Download Chapter 5 Thermochemistry

Chapter 5
Thermochemistry
Recommended Text
Problems
„
6, 25, 27, 31, 33, 35, 37, 45abd,
„
51, 53, 59, 61, 63, 67ac, 73, 75,
„
83, 90, 93, 98
System, Surroundings,
and Universe
„
Definitions:
„
The system
„
The surroundings
„
The universe
„
Energy
„
Units of Energy
„
„
The SI unit of energy is the joule (J).
An older, non-SI unit is still in widespread
use: the calorie (cal).
Energy
„
The ability to do work or transfer heat.
„
Potential energy
„
„
Energy an object possesses by virtue of its
position or chemical composition.
Kinetic energy
„
1
2
Energy an object possesses by virtue of its
motion.
Example
„
Consider the kinetic energy of a person
whose mass is 130 lb (59.0 kg)
traveling in a car at 60 mph (26.8 m/s).
Energy
„
The ability to do work or transfer heat.
„
Work
„
Energy used to cause an object that has mass to
„
„
where w is work, F is the force, and d is the distance.
Heat
„
Energy used to cause the
„
Heat flows from
Example:
„
A bowler lifts a 5.4-kg (12-lb) bowling ball from
ground level to a height of 1.6 m (5.2 feet) and then
drops the ball back to the ground.
a.
b.
c.
What happens to the potential energy of the bowling ball as
it is raised from the ground?
What quantity of work, in J, is used to raise the ball?
After the ball is dropped, it gains kinetic energy. If we
assume that all of the work done in part (b) has been
converted to kinetic energy by the time the ball strikes the
ground, what is the speed of the ball at the instant just
before it hits the ground?
More Example:
„
Two gases, A(g) and B(g), are confined in a cylinderand-piston. Substances A and B react to form a solid
product:
A(g) + B(g) → C(s).
As the reaction occurs, the system loses 1150 J of
heat to the surroundings. The piston moves
downward as the gases react to form a solid. As the
volume of the gas decreases under the constant
pressure of the atmosphere, the surroundings do 480
J of work on the system. What is the change in the
internal energy of the system?
First Law of Thermodynamics
„
Energy is neither created nor
destroyed.
„
The total energy of the universe is a
constant;
„
„
if the system loses energy, it must be
gained by the surroundings, and vice versa.
Energy can be converted from one type
to another.
Internal Energy
„
Definition:
„
„
The sum of all kinetic and potential
energies of all components of the system;
E.
The change in internal energy, ∆E,
„
The final energy of the system minus the
initial energy of the system:
∆E = Efinal − Einitial
Changes in Internal Energy
„
If ∆E > 0, Efinal >
Einitial
„
„
The system
energy
from the
surroundings.
The
process.
„
If ∆E < 0, Efinal <
Einitial
„
„
The system
energy
to the
surroundings.
The
process
Changes in Internal Energy
„
When energy is
exchanged between
the system and the
surroundings, it is
exchanged as either
heat (q) or work (w).
∆E = q + w.
Exchange of Heat between
System and Surroundings
„
Endothermic process
„
Exothermic process
Internal Energy
„
The internal energy of a system is
independent of the path by which the system
achieved that state.
„
In the system below, the water could have
reached room temperature from either direction.
State Function
„
It depends on the
„
Example:
„
Enthalpy, Entropy, Helmholtz free energy,
Temperature, Gibbs free energy, Fugacity, Density,
Internal Energy
Work
„
„
We can measure the work done by
the gas if the reaction is done in a
vessel that has been fitted with a
piston.
We know
w =Fxd
Work
„
„
If the system does work
If the work is done on the
system
Enthalpy
„
„
Enthalpy, H, is a type of chemical
energy, sometime referred to as heat
content
Enthalpy is the internal energy plus the
product of pressure and volume:
Constant Pressure
„
When the system changes at constant pressure, the
change in enthalpy, ∆H, is
Example:
a.
b.
How much heat is needed to warm
250 g of water (about 1 cup) from 22
°C (about room temperature) to near
its boiling point, 98 °C? The specific
heat of water is 4.18 J/g-K.
What is the molar heat capacity of
water?
Summary
Enthalpy of Reaction
„
„
The change in enthalpy, ∆H, is the
enthalpy of the products minus the
enthalpy of the reactants:
This quantity, ∆H, is called the enthalpy
of reaction, or the heat of reaction.
Manipulation of Enthalpy
1.
2.
3.
Enthalpy is an extensive property.
∆H for a reaction in the forward
direction is equal in size, but opposite
in sign, to ∆H for the reverse reaction.
∆H for a reaction depends on the
state of the products and the state of
the reactants.
Thermochemical Equations
„
A thermochemical equation
„
The chemical equation for a reaction (including
phase labels) in which the equation is
„
It is important to note
Remarks on Thermochemical
Equations:
„
The sign of ∆H indicates whether the reaction, when carried out
at constant pressure,
„
In interpreting a thermochemical equation
„
The phases (physical states) of all species must be specified
„
The value quoted for ∆H applies when products and reactions
are at the same temperature
Measurement of Heat Flow;
Calorimetry
„
„
Calorimetry is the science of
measuring the heat of chemical
reactions or physical changes.
