IB1 Chemistry HL Energetics - PowerPoint Presentation

Slide1

IB1 ChemistryHL Energetics

Why do chemical reactions happen?

Slide2

Topic 15: Energetics (8

hours)

15.1

Standard enthalpy changes of reaction

15.1.1

Define and apply the terms standard state, standard enthalpy change of formation (¬H ) f

Ö and standard enthalpy change of combustion (¬H ) c Ö .

15.1.2 Determine the enthalpy change of a reaction using standard enthalpy changes of formation and combustion.

15.2 Born–Haber cycle

15.2.1

Define and apply the terms lattice enthalpy and electron affinity.

15.2.2 Explain how the relative sizes and the charges of ions affect the lattice enthalpies of different ionic compounds.

15.2.3 Construct a Born–Haber cycle for group 1 and 2 oxides and chlorides, and use it to calculate an enthalpy change.

15.2.4 Discuss the difference between theoretical and experimental lattice enthalpy values of ionic compounds in terms of their covalent character.

15.3 Entropy

15.3.1

State and explain the factors that increase the entropy in a system.

15.3.2 Predict whether the entropy change (

ΔS

) for a given reaction or process is positive or negative.

15.3.3 Calculate the standard entropy change for a reaction (¬S Ö ) using standard entropy values (S Ö ) .

15.4 Spontaneity

2.5 hours

15.4.1 Predict whether a reaction or process will be spontaneous by using the sign of ¬

GÖ

.

15.4.2 Calculate ¬

GÖ

for a reaction using the equation

¬

GÖ

= ¬H Ö−

T¬S

Ö and by using values of the standard free energy change of formation,

ΔGf

Ö .

15.4.3 Predict the effect of a change in temperature on the spontaneity of a reaction using standard entropy and enthalpy changes and the equation

¬

GÖ

= ¬H Ö−

T¬S

Ö .

Slide3

Standard enthalpy change of reaction

in standard state at

101kPa

, 298K

(PV =

nRT

)

Slide4

Standard enthalpy of combustion: DHc

q

When

a substance is fully

burned

in

oxygen

for example

CH

4

+ 2O

2



CO

2

+ 2H

2

O

Slide5

Standard enthalpy of formation: DHf

q

D

H

f

q

: The energy absorbed or evolved when 1 mol of the substance is formed from its elements in their standard states. The enthalpy of formation of any element is zero.

H

2

(g

)

+

½O

2

(g

)



H

2

O

(l)

D

H

f

q

= -285

kJ/

mol

D

H =

SD

H

f

(products)

-

SD

H

f

(reactants)

Slide6

Ionic compound consist of ions arranged in a lattice

Slide7

Two or more electrons can be transferred

Different sized atoms give different mineral structures as they pack in a different way

Hexagonal Beryl crystal; Image Wikipedia

Slide8

Lattice enthalpy, DHlattice

R

elates

to the endothermic process

MX(s

)



M

+

(g)

+ X

-

(g)

in

which the gaseous ions of a crystal are separated to an infinite distance from each other.

NaCl

(s)



Na

+

(g)

+

Cl

-

(g)

D

H

lattice

= 771kJ/mol Endothermic reactions.

Slide9

Lattice enthalpy depends oncharge on an ion

size of an ion

packing arrangement

MgO

-

3791kJ

/

mol

NaCl

-

790kJ

/

mol

Slide10

Standard enthalpy of atomization: DH

at

q

D

H

at

q

: The energy

required to atomize one mole of an element.

Na

(s)



Na

(g)

+ e

-

D

H

at

q

=

+108 kJ/

mol

(physics: related to latent heat of vaporization)

Slide11

Electron affinityThe

enthalpy change

per mole when

an atom

gains

one electron in

the gaseous phase

Cl

(g

)

+ e

-

(g)

 Cl-(g)

DH

ea

= -351 kJ/mol

.

Electron

affinity can be both exothermic and endothermic depending on element.

Slide12

How does the electron affinity change across a period?

