Enthalpy and entropy

Question 1

What is the lattice enthalpy defintion?

Answer 1

Lattice enthalpy is the enthalpy change that accompanies the formation of one mole of an ionic compound from its gaseous ions under standard conditions.

Question 2

What is lattice enthalpy?

Answer 2

K+(g) + Cl-(g) –> KCl(s)
It is exothermic, so will always be a negative value.
Lattice enthalpy is a measure of the strength of ionic bonding in a giant ionic lattice.

Question 3

What are ionic compounds?

Answer 3

Solid ionic compounds are very stable - due to the strength of ionic bonds, the electrostatic attractions between oppositely charged ions.
This creates a large energy barrier that must be overcome to break down the lattice, reflected in the high melting points in many ionic copmounds.

Question 4

What is a Born-Haber cycle?

Answer 4

Lattice enthalpy cannot be measured directly.
Ionic lattice -ΔHf-> elements in standard states -ΔHat-> gaseous atoms -ΔHIE(1/2)-> gaseous ion of metal -ΔHEA(1/2)-> gaseous ions with non-metal -ΔHLE-> ionic lattice

Question 5

What are the routes of the Born-Haber cycle?

Answer 5

Route 1(clockwise): formation of gaseous atoms from their standard states - endothermic (bond breaking).
Formation of gaseous ions from atoms - overall endothermic
Lattice formation from gaseous atoms to solid ionic lattice - exothermic.
Route 2 (anti-clockwise): direct conversion of elements in lattice to standard states - exothermic.

Question 6

What is the standard enthalpy change of formation?

Answer 6

ΔfH is the enthalpy change that takes place when one mole of a compound is formed from its elements under standard conditions, with all the reactants and products in their standard states.
Na(s) + 1/2Cl2(g) –> NaCl(s)

Question 7

What is the standard enthalpy change of atomisation?

Answer 7

ΔatH is the enthalpy change that takes place for the formation of one mole of gaseous atoms from its elements in its standard state under standard conditions.
Na(s) –> Na(g)
1/2Cl2 (g) –> Cl(g)
Always endothermic as bonds are being broken.

Question 8

What is the first ionisation energy?

Answer 8

ΔIEH is the enthalpy change required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.
Na(g) –> Na+(g) + e-
It is endothermic because energy is required to overcome the attraction between a negative electron and the positive nucleus.

Question 9

What is first electron affinity?

Answer 9

ΔEAH is the enthalpy change that takes place when one electron is added to each atom in one mole of gaseous atoms to form 1 mole of gaseous 1- ions.
Cl(g) + e- –> Cl-(g)
It is exothermic because the electron being added is attracted towards the nucleus.

Question 10

What is the rule for Born-Haber cycles?

Answer 10

Between each horizontal energy level:
Only one species has changed, matching the enthalpy change that takes place.
All the species are balanced.
Sum of anticlockwise = sum of clockwise
Br2 standard state is liquid. I2 is solid.

Question 11

What are successive electron affinities?

Answer 11

When an anion has a greater charge than -1.
O (g) + e- –> O- (g)
O- (g) + e- –> O2- (g)
Second electron affinities are endothermic - a second electron is being gained by a negative ion, which repels the electron, so energy must be used to force the negative electron onto the negative ion.

Question 12

What is the definition for the standard enthalpy change of solution?

Answer 12

ΔsolH is the enthalpy change that takes place when one mole of compound is completely dissolved in water under standard conditions.
The ions from the ionic lattice finish up surrounded with water molecules as aqueous ions.

Question 13

What is enthalpy change of solution?

Answer 13

Na+Cl- (s) + aq –> Na+(aq) + Cl-(aq)
It can be endothermic or exothermic, depending on the relative sizes of lattice enthalpy, and enthalpy of hydration.
When aqueous, the 𝛿+ and 𝛿- partial charges in the water molecules are attracted towards the positive and negative ions.

Question 14

What is the enthalpy change of hydration?

Answer 14

ΔhydH is the enthalpy change when 1 mole of gaseous ions is dissolved in water to form 1 mole of aqueous ions, under standard conditions.
Na+ (g) + aq –> Na+ (aq)
Cl- (g) + aq –> Cl- (aq)

Question 15

What is the energy cycle for enthalpy of solution?

Answer 15

Gaseous ions -ΔhydH-> aqueous ion of metal -ΔhydH-> aqueous ion of both <-ΔsolH- ionic lattice -ΔLEH-> gaseous ions
The ΔsolH will be above or below the ionic lattice depending on whether it is endo or exothermic.
This is worked out by adding up the ΔhydH and comparing to LE.

Question 16

What are the properties of ionic compounds?

Answer 16

Generally:
High melting and boiling points
Soluble in polar solvents
Conduct electricity when molten or in aqueous solution.

Question 17

What is the effect of ionic size on lattice enthalpy?

Answer 17

Down the group - increased ionic size - increased ionic radius, electrostatic attraction between ions decreases, lattice enthalpy less negative, melting point decreases.
So NaCl has most negative lattice enthalpy, highest melting point, whereas RbCl has a less negative lattice enthalpy and a lower melting point.
Across the period the ions become smaller.

Question 18

What is the effect of ionic charge on lattice enthalpy?

Answer 18

Across the period, ionic charge increases, electrostatic attraction between ions increases, lattice energy becomes more negative, and melting point increases.
Na2O will have a less negative lattice enthalpy than CaO due to having a 1+ charge, not 2+.

Question 19

How do you compare ionisation energies of elements?

