New Way Chemistry for Hong Kong A-Level Book 11 Chapter 6 Energetics.

New Way Chemistry for Hong Kong A-Level Book 11

Chapter 6Chapter 6

EnergeticsEnergetics


Standard Enthalpy Changes

As enthalpy changes depend on temperature and pressure. In order to standardize the data recorded, it is necessary to define the standard conditions:

1. elements or compounds in their normal physical states;2. a pressure of 1 atm (101325 Nm-2); and3. a temperature of 250C (298 K)4. concentration of solution = 1 mol dm-3

1. elements or compounds in their normal physical states;2. a pressure of 1 atm (101325 Nm-2); and3. a temperature of 250C (298 K)4. concentration of solution = 1 mol dm-3

6.3 Standard Enthalpy Changes (SB p.141)

Enthalpy change under standard conditions denoted by symbol: H

ø


Standard Enthalpy Change of Neutralization


The standard enthalpy change of neutralization (Hneut) is the enthalpy change when one mole of water is formed from the neutralization of an acid by an alkali under standard conditions.

The standard enthalpy change of neutralization (Hneut) is the enthalpy change when one mole of water is formed from the neutralization of an acid by an alkali under standard conditions.

ø

H+(aq) + OH-(aq) H2O(l) Hneut = -57.3 kJ mol-1


Standard Enthalpy Changes of neutralization

NH3(aq) + H2O(l) === NH4+(aq) + OH-(aq)


-57.1

-57.3

-52.2

NaOH

KOH

NH3

HCl

HNO3

HCl

Hneu AlkaliAcid ø

The value is smaller if weak acids/alkalis are used because some energy is used for the complete dissociation/ionization of the weak acids/alkalis.


Standard Enthalpy Change of Formation


The standard enthalpy change of formation (Hf) is the enthalpy change of the reaction when one mole of the compound in its standard state is formed from its constituent elements under standard conditions.

The standard enthalpy change of formation (Hf) is the enthalpy change of the reaction when one mole of the compound in its standard state is formed from its constituent elements under standard conditions.

ø

Na(s) + ½Cl2(g) NaCl(s) Hf = -411 kJ mol-1

1 mole

ø


Standard Enthalpy Change of Combustion


The standard enthalpy change of combustion (Hc) of a substance is the enthalpy change when one mole of the substance is burnt completely in oxygen under standard conditions.

The standard enthalpy change of combustion (Hc) of a substance is the enthalpy change when one mole of the substance is burnt completely in oxygen under standard conditions.

ø

e.g. C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(l)

Hc = -2220 kJ mol-1ø


Experimental Determination of Enthalpy Changes

Calorimeter = a container used for measuring the temperature change of solution

6.4 Experimental Determination of Enthalpy Changes by Calorimetry (p.148)


Hess’s Law

A + B C + DRoute 1

H1

E

H2 H3

Route 2 H1 = H2 + H3 H1 = H2 + H3

Hess’s Law states that the total enthalpy change accompanying a chemical reaction is independent of the route by which the chemical reaction takes place.

Hess’s Law states that the total enthalpy change accompanying a chemical reaction is independent of the route by which the chemical reaction takes place.

Why?

6.5 Hess’s Law (p. 153)

Conservation of energy


Importance of Hess’s Law

• the reactions cannot be performed in the laboratory

• the reaction rates are too slow• the reactions may involve the formation of side

products

The enthalpy change of some chemical reactions cannot be determined directly from experiment because:

But the enthalpy change of such reactions can be determined indirectly by applying Hess’s Law.

6.5 Hess’s Law (p. 155)

Q. 5


Energetics of Formation of Ionic Compound

Na(s) + ½Cl2(g) NaCl(s)

Hf = formation of ionic bond?

7.2 Energetics of Formation of Ionic Compounds (SB p. 189)

Hf

ø

Na(g) Cl(g)

Na+(g) Cl-(g)

Hf = -411 kJ mol-1


The enthalpy change when one mole of gaseous atoms is formed from its elements in the defined physical state under standard conditions.

Questions: Why are the changes endothermic?

What type of bond is broken in each case?

7.2 Some important enthalpy terms (SB p. 190)

Na(s) Na(g) H atom [Na(s)] = +109 kJ mol-1

ø1/2 Cl2(g) Cl(g) H atom [1/2Cl2(g)] = +121 kJ mol-1ø

Standard Enthalpy Change of Atomization (H atom)ø


The amount of energy required to remove one mole of electrons from one mole of atoms or ions in the gaseous state.

Questions:Why are the changes endothermic?

Energy is needed to overcome the attractive force between the positive nucleus and the negatively charged electrons.


Ionization Enthalpy (H I.E.)

Na(g) Na+(g) + e- H I.E [Na(g)] = +494 kJ mol-1øMg(g) Mg+(g) + e- H I.E [Mg(g)] = +736 kJ mol-1

ø

Mg+(g) Mg2+(g) + e- H I.E [Mg +(g)] = +1 450 kJ mol-1

ø


The energy change when one mole of electrons is added to one mole of atoms or ions in the gaseous state.