Slide13

Born-Haber cycle

http://

en.wikipedia.org

/wiki/

File:Born-haber_cycle_LiF.svg

Slide14

Born-Haber cycles used to calculate lattice enthalpies

Can be used to find out if a bond is more or less ionic

Slide15

Born Haber cycle for sodium chlorideEnthalpy of formation of NaCl = -411kJ/mol

Enthalpy

of atomisation of Na = +103 kJ/mol

Enthalpy

of atomisation of

Cl

= +121

kJ/

mol

(½ energy

of

Cl-Cl

bond

)Electron

affinity of Cl = -364 kJ/molIonisation energy of Na = + 500 kJ/mol

Slide16

Lattice enthalpy for sodium chlorideEnergies of atomisation + Electron affinity + Ionisation energy =

= Enthalpy of formation + Lattice enthalpy

Lattice energy = 771 kJ/mol

Slide17

How ionic is the ionic lattice?

Chemistry

Data Booklet

gives

lattice enthalpies

as

:

Experimental values (obtained

from

Born-Haber cycle)

Theoretical values (calculated

using electrostatic calculations)

Greater difference between theoretical and experimental values



more covalent character of the bond

.

Slide18

Decomposition of ammonium nitrate

NH

4

NO

3

(s

)



N

2

O

(g)

+ 2 H

2O(l)

 NH4NO

3(s

)

D

H

f

q

= -366 kJ/mol

D

S

q

= 151 J/K*mol

N

2

O

(g)

D

H

f

q

= +82 kJ/mol

D

S

q

= 220 J/K*mol

H

2

O

(l)

D

H

f

q

= -285 kJ/mol

D

S

q

= 70 J/K*mol

D

H

= [

D

H

f

(N2O(g))

+

D

H

f

(H2O(l))

] – [

D

H

f

(NH4NO3(g)

] =

=[

82 + 2(-285)] - [-366] = -122 kJ/mol

Slide19

Entropy: disorder

Slide20

Which is more likely if the particles are in constant random motion?

Slide21

EntropyEntropy, S = Disorder Unit:

JK

-

1

mol

-1

D

S =

change

in

disorder

D

S =

S

p -

SrAbsolute value of

S can be measured

Slide22

Entropy and changes of stae

Solid

Liquid

Gas

H2O

Ice

Water

Steam

JK

-

1

mol

-1

48.0

69.9

188.7

Increasing entropy



Slide23

Will a reaction happen? Spontaneity

Slide24

All reactions involve changes in H and S

D

S

is probably positive

if moles

of gas

increase

and

moles

of

solid or liquid

decrease

.

NH4Cl(s)



NH

3(g)

+ HCl

(g)

D

S = +

285JK

-

1

mol

-1

Pb

2

+

(

aq

)

+ 2 I

-



PbI

2(s)

D

S = - 70

JK

-

1

mol

-1

Slide25

Spontaneity of a reaction

Nature likes low internal energy (

D

H to decrease) and high disorder (

D

S to increase)

A

reaction will occur if the final state is more

probable

than the initial state.



Decrease

in

DH



Increase

in

D

S

Slide26

Gibbs free energy, DGq

=

D

H

q

-

T

D

S

q

Temperature dependent

Spontaneous:

D

G negative (

D

G

q

<

0)

Equilibrium:

D

G

q

= 0

Not spontaneous

:

D

G

q

positive

(

D

G

q

>

0)

Slide27

D

H

D

S

D

G

Spontaneity

Negative

(Exothermic)

Positive

(More random)

D

G

< 0

Always negative

Always

spontaneous

Positive (Endothermic)

Negative

(More order)

D

G

>

0

Always positive

Never

spontaneous

Negative

(Exothermic)

Negative

(More order)

Depends on T

Spontaneous at low

Temp

Positive (Endothermic)

Positive

(More random)

Depends on T

Spontaneous at high

Temp

Slide28

Activation energy is also importantJust the fact that a reaction is spontaneous doesn’t mean that it will occur at once.

It also depends on activation energy. And we will deal with that later on in topic 7.

Slide29

LinksIonic bonding

http://www.teachersdomain.org/asset/lsps07_int_ionicbonding/

Covalent bonding

http://www.teachersdomain.org/asset/lsps07_int_covalentbond/