All of Unit 3 WJEC A-Level Chemistry

Oxidation

Loss of electrons/hydrogen, gain of oxygen

Reduction

Gain of electrons/hydrogen, loss of oxygen

Oxidising agent

Species that oxidises another species, and gains electrons in the process, becomes reduced itself

Reducing agent

Species that reduces another species, and loses electrons in the process, becomes oxidised itself

Oxidation number of elements in their standard state

0

Oxidation number of hydrogen

1 unless it’s with a Group 1 metal, then it’s -1

Oxidation number of oxygen

-2 unless it’s a peroxide when it’s -1, or reacted with fluorine, when it’s +2

Oxidation number of group 1

1

Oxidation number of group 2

2

Oxidation numbers of elements in a compound or ion

Sum = 0 or charge on the ion

Oxidation numbers and electronegativity

The most electronegative element is given the negative oxidation number.

Oxidation numbers and reduction

Decreases if reduced

Oxidation numbers and oxidation

Increases if oxidised

Standard Electrode Potential

p.d/EMF measured when a half-cell is connected to the standard hydrogen electrode

Standard conditions

• 298K temperature

• 1atm pressure

• 1 moldm^-3 comcentration

Standard hydrogen electrode components

• Platinum electrode (s)

• H+ solution (aq) 1.00 moldm^-3

• H2 (g) at 1atm (100kPa)

• 298K

Standard hydrogen electrode conditions

• H+ solution (aq) 1.00 moldm^-3

• H2 (g) at 1atm (100kPa)

• 298K

Function of platinum electrode in SHE

• Conducts electricity

• Unreactive

• Porous

Reaction between 2 SHEs

No redox reaction

Reaction between SHE and half-cell with another metal

Redox as circuit allows electrons to flow between them

How electrochemical cells are formed

2 metal/metal ion half-cells

Components of electrochemical cells

• High resistance voltmeter and wire

• Salt bridge

How electrochemical cells are connected

• High resistance voltmeter

• Salt bridge

High resistance voltmeter function in cells

Allows electrons to flow/move from the half-cell where oxidation occurs to the half-cell where reduction occurs. Measures p.d. between the two

Salt bridge function

• Completes the circuit

• Maintains electrochemical balance between the 2 half-cells

Composition of salt bridge

Usually made from filter paper soaked in saturated solution of KNO3 (aq)

Electrodes based on different oxidation states

• Where two ions are in solution (Fe2+/Fe3+ or Cr3+/Cr2O7 2-)

• Or where the OS of one species is zero (CL2/Cl-)

Components used when there are 2 ions in solution

• Both solutions included in the same half-cell solution

• Pt (s) as multiple oxidation states

Components used when the oxidation state of one species is zero

Pt (s) electrode

Direction of flow of electrons and SEP value

• L —> R = +ve

• R —> L = -ve

LOAN RRCP rule

• Left, Oxidation, Anode, Negative electrode

• Right, Reduction, Cathode, Positive electrode

Rule used for position (L/R), process and charge

LOAN RRCP

SEP values and oxidising/reducing power

• Most positive = strongest OA

• Most negative = strongest RA

Order of values written in the electrochemical series

OA + e- —> RA

Cell EMF definition

Electromotive Force. Difference between SEPs of the two half-cells

Cell EMF equation

EMF (cell) = E (most positive or reduced) - E (most negative or oxidised)

Cell EMF and feasibility

+ve = reaction is feasible

-ve = reaction is not feasible

Comparing standard electrode potentials

most +ve = reduced

most -ve = oxidised

Representing Cells using Cell Notation

X (s) (electrode) | X⁺ (aq) (solution) || Y²⁺ (aq) | Y(s)

LHS = oxidation from X to X+

RHS = reduction from Y2+ to Y

Single line (|) = change in state

Double line (||) = salt bridge

Cell notation for change in state

single line (|)

Cell notation for salt bridge

Double line (||)

Cell notation when half cell contains 2 ions in solution

Comma between ions as same state

Pt (s) electrode

Hydrogen fuel cell diagram

Principles and equations of the hydrogen fuel cell

Advantages of hydrogen fuel cells

• more efficient release of energy from fuel

• Use a renewable fuel source (hydrogen)

• Only by-product formed is H2O

• Maintain a constant voltage (compared to chemical cells)

