Module 5: Chapter 22 - Enthalpy and Entropy

Question 1

What is Lattice enthalpy?

Answer 1

The enthalpy change that accompanies the formation of one mole of an ionic compounds from its gaseous ions under standard conditions

Question 2

What is the shorthand writing of Lattice enthalpy change?

Answer 2

ΔʟᴇH

Question 3

What type of energy change is lattice enthalpy?

Answer 3

Exothermic. Lattice enthalpy will always be exothermic as it involves the formation of ionic bonds, therefore it will always have a negative value

Question 4

How can you determine lattice enthalpy?

Answer 4

Lattice enthalpy cannot be directly measured, therefore it must be calculated using other data in a born-haber cycle

Question 5

What is a born-haber cycle?

Answer 5

A type of energy cycle used to analyse reaction energies

example of a born haber cycle

Question 6

What is the standard enthalpy change of formation, ΔfH?

Answer 6

The enthalpy change that takes place when one mole of a compound is formed from its elements under standard conditions, with all reactants and products in their standard states

Question 7

What is the standard enthalpy change of atomisation, ΔₐₜH?

Answer 7

The enthalpy change that takes place for the formation of one mole of gaseous atoms from the element in its standard state under standard conditions

Question 8

What type of process is the enthalpy change of atomisation?

Answer 8

It is always an endothermic process as bonds are broken to form gaseous atoms. Therefore, it is always a positive value

Question 9

What is first ionisation energy, ΔɪᴇH?

Answer 9

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.

Question 10

What type of process is ionisation energy?

Answer 10

It is always an endothermic process as energy is required to overcome the attraction between positive nuclei and negative electon. Therefore, it always has a positive value

Question 11

What is first electron affinity, ΔᴇᴀH?

Answer 11

The enthalpy change that takes place when one electron is added to each atom in one mole of gaseous atoms to form one mole of gaseous 1- ions

Question 12

What type of process is first electron affinity?

Answer 12

First Electron affinity is always an exothermic process

Question 13

How does second electron affinity differ from first electron affinity?

Answer 13

First electron affinity is exothermic, however second electron affinity (and all subsequent electron affinitys) are endothermic. This is because the second electron is being gained by a negative anion which repels the electron away. Therefore, energy must be put in to force the negatively-charged electron onto the negative ion

Question 14

Answer 14

Lattice Enthalpy = -790 kJ mol⁻¹

Question 15

Answer 15

Lattice enthalpy = -2528 kJ mol⁻¹

Question 16

What is the standard enthalpy change of solution, ΔₛₒₗH?

Answer 16

The standard enthalpy change of solution is the enthalpy change that takes place when one mole of a solute dissolves in a solvent

Question 17

What does “+ aq” mean in a chemical equation?

Answer 17

add excess H₂O(l)

Question 18

How can you experimentally determine the enthalpy change of solution?

Answer 18

1. Weigh out a sample of the ionic solid
2. Measure out 25cm³ of water and place thermometer in the water
3. Pour all of the solid into the water, stir the mixture and record the temperature change until it no longer increases
4. Calculate the energy change using the equation “q = mcΔt” where m is the sum of the mass of the water and solid
5. Calculate number of moles of ionic substance
6. Calculate enthalpy change of solution by dividing energy by number of moles

Question 19

What are the 2 stages of dissolving an ionic substance in water?

Answer 19

• The ionic lattic breaks up
• Water molecules are attracted to, and surround, the ions

Question 20

Describe the energy changes involved with the dissolving process (enthalpy change of solution)

Answer 20

1. The ionic lattice is broken up, forming separate gaseous ions. This is the opposite process to lattice enthalpy. This process is always endothermic
2. The separate gaseous ions interact with polar water molecules to form hydrated aqueous ions. The energy change involved in this is known as enthalpy change of hydration. This process is always exothermic

Question 21

What is enthalpy change of hydration, ΔhydH?

Answer 21

The enthalpy change of hydration is the enthalpy change that accompanies the dissolving of gaseous ions in water to form one mole of aqueous ions

Question 22

Is enthalpy change of solution (ΔₛₒₗH) endothermic or exothermic?

Answer 22

It depends on the relative sizes of the reverse lattice enthalpy (always endothermic) and the enthalpy change of hydration (always exothermic)

Question 23

What factors affect lattice enthalpy and hydration enthalpy?

Answer 23

• Ionic size
• Ionic charge

Question 24

How does ionic size effect lattice enthalpy?

Answer 24

As ionic radius increases, the attraction between ions decreases, therefore lattice enthalpy becomes less negative (decreases in magnitude)

Question 25

How does ionic charge affect lattice enthalpy?

Answer 25

As ionic charge increases, the attraction between ions increases, therefore lattice enthalpy becomes more negative (increases in magnitude)

Question 26

Describe the effect of ionic charge and size on attraction across period 3 from sodium to aluminium:

Answer 26

From sodium to aluminium: * Increasing charge gives more attraction * Decreasing size gives more attraction These are 2 **supporting** effects

Question 27

Describe the effect of ionic charge and size on attraction across period 3 from chlorine to sulfur:

Answer 27

From chlorine to sulfur: * Increasing charge gives more attraction * Increasing size gives less attraction These are 2 **opposing** effects

Question 28

How can lattice enthalpy be used to predict melting points?