Answer 19

Atomic radius of the element or atomic charge (not ionic).
Greater nuclear attraction.

Question 20

How can melting points be predicted?

Answer 20

The magnitude of lattice energy indicates the melting point of an ionic compound.
Some metal oxides - MgO, Al2O3 and ZrO have very exothermic lattice enthalpies and very high melting points.
These stable metal oxides are used as a protective coating for the inside of furnaces and refractories.

Question 21

What factors affect enthalpy of hydration?

Answer 21

Ionic size - Down the group, ionic radius increases, attraction between ion and water molecules decreases, hydration enthalpy is less negative.
Ionic charge - Across the period, ionic charge increases, attraction with water molecules increases, hydration enthalpy becomes more negative.

Question 22

How is an ionic compound dissolved?

Answer 22

To dissolve an ionic compound in water, the attraction between the ions in the ionic lattice must be overcome.
This requires a quantity of energy equal to the lattice enthalpy.
Water molecules are attracted to the positive and negative ions, surrounding them and releasing energy equal to hydration enthalpy.

Question 23

How can solubility be predicted?

Answer 23

If the sum of hydration enthalpies is larger than the magnitude of the lattice enthalpy, the overall enthalpy change (ΔsolH) will be exothermic and the compound should dissolve.
However, many compounds with endothermic energy changes are soluble, so solubility also depends on temperature and entropy.

Question 24

What is entropy?

Answer 24

S is the term used for the dispersal of energy within the chemicals making up the chemical system.
Measured in J K^-1 mol^-1
The greater the entropy value, the greater that energy is spread out per Kelvin per mole.

Question 25

What are the general entropy trends?

Answer 25

The greater the entropy, the greater the dispersal of energy and the greater the disorder.
Solids have the smallest entropies.
Liquids have greater entropies.
Gases have the greatest entropies.

Question 26

What are entropy changes?

Answer 26

At 0 Kelvin there would be no energy and all substances would have 0 entropy.
Above 0 K, energy becomes dispersed amongst the particles and all substances have positive entropy.
If a system changes to become more random, energy can be spread out more, so ΔS will be positive.

Question 27

How can entropy change be predicted using changes of state?

Answer 27

Entropy increases during changes in state that give a more random arrangement of particles:
Solid –> Liquid –> Gas
So from solid to liquid to gas, entropy increases, the energy is spread out more.

Question 28

How can entropy change be predicted using a change in the number of gaseous molecules?

Answer 28

Reactions that produce a gas increase in entropy.
CaCO3(s) + 2HCl(aq) –> CaCl2(aq) + CO2(g) + H2O(l)
Production of a gas increases the disorder of particles.
Energy is spread out more and ΔS is positive.

Question 29

How is the sign for entropy change predicted?

Answer 29

If a reaction has a different number of gas molecules between the reactants and products.
N2(g) + 3H2(g) –> 2NH3(g)
There is a decrease in the randomness of particles, the energy is less spread out and ΔS is negative.

Question 30

What is standard entropy?

Answer 30

Sθ of a substance is the entropy of one mole of a substance, under standard conditions (100kPa and 298K).
Sθ have units of J K^-1mol^-1
They are always positive.
Calculated by Sθ = ∑Sθ (products) - ∑Sθ (reactants)

Question 31

What is feasibility?

Answer 31

Feasibility (spontaneous) describes whether a reaction is able to happen and is energetically feasible.
A reaction can happen if the products have a lower overall energy than the reactants.

Question 32

What is the free energy change?

Answer 32

ΔG is the overall energy change in a chemical reaction, made up from:
The enthalpy change ΔH - the heat transfer between the chemical system and the surroundings.
The entropy change at the temperature of the reaction TΔS.

Question 33

What is the Gibbs’ equation?

Answer 33

ΔG = ΔH - TΔS
Free energy is in KJmol-1
Temperature is in Kelvin.
Entropy needs dividing by 1000 to put into kJ K^1mol^-1
For a reaction to be feasible, ΔG < 0 ( decrease in free energy).

Question 34

What is the balance between ΔH and TΔS?

Answer 34

ΔH_______ΔS_______T____________ΔG________Feasible
Negative_Positive__Low or high_Negative__Yes
Positive__Negative_Low or high_Positive___No
Negative_Negative_Low_________Negative__Yes
____________________High_________Positive___No
Positive__Positive_Low__________Positive__No
_____________________High_________Negative__Yes

Question 35

What are the limitations of predictions made for feasibility?

Answer 35

Many reactions have a negative ΔG but do not seem to take place.
This is because a very large activation energy results in a very slow rate.
So, although the sign of ΔG indicates the thermodynamic feasibility, it takes no account of the kinetics or rate of reaction.

Question 36

The student carries out a second experiment using 150 cm 3 of distilled water instead of 100 cm 3 of distilled water. The mass of concentrated sulfuric acid is the same as in the first experiment.
Predict and explain the effect, if any, of the larger volume of water on the following:
* The temperature change, ΔT
* The calculated value of ΔsolH for H2SO4

Answer 36

ΔT is less and ΔsolH is the same.
ΔT less because the same heat energy is spread over a larger volume of water.
ΔsolH is the same because the same energy released per mole of H2SO4

Question 37

Student 2 carries out the same experiment but uses twice the mass of CaCl2 . All other quantities are the same.
Predict any differences between the temperature change and the calculated value of ΔsolH from the experiments of the two students. Explain your reasoning

Answer 37

Temperature change is double because twice the energy is produced in the same volume.
ΔsolH is the same because the ratio of energy produced to mass is the same.