Chemistry-Ch08 Chemical Thermodynamic

8/9/2019 Chemistry-Ch08 Chemical Thermodynamic

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88ChemicalChemical

thermodynamicsthermodynamics


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8.1 Introduction to chemical8.1 Introduction to chemicalthermodynamicsthermodynamics

Chemical thermodynamics allows us

to predict both the direction and the

extent of spontaneous chemical andphysical change under particular

conditions using a property called the

Gibbs free energy, G.


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8.2 Gibbs free energy8.2 Gibbs free energy

Chemical reactions and physical

changes almost always either absorb

or release energy as heat.

Energy may be distributed throughout

a chemical system in a large number

of different ways, some of which have

significantly higher probabilities than

others


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G =H TS

Henthalpy of the system

Function related to the heat absorbed orevolved by a chemical system

Sentropy of the system

Measure of number of ways energy isdistributed throughout a chemical system

Ttemperature in kelvin


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G =H TS

G allows us to determine whether a

particular chemical reaction or physicalchange is spontaneous

IfG < 0, the process is spontaneous

If

G > 0, the process isnonspontaneous

IfG = 0, the system is at equilibrium


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System refers to the particular

chemical species being studied

Surroundings are everything else



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Universe refers to the system and the

surroundings

Boundary defined as region acrosswhich heat flows



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Open systems

Can gain or lose mass and energy

across their boundaries Closed systems

Can absorb or release energy, but not

mass, across the boundary

Isolated systems

Can not exchange matter or energy with

their surroundings


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SI unit of energy

Joule

1 J = 1 kg m2

s-2

1 kJ = 103 J

Thermodynamic equations require the

temperature in kelvin

Temperature difference (T) of 1 K is

numerically equal to Tof 1 C


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8.3 Enthalpy8.3 Enthalpy

Internal energy (U)

The sum of energies for all of the

individual particles in a sample of matter

Interested in the CHANGE in internal

energy that accompanies a process

U = Ufinal Uinitial

U = Uproducts Ureactants


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Internal energy

Chemical system can exchange energy

with its surroundings in two ways

Either absorbing heat from or emitting

heat to the surroundings

Doing work on the surroundings ofhaving the surroundings do work on it

Work may be defined simply as motion

against an opposing force


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Internal energy

Type of work most

often encountered inchemical systems is

the compression or

expansion of gas

Often called pressure-volume orpVwork

w= pV


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First law of thermodynamic

U = q + w

q - heat w- work

Energy can be transferred between

systems as either heat or work

It can never be created or destroyed

Law of conservation of energy


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State functions

Value is dependent only on the current

state of the system

X = Xfinal Xinitial

Independence from the method or

mechanism by which a change occurs is

the important feature of all statefunctions


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State functions

q and ware NOT state functions

The values ofqand ware

dependent on

the path

of change


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Heat Capacity

Heat and temperature are not the same

thing

Heat is a transfer of energy due to a

temperature difference

q =CT

q - heat (J) C- heat capacity (J K-1)


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Heat Capacity

Heat capacity depends on the size of the

sample

A property with a value that depends on

the size of the sample is an extensive

property

A property with a value independent ofthe size of the sample is an intensive

property


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Heat Capacity

Divide heat capacity (extensive property)

by the mass of the sample to form

specific heat capacity (intensive

property)

c - specific heat capacity (J g-1 K-1)

Divide by amount instead of mass to

form molar heat capacity (J mol-1 K-1)

m

Cc !


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Heat Capacity

q = c T

If a gold ring with a mass of 5.50 g changesin temperature from 25.0 to 28.0 C, how

much heat has it absorbed?

m = 5.50 g c= 0.129 J g1 K1 T= 3 K

q = cmT

= (0.129 J g-1 K-1) (5.50 g) (3 K)

= 2.1 J


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The determination of heat

Calorimeter

Apparatus designed to

minimise heat lossbetween the system

and surroundings

Bomb calorimeter

System remains atconstant volume

U=q + w

U=qv



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Enthalpy: the heat of reaction at

constant pressure

U

= q + wU = qp pV

Inconvenient as need

to know V

Define a new thermodynamic

property called enthalpy (H)

H = qp



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constant pressure

H = U + pV

H = U + pV

Substituting U = qp pVgives

H = qp pV+ pV

H = qp


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constant pressure

The heat of reaction at constant pressureis equal to H

The heat of reaction at constant volume

is equal to U

H> 0 reaction is endothermic H< 0 reaction is exothermic

The difference between Hand U for a

reaction is pV


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Standard enthalpy change

Standard states

Pressure of 105 Pa

Concentration of 1 M

Standard enthalpy of reaction

Value ofHfor a reaction occurring understandard conditions (H(kJ or kJ mol1)

Involves the actual numbers of MOLES

specified by the coefficients of the equation


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Standard enthalpy change

N2(g) + 3H2(g) 2NH3(g)H=92.38 kJ

The above is a thermochemical equation Always gives the physical states of the

reactants and products

Its value ofH is only true when

coefficients of reactants and productsare numerically equal to the number of

moles of the corresponding substances


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Hesss law

Method for combining known

thermochemical equations in a way that

allows us to calculate Hfor another

reaction

One step

C(s) + O2(g)

CO2(g)

H

= 393.5 kJ Two step

Step 1: C(s) + O2(g) CO(g) H= 110.5 kJ

Step 2: CO(g) + O2(g) CO2(g)H= 283.0 kJ


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Hesss law2Fe(s) + 3CO2(g) Fe2O3(s) + 3CO(g) H

= +26.7 kJ

3CO(g) + O2(g) 3CO2(g) H = -849.0 kJ

2Fe(s) + O2(g)

Fe2O3(s) H

= -822.3 kJ Rules for manipulating thermochemical

equations:

1. When an equation is reversed the sign of

H must also be reversed.Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2(g) H

= -26.7 kJ

2Fe(s) + 3CO2(g) Fe2O3(s) + 3CO(g) H = +26.7 kJ



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Rules for manipulating thermochemical

equations:

2. Formulae can be cancelled from both

sides of an equation only if the substanceis an identical physical state.

