 2KCl(s)
c) C6H12O6(s) → 6C(diamond) + 6H2(g) + 3O2(g)
Practice Exercise
Write the equation corresponding to the
standard enthalpy of formation of liquid CCl4
(Carbon tetrachloride.)
We can use ∆Hf◦ of different compounds to
calculate ∆H for a reaction at standard state.
Determine ∆H◦ for the following reaction.
C3H8 (g) + 5O2(g) → 3 CO2 (g) + 4 H2O(l)
◦
∆Hrxn
On the bottom line:
◦
◦
= Σ n ∆Hf product - Σ m ∆Hf reactants
Example 5.12
Calculate the standard enthalpy of
change for the combustion of 1 mole of
benzene (C6H6) to CO2 (g) and H2O (l)
Practice
Given the following standard enthalpy of
reaction, use the standard enthalpies of
formation in table 5.3 to calculate the standard
enthalpy of formation of CuO(s)
Food and Fuels
Section 8
• 89.)What is a fuel value?
•
The energy released when 1 gram of a
material is combusted.
The fuel value for any food or fuel can be
measured by calorimetry.
• Protein has 17 kJ/g
• Carbohydrates 17 kJ/g
• Fats 38 kJ/g
90. Fats are the major energy storage molecules in
the body for two reasons:
• 1. insoluble in water
• 2. produce more energy per gram than
carbohydrates or proteins.
Read exercise 5.14 page 182
• A plant such as celery contains carbohydrates in the form
of starch and cellulose. These two kinds of carbohydrates
have essentially the same fuel values when combusted in a
bomb calorimeter. When we consume celery, however, our
bodies receive fuel values from the starch only. What can
we conclude about the starch stored as cellulose as foods?
• A 28 –g ( 1 oz) serving of a popular breakfast cereal served with 120 mL of
skim milk provides 8 g of protein, 26 g of carbohydrates and 2 g of fat. A)
Using the average fuel values of these kinds of substances, estimate the
amount of food energy in this serving. B.) A person of average weight uses
about 100 Cal/mile when running or jogging. How many servings of this
cereal provide the fuel value requirement for running 3 miles.
• Practice Exercise
• Dry red beans contain 62% carbohydrate, 22%
protein, and 1.5% fat. Estimate the fuel value of
these beans. B.) Very light activity like reading or
watching television uses about 7 kJ/min. How
many minutes of such activity can be sustained by
the energy provided by a serving of chicken
noodle soup containing 13 g or protein, 15 g
carbohydrate and 5 g of fat.
Fuels:
The greater the % of carbon and hydrogen in a
fuel the higher its fuel value
• Fossil fuels- coal, petroleum and natural gas
• Natural gas- gaseous hydrocarbons (primarily
methane) ethane, propane and butane.
• Petroleum – a liquid composed of hundreds of
compounds most of which are hydrocarbons
containing sulfur, nitrogen and oxygen
• Coal- solids containing hydrocarbons of high sulfur
molecular weight as well as compounds containing
sulfur, oxygen and nitrogen.
Compare the quantity of heat produced by the
combustion of 1 gram of propane compared
to 1 gram of benzene.
Exercise 5.13
The standard enthalpy change for the reaction
CaCO3(s) → CaO(s) + CO2(g)
Is 178.1 KJ. From the values for the standard enthalpies
of formation of CaO(s) and CO2(g) given in table 5.3
calculate the enthalpy of formation of CaCO3(s).