PRINCIPLES OF CHEMISTRY II CHEM 1212

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Oct 27, 2013 (4 years and 2 months ago)

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PRINCIPLES OF CHEMISTRY
II


CHEM 1212



CHAPTER
17



DR. AUGUSTINE OFORI AGYEMAN

Assistant professor of chemistry

Department of natural sciences

Clayton state university

CHAPTER 17


CHEMICAL THERMODYNAMICS

Used to


-

Describe heat transfer in chemical systems


-

Evaluate the heat evolved or absorbed by a reaction


-

Predict the maximum energy that can be produced by a reaction


-

Predict whether a proposed reaction is feasible


-

Predict whether a process is spontaneous or not

CHEMICAL THERMODYNAMICS

Spontaneous Process

-

Takes place with no apparent cause

Nonspontaneous Process

-

Requires something to be applied in order for it to occur

(usually in the form of energy)

CHEMICAL THERMODYNAMICS

-

The ability to do work or to transfer heat



-

Energy is necessary for life: humans, plants, animals, cars



-

Forms of energy are interconvertible


-

Units: kg

m
2
/s
2

or Joules (J)

ENERGY

SYSTEM AND SURROUNDINGS

System

-

The limited and well
-
defined portion of the

universe under study



Surroundings


-

Everything else in the universe



Studying energy changes in a chemical reaction

-

The reactants and products make up the system

-

The reaction container and everything else make up

the surroundings

SYSTEM AND SURROUNDINGS

Open System

-

Matter and energy can be exchanged with the surroundings


(water boiling on a stove without a lid)


Closed System

-

Energy but not matter can be exchanged with the surroundings


(two reactants in a closed cylinder reacting to produce energy)


Isolated System

-

Neither matter nor energy can be exchanged with the surroundings


(insulated flask containing hot tea)

-

The energy transferred when a force moves an object



-

The product of
force (F)

and
distance (d)

through which

the object moves


w = F x d


-

Units: kg

m
2
/s
2

or Joules (J)

WORK (w)

-

The energy transferred between a system and its surroundings


due to difference in temperature



-

A form of energy necessary to raise the temperature


of a substance



-

Units: kg

m
2
/s
2

or Joules (J)

HEAT (q)

-

Sum of all potential and kinetic energies of all components


-

Change in internal energy = final energy minus initial energy



E

= E
final

-

E
initial



-

Energy can neither be created nor destroyed


-

Energy is conserved

INTERNAL ENERGY (
E)


E

= E
final

-

E
initial




If E
final

> E
initial


E is positive and system has gained energy from its surroundings



If E
final

<
E
initial


E is negative and system has lost energy to its surroundings

INTERNAL ENERGY


E = q + w

q = heat added to or liberated from a system

w = work done on or by a system


Internal energy of a system increases when


-

Heat is added to the system from surroundings (positive q)

-

Work is done on the system by surroundings (positive w)

INTERNAL ENERGY

+w

+q

system


E = q + w

q = heat added to or liberated from a system

w = work done on or by a system


Internal energy of a system decreases when


-

Heat is lost by the system to the surroundings (negative q)


-

Work is done by the system on the surroundings (negative w)

INTERNAL ENERGY

-
w

-
q

system

Endothermic Process

-

Process in which system absorbs heat (endo
-

means ‘into’)

-

Heat flows into system from its surroundings


(melting of ice
-

the reason why it feels cold)



Exothermic Process

-

Process in which system loses heat

-

Heat flows out of the system (exo
-

means ‘out of’)


(combustion of gasoline)

INTERNAL ENERGY

State Function



-

Property that depends on initial and final states of the system


-

Does not depend on path or how a change occurs


-

Internal Energy depends on initial and final states


-

Internal Energy is a state function


-

q and w, on the other hand, are not state functions

INTERNAL ENERGY

Internal energy is influenced by


-

Temperature


-

Pressure


-

Total quantity of matter




-

Internal Energy is an extensive property

INTERNAL ENERGY

INTERNAL ENERGY

Calculate

E

for a system absorbing 22 kJ of heat from its

surroundings while doing 11 kJ of work on the surroundings.

State whether it is an endothermic or an exothermic process


q = +22 kJ (heat is added to the system from surroundings)

w =
-
11 kJ (work is done by the system on the surroundings)



E = q + w


E = (+ 22


11) kJ = +11 kJ

Endothermic

PRESSURE
-
VOLUME WORK

A gas against constant pressure

w =
-

P

V


w = work and P = pressure



V = change in volume = V
final



V
initial


w is positive when the gas contracts (negative

V)


w is negative when the gas expands (positive

V)


Units:
L∙atm (1 L∙atm = 101.3 J)