Hesss law2Fe(s) + 3CO2(g) Fe2O3(s) + 3CO(g) H

= +26.7 kJ

3CO(g) + O2(g) 3CO2(g) H = -849.0 kJ

2Fe(s) + O2(g) Fe

2O

3(s) H = -822.3 kJ



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Rules for manipulating thermochemical

equations:

3. If all the coefficients of an equation are

multiplied or divided by the same factor,the value ofH must likewise be

multiplied or divided by that factor.

2Fe(s) + 3CO2(g) Fe2O3(s) + 3CO(g) H = +26.7 kJ

3CO(g) + O2(g) 3CO2(g) H = -849.0 kJ

2Fe(s) + O2(g) Fe2O3(s) H = -822.3 kJ

CO(g) + O2(g) CO2(g) H = -283.0 kJ

3CO(g) + 3/2O2(g) 3CO2(g) H = -849.0 kJ

Hesss law



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Standard enthalpy of combustion

cH

Enthalpy change at temperature Twhen1 mole of a substance is completely

burned in pure oxygen gas

Combustion reactions are always

exothermiccH

always negative

kJ mol1


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Standard enthalpy of formation (fH)

Enthalpy change when 1 mole of

substance is formed at 105 Pa and the

specified temperature from its elements in

their standard states

An element is in its standard state when it

is in its most stable form and physical stateat 105 Pa and the specified temperature

fH for the elements in their standard

states are 0


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Standard enthalpy of formation

aA + bB cC + dD

Hesss law equation

Use either enthalpies of combustion or enthalpies of

formation

f f f f D A BCH c H d H a H b H

U U U U U ( ! ( ( ( (-


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Bond enthalpies

A bond enthalpy is the enthalpy change on

breaking 1 mole of a particular chemical

bond to give electrically neutral fragments

Atomisation enthalpy (atH) is the enthalpy

change that occurs on rupturing all the

chemical bonds in 1 mole of gaseousmolecules


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Bond enthalpies and Hesss law



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Lattice enthalpies and Hesss law

Lattice enthalpies for ionic solids calculable

using Hesss law and thermodynamic data


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Entropy and probability

Spontaneous processes tend to proceed from states of

low probability to states of higher probability.

Spontaneous processes tend to disperse energy

8.4 Entropy8.4 Entropy

No energy is transferred

One unit of energy is

transferred

Two units of energy

are transferred

Three units of energy

are transferred


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Entropy and entropy change

Entropy (S) describes the number of

equivalent ways that energy can be

distributed in the system

Entropy is a state function

S = Sproducts Sreactants

Any event that is accompanied by an

increase in the entropy of the system has

a tendency to occur spontaneously



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Entropy and entropy change



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Factors that affect entropy

Often possible to predict whetherSis positive

or negative for a particular change

Volume

For gases the entropy increases with increasing

volume



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Temperature

The higher the temperature the higher the entropy



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Physical state

One of the major factors that affects the entropy of a

system is its physical state



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Number of particles

When all other things are equal, reactions that

increase the number of particles in the system tendto have a positive entropy change



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The second law of thermodynamics

Whenever a spontaneous event takes

place in our universe, the total entropy of

the universe increases (Stotal>0)

Stotal= Ssystem + Ssurroundings

qsurroundings = -qsystem

qsystem =H

suroundingssurroundings

qS

T( !

systemsurr undin s

HS

T

(( !


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Absolute entropy and the third law of

thermodynamics

At absolute zero, the entropy of a perfectly

ordered pure crystalline substance is 0

S= 0 at T= 0 K

Point at which entropy equal to 0 is known,hence by experimental measurement and

calculation entropy can be determined for

a substance at temperatures above 0 K



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Absolute entropy and the third law of

thermodynamics

Standard entropy (S) entropy of

1 mole of a substance determined under

standard conditions at temperature

of 298 K

Standard entropy of reactionaA + bB cC + dD

f f f f D A BCS c S d S a S b S U U U U U ( ! ( ( ( (-


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The sign ofG

G = H TS

G < 0 reaction

spontaneous

8.5 Gibbs free energy and8.5 Gibbs free energy andreaction spontaneityreaction spontaneity


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Standard Gibbs free energy change

When G is determined at 105 Pa, this is

called the standard free energy change

(G)

There are several ways for determining

G for a reaction

G

=

H

T

S

aA + bB cC + dD

f f f f D A BCG c G d G a G b GU U U U U ( ! ( ( ( (-


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Gibbs free energy and work

Maximum conversion of chemical energy to work

occurs if a reaction is carried out under conditions

that are said to be thermodynamically reversible



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The maximum amount of energy produced

by a reaction that can be

THEORETICALLY harnessed as work is

equal to G

G = 0 the system is in a state of

equilibrium

Gproducts =Greactants

When G = 0 the amount of work

available is 0



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For phase changes such as

H2

O(l) H2

O(s)

equilibrium can be established only at

one particular temperature at

atmospheric pressure

G = 0 =H TSH = TS

HS

T

(( !

HT

S

(!

(


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Gibbs free

energy and

work


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