Calorimeter
„
The wall of the calorimeter is isolated
„
„
No heat enters or escapes from the
calorimeter
It contains water and/or other materials
of known heat capacity
Heat Capacity and Specific
Heat
„
Heat capacity
„
Specific heat capacity (specific heat)
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:
Example:
„
When a student mixes 50 mL of 1.0 M HCl
and 50 mL of 1.0 M NaOH in a coffee-cup
calorimeter, the temperature of the resultant
solution increases from 21.0 °C to 27.5 °C.
Calculate the enthalpy change for the
reaction in kJ/mol HCl, 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.
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.
Example:
„
Methylhydrazine (CH6N2) is used as a liquid rocket
fuel. The combustion of methylhydrazine with oxygen
produces N2(g), CO2(g), and H2O(l):
2 CH6N2(l) + 5 O2(g) → 2 N2(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 25.00 ˚C to 39.50 ˚C. In a separate
experiment the heat capacity of the calorimeter is
measured to be 7.794 kJ/ ˚C. Calculate the heat of
reaction for the combustion of a mole of CH6N2.
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.
We can estimate ∆H using published ∆H
values and the properties of enthalpy.
„
Why & how?
Hess’s Law
„
Hess’s law of heat summation
states
„
„
for a chemical equation
the enthalpy change for the overall
equation
Rules For Manipulating
Thermochemical Equations:
∆ Ho = -890 kJ/mol
When a thermochemical equation is multiplied by
any factor,
CH4 + 2O2 → CO2 + 2H2O
„
„
„
When a chemical equation is reversed,
„
„
the value of ∆H for the new equation is obtained by
the value of ∆H is
The value of ∆H for a reaction is the same whether it
occurs in one step or in a series of steps
Example:
„
The enthalpy of reaction for the
combustion of C to CO2 is –393.5
kJ/mol C, and the enthalpy for the
combustion of CO to CO2 is –283.0
kJ/mol CO:
More Example:
„
„
Calculate ΔH for the reaction
2 C(s) + H2(g) → C2H2(g)
given the following chemical equations and their
respective enthalpy changes
Calculate ΔH for the reaction
NO(g) + O(g) → NO2(g)
given the following information:
Enthalpies of Formation
„
An enthalpy of formation, ∆Hf,
„
„
The enthalpy change for the reaction in
which a compound is made from
in their
elemental forms.
Standard enthalpies of formation, ∆Hf°,
„
Measured under standard conditions (25 ˚C
and 1.00 atm pressure).
Example:
„
For which of the following reactions at 25˚C
would the enthalpy change represent a
standard enthalpy of formation? For each that
does not, what changes are needed to make
it an equation whose ΔH is an enthalpy of
formation?
Calculation of ∆H
„
We can use Hess’s law in this way:
Example
„
Large quantities of ammonia are used to
prepare nitric acid according to the following
equation:
„
What is the standard enthalpy change for this
reaction?
Calculation of ∆H
C3H8 (g) + 5 O2 (g) ⎯→ 3 CO2 (g) + 4 H2O (l)
Calculation of ∆H
C3H8 (g) + 5 O2 (g) ⎯→ 3 CO2 (g) + 4 H2O (l)
C3H8 (g) ⎯→ 3 C (graphite) + 4 H2 (g)
3 C (graphite) + 3 O2 (g) ⎯→ 3 CO2 (g)
4 H2 (g) + 2 O2 (g) ⎯→ 4 H2O (l)
C3H8 (g) + 5 O2 (g)
⎯→ 3 CO2 (g) + 4 H2O (l)
Example
„
For example, suppose you are given the
following data:
• Could you use these data to obtain the
enthalpy change for the following reaction?
Practice
„
From the following data,
CH4 + 2O2 → CO2 + 2H2O
H2O(l) → H2O(g)
∆Ho = -890 kJ/mol
∆Ho = 44 kJ/mol at 298 K
Calculate the enthalpy of the reaction
CH4 + 2 O2(g) → CO2(g) + 2 H2O(g)
∆Ho = ?
Challenging
„
From the following enthalpies of reactions:
1.
2.
3.
4.
5.
6.
7.
8.
„
2 O(g) -> O2(g)
H2O(l) -> H2O(g)
2 H(g) + O(g) -> H2O(g)
C(graphite) + 2 O(g) -> CO2(g)
C(graphite) + O2(g) -> CO2(g)
C(graphite) + 2 H2(g) -> CH4(g)
2 H(g) -> H2(g)
H2O(l) -> H2O(g) ∆H = 41 kJ/mol at
∆Ho = -249 kJ/mol
∆Ho = 44 kJ/mol at 298 K
∆Ho = -803 kJ/mol
∆Ho = -643 kJ/mol
∆Ho = -394 kJ/mol
∆Ho = -75 kJ/mol
∆Ho = -436 kJ/mol
373 K, non-standard condition
Calculate the heat of combustion of methane into gaseous H2O.
Energy in Foods
„
„
Most of the fuel in the
food we eat comes from
carbohydrates and fats.
The vast majority of the
energy consumed in this
country comes from
fossil fuels.
Definitions/Laws
„
„
„
„
„
„
„
„
„
„
The system
The surroundings
The universe
Potential energy
Kinetic energy
Energy
First law of
thermodynamics
Endothermic process
Exothermic process
State function
„
„
„
„
„
Heat capacity
Specific heat capacity
(specific heat)
Hess’s law of heat
summation
Manipulating
thermochemical
equations
Enthalpy of formation,
∆Hf,
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