Thermodynamics

Define enthalpy change

Heat energy change at a constant pressure

Define lattice formation enthalpy

• The enthalpy change when 1 mole of a solid ionic compound is formed from its gaseous ions

• E.g. Na⁺ (g) + Cl^- (g) →NaCl (s)

Is lattice formation enthalpy exothermic or endothermic - explain why?

• Exothermic

• In the ionic compound formed new attractions between the ions are being formed

• Energy released when attractions formed

Define lattice dissociation enthalpy

• Enthalpy change when 1 mole of a solid ionic compound is completely dissociated into its gaseous ions.

• E.g. NaCl (s) → Na⁺ (g) + Cl^- (g)

Is lattice dissociation enthalpy exo or endo - explain why?

• Endothermic

• Breaking attractions requires energy

What does the value of lattice enthalpies depend on?

• Strength of electrostatic attraction between ions

What does the strength of electrostatic attraction depend on and explain?

• Size of ion→ Smaller the radius of the ion, stronger the attraction (smaller ions can pack closely together in ionic lattice)

• Magnitude of charge→ Larger charge, stronger the attraction

(Basically it depends on charge density)

Define charge density

• How concentrated the charge is in an ion

• The higher the charge density→ Stronger attraction between ions

• More exothermic the lattice formation enthalpy

What is required to dissolve an ionic compound?

• Where does it come from

• And what proceeds to happen

• Energy required to break apart the lattice (Lattice dissociation enthalpy)

• Energy comes from the water (solvent)

• Water forms new attractions to the ions in the lattice→ Creating an aqueous solution

• Ions dissolved until there is infinite dilution meaning no electrostatic attraction between + and - ion, attraction is now between ions and water solvent.

Define enthalpy of hydration

• Enthalpy change when 1 mole of gaseous ions become aqueous ions

• E.g. Na⁺ (g)→ Na⁺ (aq)

• Exothermic

Define enthalpy of solution

• Enthalpy change when 1 mole of an ionic solid dissolves in enough solvent to form an infinitely dilute solution

• NaCl (s) + aq → Na⁺ (aq) + Cl^- (aq)

What does the enthalpy of solution depend on?

• Balance between the lattice dissociation enthalpy which is endothermic and hydration enthalpy which is exothermic

• If magnitude of hydration enthalpy> lattice dissociation enthalpy → the enthalpy of solution will be exothermic / vice versa

Hess cycle for enthalpy of solution

• Switch direction of arrow for lattice dissociation enthalpy

Hoe does the Born Haber cycle work

• Works in the same way as Hess cycle

• Enthalpy change is still independent to route taken

• Born Haber cycles have energy as a y axis

• Arrow pointing up = Endothermic

• Arrow pointing down= Exothermic

What is lattice enthalpy

• A measure of attraction between ions

• (Cannot be measured directly in an experiment)

What is theoretical lattice enthalpy?

• Lattice enthalpy value you calculate using a model that assumes ions are perfect spheres and considers size, charge and arrangement of ions in the lattice.

What is experimental lattice enthalpy?

• Carry out a series of experiments

• Construct a Born-Haber cycle

• To workout lattice enthalpy

What is enthalpy of formation?

• Enthalpy change when 1 mole of a compound is formed from its elements in their standard states.

• (In BH cycle its 1 mole of an ionic compound)

• Exothermic

• E.g. Na (S) + ½ Cl₂ → NaCl (S)

What is first ionisation energy?

• Energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous ions with a +1 charge.

• Endothermic

• E.g. Mg (g) → Mg⁺ (g) + e^-

What is second ionisation energy?

• Energy required to remove 1 mole of electrons from 1 mole of gaseous 1+ ions to produce one mole of gaseous 2+ ions

• Endothermic

• E.g. Mg⁺ (g) → Mg²⁺ (g) + e^-

What is first electron affinity?

• Enthalpy change when 1 mole of gaseous atoms gain 1 mole of electrons to form 1 mole of gaseous ions with a -1 charge

• Exothermic (New attraction formed between the electron and nucleus of atom)

• E.g. o (g) + e^- → o^- (g)

What is second electron affinity?

• Enthalpy change when 1 mole of gaseous 1- ions gain one electron per ion to produce gaseous 2- ions

• Endothermic ( The o- and e- repel each other - energy supplied/needed to overcome that repulsion)

• E.g. o^- (g) + e- → o^2- (g)

What is atomisation (for an element)

• Enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard state.

• Endothermic

• E.g. Na (s) → Na (g)

Example of atomisation for a diatomic element ?

• ½ Cl₂ (g) → Cl (g)

• ( Half the moles of Cl₂ has the same number of Cl atoms as 1 mole of Cl (g) )

What is bond disassociation enthalpy

• Standard molar enthalpy change when one mole of a covalent bond is broken into 2 gaseous atoms

• E.g. Cl₂ (g) → 2Cl (g)

What is atomisation ( For a compound)

• Enthalpy change when 1 mole of a compound in standard state is converted into gaseous atoms

• E.g. NaCl (s) → Na (g) + Cl (g)

Enthalpy change of formation = sum of all enthalpy changes when..

• Born Haber cycle includes lattice formation enthalpy

General format of Born Haber cycle

1) Enthalpy change of formation

2) Atomisation of x

3) Atomisation of y

4) F.I.E of x

5) Electron affinity of y

6) Lattice formation enthalpy

What is a flaw of theoretical lattice enthalpy?