Questions: Why does the 1st E.A. of O is negative while the second one is positive?


Electron Affinity (H E.A.)

ø

First electron affinity

O(g) + e- O-(g) H E.A [O(g)] = - 142 kJ mol-1

Second electron affinity

O-(g) + e- O2-(g) H E.A [O-(g)] = + 844 kJ mol-1

øø


The energy change when one mole of an ionic crystal is formed from its constituent ions in the gaseous state under standard conditions


Lattice Enthalpy (H L.E.)

ø

Na+ (g) + Cl-(g) NaCl(s) H lattice [Na+Cl-(s)]ø

It is a measure of the strength of ionic bond.

+ –

– +


L.E. can be determined indirectly by either:(1) calculations basing on the knowledge of electrostatics in Physics (assuming ions are point charges) – Theoretical

Value or

(2) calculations basing on Hess’s Law --- Experimental Value

L.E. can be determined indirectly by either:(1) calculations basing on the knowledge of electrostatics in Physics (assuming ions are point charges) – Theoretical

Value or

(2) calculations basing on Hess’s Law --- Experimental Value


Na+ (g) + Cl-(g) NaCl(s) H lattice [Na+Cl-(s)]

ø

Q. 8


Born-Haber Cycle for the formation of sodium chloride


Hatom[Na(s)]

HI.E.

Do Q. 6(a) on p. 181Do. Q. 49 on p. 325Do Q. 9(c) i. on p. 316


Strength of ionic lattice7.2 Strength of Ionic lattice (SB p. 196)

1. Ionic SizeGreater the size lower charge density weaker attraction lower (less negative) lattice enthalpy

NaCl: -771 kJ mol-1; KCl: -701 kJ mol-1

2. Ionic ChargeHigher charge stronger attraction higher lattice enthalpy

CaO: -3513 kJ mol-1; CaCl2: -2237 kJ mol-1


Discrepancy between calculated and experimental value

7.2 Strength of Ionic lattice (not typically mentioned in book)

Assumption in calculating lattice enthalpy:. Ions are hard sphere. Completely transfer of electrons. Charge density distributes evenly on the sphere

compound calculated experimental difference

NaCl -770 -780 10

KCl -702 -711 9

AgCl -833 -905 72

AgI -778 -889 111

Incomplete transfer of electrons covalent characters


Bond Enthalpy (for covalent bond) is the energy associated with a chemical bond.

When a chemical bond is broken, a certain amount of energy is absorbed.

When a chemical bond is formed, a certain amount of energy is released.

Bond Enthalpies – Covalent Compounds

8.3 Bond Enthalpies (SB p.220)


8.3 Bond Enthalpies (SB p.221)

Bond Dissociation EnthalpiesB.D.E of a certain bond is the amount of energy required to break one mole of that bond in a particular compound under standard conditions.

Bond EnthalpiesAverage bond enthalpy (or simply bond enthalpy) is the average of the bond dissociation enthalpies required to break a particular chemical bond.

Q. 12 b, Q. 13


Bond length (for covalent bond)Bond length (for covalent bond)

8.6 Relationship between Bond Enthalpies and Bond Lengths (SB p.228)

Bond Lengths


Any conclusion for the relationship between bond length & bond enthalpy?

Any conclusion for the relationship between bond length & bond enthalpy?

Usually a longer bond length corresponds to a lower value of bond enthalpy (weaker bond)

Usually a longer bond length corresponds to a lower value of bond enthalpy (weaker bond)


Bond Bond length (nm)

Bond enthalpy

(kJ mol-1)

H-H

Cl-Cl

Br-Br

I-I

H-F

H-Cl

H-Br

H-I

0.074

0.199

0.228

0.266

0.092

0.127

0.141

0.161

436

242

193

151

565

431

364

299

Bond enthalpies and bond lengths

Distance between shared electrons pair and nuclei increases attraction decreases bond strength decreases

Distance between shared electrons pair and nuclei increases attraction decreases bond strength decreases


(often referred as ‘Atomic radius’ ???)

The space occupied by an atom in a covalently bonded molecule in the direction of the covalent bond (generally taken as half of the bond length)

The space occupied by an atom in a covalently bonded molecule in the direction of the covalent bond (generally taken as half of the bond length)


Covalent Radius



Bond Calculated bond length (nm)

Experimentally determined bond length (nm)

C-O

C-F

C-Cl

C-Br

C-C

H-Cl

C-H

N-Cl

0.150

0.149

0.176

0.191

0.154

0.136

0.114

0.173

0.143

0.138

0.177

0.193

0.154

0.128

0.109

0.174

Calculated and experimentally determined bond length

Difference in electronegativities polar bond ionic character


Standard Enthalpy Change of Solution -- dissolving

Note that enthalpy changes of solution associate with physical changes.

