Transcript Thermodynamics

Thermodynamics
Thermodynamics responds
to four questions:
Can the reaction occur spontaneously or
must it be driven?
How much useful work is produced or
consumed when the reaction takes place?
How far will the reaction go?
How can the reaction be made to advance
further?
Chemical reactions involve
changes in energy.
Thermodynamics is the study of energy and
its transformations.
Force is a push or pull on an object.
Work is the product of force applied to an object over a
distance.
W = F x d.
Energy is the work done to move an object against a
force.
Heat is the transfer of energy between two objects.
Energy is the capacity to do work or transfer heat.
Heat vs. Temperature
Heat is energy that is transferred from one
object to another because of a difference in
temperature.
Temperature is a measure of the average
kinetic energy of the molecules of a substance.
Kinetic Energy
Energy of motion
1/2 mv2
depends on mass and velocity of object
Potential Energy
Energy by virtue of position relative to
other objects (mgh)
Stored energy
Arises when there is a force acting on an
object
Potential energy can be converted into
kinetic energy. Ex: a bicyclist at the top
of a hill.
Kinetic and Potential Energy
Electrostatic potential energy, Ed, is the
attraction between two oppositely charged
particles, Q1 and Q2, a distance d apart:
Ed = κQ1Q2
d
The constant κ = 8.99 x 109 J-m/C2
If the two particles are of similar charge, then Ed
is the electrostatic repulsion between them.
Units of Energy
SI unit = Joule (J)
1 J = 1 kg-m2/s2
 2 kg object moving at a speed of 1 m/s has a kinetic
energy of 1 J.
1 cal = 4.184 J
A nutritional Calorie: 1 Cal = 1000 cal = 1kcal.
System and Surroundings
System = part of the universe we are
interested in; what is under
study.
Surroundings = everything else in the
universe.
Energy lost by system =
Energy gained by surroundings.
 Energy out of system = - E
 Energy into system = + E
First Law of
Thermodynamics
Internal energy: total energy of a system.
Cannot measure absolute internal energy.
Change in internal energy:
E = Efinal - Einitial
E - system loses
E + system gains
First Law of
Thermodynamics
Energy can neither be created nor
destroyed.
Systems tend to lose energy to be as
stable as possible.
Energy of system + surroundings is
constant.
Any energy lost by a system must be
transferred to the surroundings.
Energy is exchanged by changes in heat or work.
E = q + w The internal energy is the sum of the kinetic
and potential energies of all the particles in a system.
Internal energy can be changed by work or heat flow
or both.
q + heat in
q - heat out
w + work done on
w - work done by
Calculate the change in energy of the system for a
process where the system absorbs 140 J of heat
from the surroundings and does 85 J of work on the
surroundings.
State v. Path Functions
State:
•property of a system not determined by its history
•depends only on its present condition
•E is a state function
Path:
•depends on manner in which process is carried out
•q and w are path functions
State functions depend only on the initial and final
states of the system, not on how the internal energy
is used.
Enthalpy
Chemical reactions have the ability to do
work.
In chemistry most mechanical work is PV
work.
When a gas is produced in a reaction, the
gas can be used to push a piston, thus
doing work.
For a constant pressure process, the
changes in internal energy and enthalpy
are related by the equation: H=E +PV
Enthalpy
When the pressure is constant, w = -PΔV
Like internal energy, enthalpy cannot be
determined exactly, but changes in
enthalpy (H) can be determined.
Consider work done under constant
pressure:
E = qp - P V
or
qp = E + P V
H = E + P V
so qp = H
H = qp
H, enthalpy, like internal energy is a state function.
AKA heat of reaction, the enthalpy of reaction.
Enthalpy
Enthalpy is the heat exchanged between the system and
surroundings at constant pressure, as the result of a
chemical reaction.
H = H(products) - H(reactants)
The enthalpy change of a reaction is equal in magnitude but
opposite in sign of the enthalpy change of the reverse reaction:
Hforward = - Hreverse
The enthalpy change of a reaction depends on the states of
matter of the products and reactants.
Exothermic Reactions
Heat flows out of system
Heat is a product
System has less energy
H is negative to indicate direction of
heat flow
Energy released making bonds > energy
absorbed to break bonds
Product bond energy > reactant bond
energy
Endothermic Reactions
Heat flows into system
Heat is a reactant
System has more energy
H is positive to indicate direction of heat
flow
Energy absorbed to break bonds> energy
released making bonds
Reactant bond energy > product bond
energy
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 individual steps.
When solving Hess’s Law problems keep the following in mind:
1. If the reaction is reversed, the sign of H must also be
reversed.