• Makes the assumption that ions are perfectly spherical

• But ions don’t exist as isolated perfect spheres

• Instead electron clouds around ions are attracted towards each ion

What is meant by polarisation of negative ion?

• Electron cloud around negative ion is attracted towards positive ion

• So there is electron density between the nuclei of the ions

What effects the polarisation of a negative ion?

• Higher charge density of the positive ion →More polarisation of negative ion

• A large singly charged negative ion will be more polarisable→ Electron cloud is further away from nucleus so it is more easily distorted by positive ion → A singly charged ion has more electrons than protons so the nucleus cannot hold them as tightly. (But in a small singly charged negative ion the cloud is closer to nucleus so less polarisable)

• (Electron cloud refers to all electrons but the outermost electrons are mainly involved in polarisation)

What does the experimental lattice enthalpy value reflect?

• Polarisation of negative ions

• Lattice has covalent character (if there a higher degree of polarisation)

Why does a lattice have covalent character in experimental lattice enthalpy value?

• Because of the electron density between the nuclei of the 2 ions during polarisation

• (Basically what happens in covalent bond - i.e shared electron attracted to nuclei of both atoms)

• So ions are held more tightly together

• And the experimental lattice enthalpy (of formation) is more exothermic

-Check notes for additional guidance

Chemical reactions that happen spontaneously are …

• Feasible reactions

When does a reaction occur and how does this happen?

• When it is overwhelmingly probable by chance alone

• Happens because a chemical system will change from one where there is limited ways of arranging energy to a system with more ways of arranging energy.-Occurs by random changes

• Trend is from order to disorder

• (When we consider whether or not a reaction is going to be feasible we are basically asking whether this will result in an increase in disorder)

Why are exothermic reactions more likely to be feasible?

• Heat energy is released to surrounding gas molecules

• Warmer gas molecules→ have more kinetic energy→ move around more→ become more disordered.

Define entropy

• Entropy of a chemical system is a measure of the energy that a system has at a particular temp, per mole of each chemical

Units of entropy (s)

• Joules per kelvin per mole

• J K ^-1 mol^-1

Which way do feasible reactions go spontaneously?

• In the direction of increasing disorder i.e increasing entropy

What does an increase in moles mean for entropy?

• More moles of a chemical→ more disordered than a few moles

• Increase in moles→ Increase in entropy

• E.g. 1mole (of reactants) → 2moles of products : Increasing entropy i.e positive entropy change

• 2moles → 1 mole : Decreasing entropy, negative entropy change (This reaction is less likely to be feasible as there’s a decrease in disorder)

Equation to calculate entropy change

What is Gibbs Free- energy change ∆G

• Balance between entropy and enthalpy in a system

• The balance determines feasibility of a reaction

• (Reaction will only be feasible if there is an overall movement of “free-energy” [all types of energy] out of the chemical system to the surrounding)

Gibbs free energy equation and units

A reaction is feasible at a specific temperature if ..

• Gibbs free energy change is equal to or below zero (negative)

∆G is more likely to be negative and the reaction feasible if..

• ∆H is negative ( Exothermic reaction)

• ∆S is positive (Increased disorder)

When is a reaction never feasible?

• When ∆S is negative and when ∆H is positive (endo)

If a reaction has -∆H (exo) and -∆S (decreasing disorder)..

• the reaction will be feasible if ∆H> T∆S

• it gets less feasible as temp increases because enthalpy value is what will cause this reaction to be feasible so its size needs to be greater.

If a reaction has +∆S and +∆H..

• Reaction will be feasible if T∆S>∆H

• So as temp increases this reaction will get more feasible.

Why may a reaction not always be spontaneous even if it is feasible?

• ∆G can be negative → making reaction feasible but activation energy must be very high

• so even if ∆G is negative, reaction may not be spontaneous if Ea. is very high

What temperature does a reaction become feasible?

• This temperature is the point where the reaction goes from being infeasible to feasible (∆G= 0)

• So when asked to workout the temperature for when the reaction becomes feasible assume ∆G is 0

Rearrange Gibbs free energy equation to find temperature

• T = ∆H / ∆S

• Before substituting in entropy change, you must convert it into KJ K^-1 Mol^-1

A situation where ∆G = 0 and what this allows you to do?

• A change of state

• This is because at MP (S→L) and BP (L→G) forward and reverse reactions occur at the same rate.(Dynamic eqm)

• Can use Gibbs free energy equation to workout out at what temp MP or BP should occur

Gibbs free energy equation into straight line equation

• ∆G = -∆ST +∆H

• Y = m x + c

• Gradient is - ∆S
  

Draw and explain a free energy graph with a positive gradient

• Refer to notes.

Draw and explain a free energy graph with a negative gradient

• refer to notes

Draw a graph for a reaction that is never feasible and one where the reaction is always feasible

• Refer to notes

More negative a value for ∆G …

• More feasible the reaction

• More likely that the reaction goes to completion (almost all of the reactants are converted into products)

More positive a value for for ∆G …

• The less likely the reaction will go to completion

• Reverse reaction is more feasible

• ( If the ∆G value for the forward reaction is +250 the value for the reverse reaction will be -250)