Calculate the work associated with the expansion of a gas from

32 L to 58 L at a constant pressure of 12 atm



w =
-

P

V


w =
-

(12 atm)(58 L
-

32 L) =
-

310 L∙atm


Gas expands hence work is done by system on surroundings

PRESSURE
-
VOLUME WORK

Calculate the work associated with the compression of a gas from

58 L to 32 L at a constant pressure of 12 atm



w =
-

P

V


w =
-

(12 atm)(32 L
-

58 L) = + 310 L∙atm


Gas contracts hence work is done on system by surroundings

PRESSURE
-
VOLUME WORK

Expansion of Volume

-


V is a positive quantity and w is a negative quantity

-

Energy leaves the system as work

-

Work is done by the system on the surroundings



Compression of Volume

-


V is a negative quantity and w is a positive quantity

-

Energy enters the system as work

-

Work is done on the system by the surroundings

PRESSURE
-
VOLUME (P
-
V) WORK

THE FIRST LAW OF THERMODYNAMICS

-

The law of conservation of energy


-

Energy can be neither created nor destroyed


-

Concerned with change in energy


In an isolated system

-

Neither matter nor energy can enter or leave


E = 0


In a closed system

-

Energy can enter or leave in the form of heat and work


E = q + w

ENTHALPY (H)

-

Heat flow in processes occurring at constant pressure


-

Only P
-
V work are performed


H = E + PV


H, E, P, and V are all state functions



Change in Enthalpy



H =

(E + PV)

ENTHALPY (H)

Change in Enthalpy

at Constant Pressure



H =

E + P

V



E = q + w


P

V =
-

w


Implies


H = (q
p

+ w)
-

w = q
p


q
p

= heat at constant pressure

ENTHALPY (H)


H = q
p


Change in enthalpy = heat gained or lost at constant pressure


Positive

H


-

System gains heat from the surroundings

-

Endothermic process


Negative

H


-

System releases heat to the surroundings

-

Exothermic process

ENTHALPY (H)


H =

E + P

V


For reactions involving solids and liquids


V ≈ 0


H ≈

E


For gases


n = n
final



n
initial



n = total gas moles of products


total gas moles of reactants


Implies

P

V = (

n)RT


H =

E + (

n)RT (R = 8.314 J/mol∙K)

ENTHALPY (H)


E = q + w


At constant volume


V = 0

w =
-

P

V = 0

and


E = q
v


q
v

= heat gained or lost at constant volume

ENTROPY (S)

-

The amount of disorder in a process


-

Is a measure of randomness


-

Many spontaneous reactions are accompanied by release of
energy (exothermic processes)


-

Some endothermic processes, however, are spontaneous

(dissolution of some salts such as barium hydroxide)


-

Disorder plays an important role in predicting the
spontaneity of a reaction

ENTROPY (S)

-

Entropy is a state function



Change in entropy


ΔS = S
final



S
initial

ENTROPY (S)

Some General Concepts


-

The entropy of a substance increases as the substance
changes from solid to liquid to gas


-

Due to increase in randomness of the molecules


S
solid

< S
liquid

< S
gas


Generally

ΔS from liquid to gas > ΔS from solid to liquid

ENTROPY (S)

Some General Concepts


-

The entropy of a substance increases when a molecular
solid or liquid dissolves in a solvent

Solid → Aqueous


-

Increase in disorder when a solute dissolves in a solvent


-

Increase in disorder of the solute is greater than the
increase in disorder of the solvent


-

Net decrease in entropy occurs with some ionic solids

(solids with highly charged ions due to hydration)

ENTROPY (S)

Some General Concepts


-

The entropy decreases when a gas dissolves in a solvent

CO
2
(g) → CO
2
(aq)


-

Solute molecules go from the gas phase to the liquid phase


-

Solute molecules become less random


-

Increase in disorder of the solvent is small


-

Hence a net decrease in disorder

ENTROPY (S)

Some General Concepts


-

Entropy increases with increase in temperature



-

Kinetic energy of particles increase as temperature
increases



-

Disorder increases as kinetic energy (energy of motion)
increases

ENTROPY (S)

Other Concepts


-

Entropy increases in a chemical reaction when ∆n is
positive


-

Entropy decreases in a chemical reaction when ∆n is
negative


-

Entropy of a system increases with increasing volume


-

Entropy of a solution increases with dilution


-

Osmosis (spontaneous process) is driven by positive
Δ
S

-

For a spontaneous process there is always an increase in
the entropy of the universe


Δ
S
univ

=
Δ
S
sys

+
Δ
S
surr



Δ
S
univ

> 0 for a spontaneous process


If
Δ
S
sys

is negative,
Δ
S
surr

must be large and positive to make

Δ
S
univ

positive


Example

-

Combustion of hydrocarbons

Δ
S
sys

is negative but
Δ
S
univ

is positive

THE SECOND LAW OF THERMODYNAMICS

If
Δ
S < 0


-

The reverse reaction is spontaneous




If
Δ
S = 0


-

The system is at equilibrium

-

The process is not spontaneous in either direction

-

The process has no tendency to occur

THE SECOND LAW OF THERMODYNAMICS

-

The entropy of any pure crystalline substance at a
temperature of 0 K is zero



-

Absolute value of S can be measured



-

As the temperature of a substance is increased from 0 K,
the motion of the particles increases and entropy increases