Answer 28

**Usually** the more exothermic the lattice enthalpy is, the higher the melting point. However it can also depend on other factors such as the packing of ions in an ionic lattice

Question 29

How does ionic size affect hydration enthalpy?

Answer 29

As ionic radius increases, the attraction between the ion and the water molecule decreases, so hydration energy becomes less negative

Question 30

how does ionic charge affect hydration enthalpy?

Answer 30

As ionic charge increases, the attraction with water molecules increases, therefore hydration energy becomes more negative

Question 31

How can you predict solubility?

Answer 31

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. Therefore, if the sum of the hydration enthalpies is larger than the magnitude of the lattice enthalpy, the overall energy change will be exothermic and the compound *should* dissolve. However temperature and entropy can also affect this and therefore some compounds are soluble with endothermic enthalpy changes

Question 32

What is entropy?

Answer 32

The natural tendency for energy to spread out rather than to be concentrated in one place. The greater the entropy, the greater the dispersal of energy and the greater the disorder

Question 33

What are the units of entropy?

Answer 33

J K⁻¹ mol⁻¹

Question 34

In general, how does entropy differ between states of matter?

Answer 34

* Solids have the smallest entropies * Liquids have greater entropies * Gases have the greatest entropies

Question 35

What can entropy be used to explain?

Answer 35

* A gas spreading out through a room * Heat from a fire spreading through a room * Ice melting in a hot room In all of these examples energy is being dispersed and becoming more spread out

Question 36

Explain entropy values at and above 0K

Answer 36

* At 0K a substance would have an entropy value of 0 * Above 0K energy becomes dispersed amongst the particles and all substances have a positive **entropy** value

Question 37

Explain the entropy change (ΔS) if a system changes to become more random:

Answer 37

If a system changes to become more random, energy can be spread out more, therefore there will be a positive entropy change

Question 38

Explain the entropy change (ΔS) if a system changes to become less random:

Answer 38

If a system changes to become less random, energy can be spread out less, therefore there will be a negative entropy change

Question 39

How can you predict entropy changes in physical changes/chemical reactions?

Answer 39

The side of the equation with a greater number of moles of gas has greater entropy

Question 40

Explain how melting/boiling causes an increase in entropy:

Answer 40

Melting and boiling increases the randomness of particles, energy is being spread out more and therefore there is an increase in entropy

Question 41

What is the Standard entropy (Sθ) of a substance?

Answer 41

The entropy of one mole of a substance, under standard conditions of 100kPa and 298K

Question 42

How can you calculate entropy changes?

Answer 42

ΔSθ = ΣSθ(products) - ΣSθ(reactants)

Question 43

Can entropy have a negative value?

Answer 43

No, entropy values are always positive (except 0 at 0K). However entropy **changes** can be negative

Question 44

What is the feasibility of a reaction?

Answer 44

The term used to describe whether a reaction is able to happen and is energetically feasible

Question 45

What does spontaneous mean?

Answer 45

Energetically feasible

Question 46

What is free energy change ΔG?

Answer 46

The overall energy change during a chemical reaction

Question 47

What does free energy change consist of?

Answer 47

* The enthalpy change ΔH. THis is the heat transfer between the chemical system and the surroundings * The entropy change at the temperature of the reaction TΔS. This is the dispersal of energy within the chemical system itself

Question 48

How can you calculate the free energy change?

Answer 48

The Gibbs' free energy equation

Question 49

What is the Gibb's equation?

Answer 49

ΔG = ΔH - TΔS

Question 50

What are the conditions for energetic feasibility of a reaction?

Answer 50

* There must be a decrease in free energy * ΔG < 0

Question 51

What are the units of ΔS **in the Gibb's equation**?

Answer 51

**They must be *changed* to kJ K⁻¹ mol⁻¹**. Normally the units are J K⁻¹ mol⁻¹, however they must be changed to kJ K⁻¹ mol⁻¹ to match the kJ in the ΔH

Question 52

What are the limitations of predictions made for feasibility based on ΔG?

Answer 52

Despite indacting thermodynamic feasibility, **it does not account for the kinetics or rate of reaction**. Therefore there may be an extremely high activation energy and the rate of reaction may be extremely low so it may not seem like the reaction is occuring

Question 53

Calculate the lattice enthalpy for sodium bromide

Answer 53

-752 kJ mol⁻¹

Question 54

Calculate the lattice enthalpy of sodium oxide

Answer 54

-2520 kJ mol⁻¹

Question 55

Calculate the lattice enthalpy of calcium fluoride

Answer 55

-2611 kJ mol⁻¹

Question 56

Explain the difference between the lattice enthalpies of NaCl and MgCl₂

Answer 56

Mg²⁺ ions are smaller and have a greater charge than Na⁺, therefore they have a greater charge density. As a result, there is a greater attraction between ions in MgCl₂ compared to NaCl. Therefore, the lattice enthalpy is more exothermic

Question 57

Explain the difference between the lattice enthalpies of NaBr and KCl

Answer 57

Na⁺ ions are smaller than K⁺ ions and therefore have a greater attraction, however Cl⁻ ions are smaller than Br⁻ ions, therefore there are competing factors and it is hard to determine which lattice enthalpy will be more exothermic

Question 58

Explain the difference between the lattice enthalpies of MgO and MgF₂

Answer 58

F⁻ ions are smaller than O²⁻ ions but have a smaller charge, therefore there are competing factors and it is hard to determine which lattice enthalpy will be more exothermic