6.3 Standard Enthalpy Changes (SB p.146, NB p. 33))

The standard enthalpy change of solution (Hsoln) is the enthalpy change when one mole of a solute is completely dissolved in a sufficient large volume of solvent to form an infinitely dilute solution under standard conditions.

The standard enthalpy change of solution (Hsoln) is the enthalpy change when one mole of a solute is completely dissolved in a sufficient large volume of solvent to form an infinitely dilute solution under standard conditions.

ø

NaCl(s) + water Na+(aq)+Cl-(aq) Hsoln=+3.9 kJ mol-1ø

LiCl(s) + water Li+(aq) + Cl-(aq) Hsoln=-37.2 kJ mol-1ø


According to Hess’s law,

The enthalpy change of solution is :

Hsoln = Hhyd – Hlattice


This change involves 2 processes:

1st : NaCl(s) Na+(g) + Cl–(g) H = +776 kJ mol–1

The enthalpy change involved in this process is the reverse

of lattice enthalpy. The lattice enthalpy is –776 kJ mol–1.

2nd : hydration enthalpy

Na+(g) + Cl–(g) Na+(aq) + Cl–(aq)

Hhyd = –772 kJ mol–1

NaCl(s) Na+(aq) + Cl–(aq) Hsoln = +776 – 772

= +4 kJ mol–1


Effect of charge and size of ions on Hhyd and Hlattice

rr

ZZlatticeH

r

Z

r

ZhydH

6.3 Standard Enthalpy Changes (NB p. 34)


Generalizations (Not group trend):

When there is a mismatch in ionic size, the salt is expected

to be fairly soluble.

Reasons:

i. The magnitude of lattice enthalpy is not great because of the

large ion, making (r++ r-) relatively large.

ii. The magnitude of the hydration enthalpy is still large due to

the presence of the small counter ion.

Fast Prediction of solubility (NB. P.38)


Relative Solubility of the alkalis metal halides

Group trend of solubility (NB. P.36)

LiI

NaI

KI

RbI

CsI

rr

ZZlatticeH

r

z

r

zhydH

less negativeweaker bondmore soluble

less negativesmaller attraction to waterless soluble


• For Group I iodide, cations are much smaller than anions

• The Hlattice is determined by the reciprocal of the sum of ca

tionic and anionic radii (i.e. )

Large anionic radius makes the relatively small catio

nic radius insignificant w.r.t the sum of r+ and r–

r+ + r- ~ r-

Down the group, increase in cationic size does not m

ake a significant decrease in the magnitude of Hlattice

rr

1

Solubility of Group I metal iodides – group trend



• The hydration enthalpy is determined by

• An increase in cationic size causes Hhyd to become less a

nd less negative significantly

Hsoln becomes less and less exothermic

The solubility of Group I metal iodides decreases do

wn the group


rr

11Hhyd

rH

rr

hyd

1

11


How about the solubility trend of Group I metal Fluorides?


For small anion like F-, as the size of the cation increases, the magnitude of the lattice enthalpies decrease quickly (become less negative).

On the other hand, the increase in cationic size does not cause a great decrease in the sum of the hydration enthalpies because of the great magnitude of hydration enthalpy of the small anion.

As a result, the solubility increases with an increase in cationic radius.


Questions:

1. The solubility of CsF is relatively high. Explain.

2. The solubility of LiF is relatively low. Explain.

3. The solubility of CsI is relatively low. Explain.



Entropy6.7 Entropy Change (p. 164)

Scientists want to find out what governs a spontaneous reaction:

• an exothermic reaction is a spontaneous reaction while an endothermic reaction is not (exothermicity)

any exception?

• Some spontaneous change is endothermic, e.g. melting of ice at room temperature.

besides enthalpy change, there is another factor that determine a chemical reaction. Entropy

Entropy is a measure of the randomness or the degree of disorder of a system.

Entropy is a measure of the randomness or the degree of disorder of a system.

Q. 11


Entropy6.7 Entropy Change (p. 164)

In any spontaneous process, there is always an increase in the entropy (disorder) of the system and its surroundings --- The second law of thermodynamics.

In any spontaneous process, there is always an increase in the entropy (disorder) of the system and its surroundings --- The second law of thermodynamics.

Melting of ice at room temperature increase in entropy of water molecules

Dissolving of sodium chloride in water increase in entropy


6.7 Entropy Change (p. 164)

Entropy Change S

S = Sfinal - Sinitial

Entropy change is temperature dependent:High temperature entropy increasesLow temperature entropy decreases

Unit of entropy: kJ mol-1 K-1


Free Energy6.7 Free Energy Change (p. 168)

When predicting whether a reaction is spontaneous, we need to consider both enthalpy change and entropy change (at constant temperature).

A new term is developed to include both enthalpy and entropy free energy:

G = H – TS where G = change in free energy

A process is spontaneous when G is negative (H is –ve, S is +ve)A process is not spontaneous when G is positive.


The END

New Way Chemistry for Hong Kong A-Level Book 11 Chapter 6 Energetics.

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