THE THIRD LAW OF THERMODYNAMICS

-

Change in entropy is proportional to the added energy


-

The transfer of a given amount of energy as heat has
greater impact on entropy at lower temperatures than at
higher temperatures



-

Δ
S is directly proportional to quantity of heat transferred
(q) and inversely proportional to temperature (T)


Units: J/K

THE THIRD LAW OF THERMODYNAMICS

-

The entropy change of a chemical reaction



∆S
rxn

=
Σ
nS
o
[products]
-

Σ
mS
o
[reactants]


S
o

= standard molar entropy of a substance (at 298 K)


-

n and m are the number of moles of products and reactants


-

The standard molar entropy of an element in its standard
state is not zero

THE THIRD LAW OF THERMODYNAMICS

Calculate ∆S
o

at 25
o
C for the reaction

2NiS(s) + 3O
2
(g) → 2SO
2
(g) + 2NiO(s)


Obtain S
o

values from appendix


∆S
o


= [(2 mol)(248.11 J/mol∙K) + (2 mol)(37.99 J/mol∙K)]




[(2 mol)(52.99 J/mol∙K) + (3 mol)(205.03 J/mol∙K)]


=
-
148.87 J/K

(∆S
o

is negative as ∆n is negative)

THE THIRD LAW OF THERMODYNAMICS

G = H


TS


-

Change in Gibbs free energy at constant temperature and
pressure


∆G = ∆H


T∆S


-

The absolute value of G cannot be measured but ∆G can
be measured


-

∆G is a state function

GIBBS FREE ENERGY (G)

If ∆G < 0

-

Forward reaction is spontaneous



If ∆G = 0

-

System is at equilibrium



If ∆G > 0

-

Reverse reaction is spontaneous

GIBBS FREE ENERGY (G)

Standard Gibbs Free energy of formation (∆G
f
o
)


-

The Gibbs free energy change during the formation of one mole
of a substance in its standard state from its constituent elements
in their standard states



∆G
f
o

= ∆H
f
o



T∆S
f
o

GIBBS FREE ENERGY (G)

Change in Gibbs free energy of a chemical reaction


∆G
o
rxn

=
Σ
n∆G
f
o
[products]
-

Σ
m∆G
f
o
[reactants]



-

n and m are the number of moles of products and reactants



-

The standard Gibbs free energy of an element in its standard
state is zero

GIBBS FREE ENERGY (G)

Influence of Temperature


∆G = ∆H
-

T∆S


From the equation above


-

Decrease in ∆H (more negative) favor spontaneous change


-

Increase in ∆S (more positive) favor spontaneous change


-

∆G is strongly influenced by temperature through the T∆S term

GIBBS FREE ENERGY (G)

Influence of Temperature


At Low Temperatures

-

The sign of ∆H determines the sign of ∆G


At High Temperatures

-

The sign of ∆S determines the sign of ∆G


-

When ∆H and ∆S have opposite signs temperature change does
not influence the direction of spontaneity


-

It is assumed that the numerical values of ∆H and ∆S are not
affected by temperature change

GIBBS FREE ENERGY (G)

Equilibrium Constant


∆G = ∆G
o

+ RTlnQ


∆G = Gibbs free
-
energy change at non
-
standard
-
state conditons


∆G
o

= standard Gibbs free
-
energy change


R = ideal gas constant (8.314 J/mol∙K)


T = temperature (K)


Q = reaction quotient

GIBBS FREE ENERGY (G)

Equilibrium Constant


∆G = ∆G
o

+ RTlnQ


At equilibrium

∆G = 0 and Q = K
eq


Implies

∆G
o

=
-

RTlnK
eq


or

GIBBS FREE ENERGY (G)

In Summary


-

Reaction is spontaneous in the forward direction if ∆G is negative


-

Reaction is spontaneous in the reverse direction if ∆G is positive


-

Reaction is spontaneous in the forward direction if Q < K
eq


-

Reaction is spontaneous in the reverse direction if Q > K
eq


-

At equilibrium ∆G approaches zero and Q approaches K
eq

GIBBS FREE ENERGY (G)

Temperature and Equilibrium Constant


Combining

∆G
o

=
-

RTlnK
eq

and

∆G
o

= ∆H
o

-

T∆S
o


We obtain

∆H
o

-

T∆S
o

=
-

RTlnK
eq


and

GIBBS FREE ENERGY (G)

Temperature and Equilibrium Constant


-

Equilibrium constant changes with temperature


-

If K
1

and K
2

are the equilibrium constant values at temperatures T
1

and T
2
, respectively

GIBBS FREE ENERGY (G)

-

Known as the Clausius
-
Clapeyron equation

Useful Work


-

For a spontaneous process at constant temperature and pressure


-

The maximum useful work that can be performed equals the
change in Gibbs free energy


w
max

= ∆G


-

The energy that is free to perform useful work


-

The equation gives the minimum amount of work required to
cause a change when ∆G is positive

GIBBS FREE ENERGY (G)