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AQA A-Level Physical Chemistry > 1.8 Thermodynamics > Flashcards

1.8 Thermodynamics Flashcards

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1
Q

What is a Born-Haber cycle?

A

An application of Hess’ law

Where there are multiple steps in one leg of the cycle

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2
Q

How are ionic solids produced?

A

Between ions in the gas phase

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3
Q

What is exothermic and endothermic in terms of bonds? Arrow representation?

A

Exothermic - bond forming (down arrow)

Endothermic - bond breaking (up arrow)

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4
Q

What is significant about the ionic solid?

A

The ions in an ionic lattice are electrostatically bonded by ionic bonds

The standard lattice enthalpy of formation is a measure of the strength of an ionic bond

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5
Q

What is standard enthalpy of formation? ΔHθf

A

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

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6
Q

What is standard enthalpy of combustion? ΔHθc

A

The enthalpy change when one mole of a substance reacts completely in excess oxygen under standard conditions with all reactants and products in their standard states

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7
Q

What is ionisation enthalpy? ΔHθi

A

The standard enthalpy change when one mole of GASEOUS ATOMS is converted into one mole of GASEOUS IONS with a single POSITIVE charge

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8
Q

What is bond dissociation enthalpy? ΔHθdiss

A

The enthalpy change when one mole of GASEOUS molecules each breaks a covalent bond to form 2 free radicals, averaged over a range of compounds

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9
Q

What is enthalpy of atomisation? ΔHθat

A

The standard enthalpy change that accompanies the formation of one mole of GASEOUS ATOMS from the element in its standard state under standard conditions

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10
Q

What is electron affinity? ΔHθea

A

The standard enthalpy change when one mole of GASEOUS ATOMS is converted into one mole of GASEOUS IONS, each with a single NEGATIVE charge

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11
Q

What is the lattice enthalpy of formation? ΔLHθ

A

The standard enthalpy change that accompanies the formation of one mole of a solid ionic compound from its gaseous ions

Exothermic

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12
Q

What is the enthalpy of lattice dissociation? ΔLHθ

A

The standard enthalpy change that accompanies the separation of one mole of a solid ionic compound into its gaseous ions

Endothermic

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13
Q

How can we tell the difference between ΔLHθ with the formation and dissociation of lattices?

A

We can only tell from the positive or negative values due to bond forming/making therefore exo/endo

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14
Q

What is the enthalpy of hydration? ΔhydHθ

A

The standard enthalpy change for the process where one mole of gaseous ions are surrounded by a maximum of water molecules to produce one mole of hydrated ions

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15
Q

What is enthalpy of solution? ΔsolHθ

A

The standard enthalpy change that occurs when one mole of an ionic solid dissolves completely in sufficient solvent to form a solution in which the molecules or ions are far enough apart not to interact with each other

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16
Q

In ionisation enthalpy what must you specify?

A

Whether it is first or second

Both are endothermic

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17
Q

In electron affinity what must you specify?

A

Whether it is first or second electron affinity

First - exothermic
Second - endothermic

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18
Q

Why is the first and second electron affinity exothermic and then endothermic?

A

First is exothermic as an attraction is formed between a neutral atom and an electron

Second is endothermic as energy needs to be put in to over come the repulsion of two negative ions

19
Q

What are the intermediates in a simple hess cycle for Enthalpy of lattice formation/dissociation?

A

Elements
The ionic solid

At the top the gaseous ions of the elements

20
Q

What are lattice enthalpy values based on?

A

The Born-Haber cycles

They are experimental values

21
Q

What is the perfect ionic model?

A

They are theoretical values

22
Q

What does the perfect ionic model presume?

A

Assumes there is 100% ionic bonding

The ions are perfect spheres with a uniform distribution of ion charge

23
Q

Why are experimental values and theoretical values of lattice enthalpy different?

A

In reality ions are perfectly spherical due to different charge densities and electrons aren’t always completely donated
Therefore there can be a degree of covalency
= not perfect ionic bonding

Mor discrepancy = greater degree of covalency

25
Q

What happens to a positive ion in water? (Enthalpy of solution)

A

The δ- oxygen of the polar water molecule is electrostatically attracted to the positive metal ion

26
Q

What happens to a negative ion in water? (Enthalpy of solution)

A

The δ+ hydrogen of the polar water molecule is electrostatically attracted to the negative metal ion

27
Q

What is significant about water surrounding ions in solution?

A

As soon as the gas ions are surrounded they are classed as aqueous

But there is a limit to the number of water molecules to surround the ion
= infinite dilution

28
Q

Is Enthalpy of solution exothermic or endothermic?

A

It could be either

Depends on the reaction

29
Q

How does solubility of a solution related to being exo/endothermic?

A

More exothermic = more soluble

More endothermic = less soluble

30
Q

What are the intermediates in a simple hess cycle for Enthalpy of solution?

A

The ionic solid
The gaseous ions
The aqueous ions

Check the arrow direction!!!

31
Q

Why does enthalpy of hydration decrease down a group?

A

Ionic radius increases
= charge density decreases
= the electrostatic attraction to the δ- oxygen of the polar molecule decreases
= less exothermic as the attraction is weaker

32
Q

What is entropy?

A

A measurement of disorder

The more disordered something is the more entropy it has

Nature likes disorder

33
Q

How does entropy change with states?

A

Solid (low entropy)
Liquid
Gas (high entropy)

34
Q

What is the explanation for the entropy graph when increasing temperature- changing state?

A

Increase in temp
Increases kinetic energy
Particle start to move/vibrate
And there is more disorder

But there is greater disorder going from liquid to gas as ALL the particles separate completely

35
Q

How can entropy increase?

A

More particles from reactants to products eg 2 -> 7
Changing state towards a gas
A molecule with more electrons (as they are moving around)

36
Q

What is entropy change?

A

∆S system =
ΣS products - ΣS reactants

If ∆S is positive - increase in entropy = increase in disorder

37
Q

What is significant about spontaneity, enthalpy and entropy?

A

Exothermic changes are spontaneous

Endothermic aren’t normally but if they have a high entropy they can be spontaneous

38
Q

When does something have zero entropy?

A

At absolute 0 (0 kelvin)

Particles have zero energy so entropy is also zero

39
Q

What is Gibbs free energy?

A

∆G
Combines the influence of both enthalpy and entropy

The energy available to do useful work

40
Q

What is the Gibbs free energy equation?

A

∆G = ∆H - T∆S

∆G - Gibbs free energy
∆H - enthalpy change
T - Temperature
∆S - entropy change

Always write an equation out!

41
Q

How does the value of ∆G affect spontaneity and feasibility?

A

∆G less than 0
Spontaneous and feasible

∆G = zero
Can be feasible or unfeasible

∆G more than 0
Not spontaneous, needs more energy to be feasible

42
Q

When is ∆G = zero?

A

During a change of state

43
Q

What can we assume when ∆G = zero?

A

∆H = T∆S

So

T = ∆H/∆S

44
Q

What do you need to remember with all Gibbs free energy equations?

A 
∆H = KJ mol-1
∆S = J mol-1

You need to convert ∆S to KJ!!!!

45
Q

How can we link y=mx+c to Gibbs free energy equation?

A

Y = M x +c

∆G = -∆S T +∆H

AQA A-Level Physical Chemistry (12 decks)

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