2. If the coefficients in a balanced equation are multiplied by
an integer, the value of H must also be multiplied by the
same integer. Likewise, if the coefficients are divided,
then H must also be divided.
Hess’s Law
For example:
CH4(g) + 2O2(g)  CO2(g) +2H2O(g)
ΔH = -802 kJ
2H2O(g)  2H2O(l)
ΔH = -88 kJ
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
ΔH = -890 kJ
Given:
2C(s) + 2H2(g)  C2H4(g)
H = ?
Use these:
C(s) + O2(g)  CO2(g)
C2H4(g) + 3O2(g)  2CO2(g) + 2H2O
H2(g) + ½ O2(g)  H2O(l)
H = -393.5 kJ
-1410.9
-285.6
Diborane, B2H6, is a highly reactive boron hydride, which was
once considered a possible rocket fuel. Calculate the H for
the synthesis of diborane from its elements
using the following information:
H
2B(s) + 3/2O2(g)  B2O3(g)
-1273 kJ
B2H6(g) + 3O2(g)  B2O3(g) + 3H2O(g)
-2035 kJ
H2(g) + 1/2O2(g)  H2O(l)
-286 kJ
H2O(l)  H2O(g)
44 kJ
Standard enthalpy of
formation (Hof)
…the energy associated with the formation of a compound
from its elements in their standard state.
Standard state is 25oC and 1 atm.
Hof for the most stable form of an element = 0
Horxn = nprod Hof (prod) - nreact Hof (reac)
Calculate the change in enthalpy for:
4NH3(g) + 7O2(g)  4NO2(g) + 6H2O(l)
Calorimetry
Method of measuring heat flow between
systems and surroundings.
Heat capacity = the amount of heat
necessary to raise the temperature of an
object by one degree Celsius or Kelvin
The heat capacity of one mole of a
substance is called its molar heat
capacity, joules per mole per degree
Celsius or Kelvin
Specific heat also known as specific heat capacity is
defined as the amount of heat necessary to raise the
temperature of 1 g of a substance by one degree.
Note that heat capacity applies to objects, while
molar and specific heat capacities apply to
substances.
The device used to measure heat flow between
systems and surroundings is called a calorimeter.
Constant Pressure
Calorimetry/Coffee Cup
qp = (m) (Cp) (T)
qp = H assumed:
solutions have the
density of water
solutions have the
specific heat of water
no heat lost to
surroundings
Constant Volume Calorimetry/Bomb Calorimeter
qv = E
If given heat capacity
of calorimeter and
water as one quantity
qv = C x T
If the heat capacity of
the bomb and the
amount of water is
given qv =(mCT)water
+ (C x T)bomb
A 1.5 gram sample of methane gas was burned with an
excess of oxygen in a bomb calorimeter with a heat
capacity of 11.3 kJ/oC, the temperature increase was
14.3oC. Calculate the energy of combustion per gram of
methane.
Combustion of a 0.1584 gram sample of vanillin, C8H8O3,
in a calorimeter containing 987 grams of water with a heat
capacity of 6.66 J/oC, increased the temperature of the
bomb calorimeter 3.25oC. What is the enthalpy of
combustion per mole of vanillin?
The complete combustion of octane to produce gaseous
carbon dioxide and liquid water at standard conditions gives
off 47.9 kJ of heat per gram of octane. Write a chemical
equation to represent this information.
1. What transfer of energy takes place when:
a. a saucepan of water is heated to boiling
b. a tennis ball bounces to rest on the floor
c. two automobiles collide head on
d. a flashlight is left on until the battery runs down
e. a fire burns
f. a wind turbine generates electricity
g. a chemical reaction liberates heat
2. If the temperature of a 50.0-g block of aluminum increases by
10.9 K when heated by 500 J, calculate the
a. heat capacity of the aluminum block
b. molar heat capacity of aluminum
c. specific heat of aluminum
3. If the enthalpy change for the combustion of propane is -2220
kJ/mol propane, what quantity of heat is released when 1 kg
of propane is burned?
If you have 10 moles each of copper and water– one is very
hot and the other is very cold and you simultaneously toss them
into the same insulated container,
a. Which will gain or lose more heat?
b. What will be true of the magnitudes of the ΔT’s?
c. If the copper was originally the hotter object, what will be
the sign of ΔT for each substance?
d. Is this a chemical or physical change?
Consider the process: Z2 (g)  2Z (g) which takes place inside a
rigid, sealed container at constant temperature. Answer the
Following using +, -, or 0.
q=
w=
ΔE =
ΔEuniv=