• Less heat energy is lost compared to the standard hydrogen combustion engine

Disadvantages of hydrogen fuel cell

• Expensive to produce

• Toxic chemical used in production

• H2(g) is expensive to produce

• One source of H2(g) is fossil fuels

• Non-renewable

• Storing H2(g) can be difficult & dangerous

Metal used in hydrogen fuel cells + why

platinum as it is unreactive

Constructing ion/electron half-equations

1. Write the reagents and products

2. Balance the atoms present

3. Add any extra hydrogen ions that can be used to form water

4. Find the difference in charge between the start and end to find the number of electrons needed

Combining half-equations

1. Write then with one beneath the other so the arrows align and are in the correct directions for the reactions

2. Multiply one or both of them to balance the electrons so both have the same number of electrons

3. Combine into one equation with all reactants on the left and all products on the left

4. Cancel the electrons

5. H+ or H2O can be cancelled too if needed

Need for acidic conditions in redox reactions

• Oxidising agent contains oxygen

• H+ needed to combine with the oxygen to produce H2O as a product

• Number of H+ depends on number of water molecules

Equation for reduction of acidified dichromate

Colour change in reduction of acidified dichromate

Orange to green

Equation for reduction of acidifed manganate

Colour change for reduction of acidified permanganate

Purple to pale pink/colourless

Equation for reduction of iodine

Colour change for reduction of iodine

Brown to colourless

Blue-black to colourless if starch indicator is added

Equation for oxidation of thiosulfate

Colour change for reduction of thiosulfate

None

Equation for oxidation of oxalate

Equation for oxidation for Fe3+

Colour change for oxidation of iron II

Pale green to pale yellow

Equations for redox reaction between Cu²⁺ and I^-

Reaction between aqueous thiosulfate ions and aqueous iodine

• Thiosulfate oxidised by iodine

• Thiosulfate in butte

• Iodine pipetted into conical flask

• Thiosulfate run into flask until colour due to iodine fades to pale yellow

• Starch solution added as indicator

• Colour change to dark blue

• End point = blue colour decolourised

Thiosulfate and iodine equation

Reaction between aqueous iodide ions and aqueous copper (II) ions

• Blue solution loses its colour

• White ppt of CuI forms

• Brown solution of I2

• I2 titrated with thiosulfate ions

Iodide and copper (II) equation

Aqueous dichromate ions and iron (II) ions reaction

• Aqueous dichromate ions placed in burette

• Added to conical flask

• Oxidises aqueous Fe2+

• Indicator required

Dichromate and iron (II) equation

Dichromate to chromate(VI) ions reaction

Not redox as chromium does not change is oxidation number

Dichromate to chromate(VI) colour change

Orange to yellow

Dichromate to chromate(VI) equation

Aqueous manganate(VII) ions and iron(II) ions reaction

• Aqueous MnO4^- in burette

• Fe2+ solution pipetted into conceal flask with excess aqueous sulfuric acid

• No indicator required

• End point; all Fe2+ has been oxidised —> next drop of MnO4^- gives a pink colour

Manganate and iron(II) equation

Redox titrations vs acid base

• Similarity; use a solution of known concentration of ions to find the contraction of ions in another

• Difference; redox involving transition metals usually involve a colour change to no need for indicator to be used

p block electron configuration

Outermost electrons in a p orbital

Oxidation states of p block elements

Maximum oxidation state = group number

Lower = group number-2

Amphoteric

Metals that react with/act as an acid and a base

Examples of amphoteric behaviour

• Al³+/Al

• Pb²+/Pb

Aluminium Oxide acting as a base

Al₂O₃ + 6HCl —> 2AlCl₃ + 3H₂O

Aluminium oxide acting as an acid

Al₂O₃ + 2NaOH + 3H₂O —> 2Na⁺ + 2[Al(OH)₄]^- (2Na[Al(OH)₄]

Aluminium Hydroxide acting as a base

Al(OH)₃ + 3H⁺ —> Al³⁺ + 3H₂O

OR Al(OH)₃ + 3HCl —> AlCl₃ + H₂O

Aluminium Hydroxide acting as an acid

Al(OH)₃ (s) + OH^- —> [Al(OH)₄]^- (aq)

Lead Oxide acting as a base

PbO + 2HCl —> PbCl₂ + 2H₂O

OR PbO + 2HNO₃ —> Pb(NO₃)₂ + H₂O

Lead Oxide acting as an acid

PbO + 2NaOH + H₂O —> 2Na⁺ + [Pb(OH)₄]^2- (Na₂[Pb(OH)₄])

Lead Hydroxide acting as an acid

Pb(OH)₂ (s) + 2OH^- —> [Pb(OH)₄]^2- (aq)

Lead hydroxide acting as a base

Pb(OH)₂ + 3HCl —> PbCl₂ + H₂O

OR Pb(OH)₂ + 2H⁺ —> Pb²⁺ + 2H₂O

Amphoteric metal solutions and NaOH

Solid white precipitate of metal hydroxide

Precipitate then redissolves of excess to form a colourless solution

Al³⁺ and NaOH (aq) observations

White precipitate of aluminium hydroxide formed

Redissolves in excess to form a colourless solution

Al³⁺ and NaOH (aq) equation

Al³⁺ (aq) + 3OH^- (aq) —> Al(OH)₃ (s) white precipitate then + OH^- (aq) —> [Al(OH)₄]^- (aq) colourless solution

Pb²⁺ and NaOH (aq) observations

White precipitate of lead hydroxide formed

Redissolves in excess to form a colourless solution

Pb²⁺ and NaOH (aq) equation

Pb2+ (aq) + 2OH- —> Pb(OH)2 (s) white precipitate then +2OH^- —> [Pb(OH₄)]^2- (aq) colourless solution

Stable oxidation states of group 3

+3 and +1Stable ox

Stable oxidation states of group 4

+4 and +2

Stable oxidation states of group 5

+5 and +3

Trend in oxidation states down the group

Elements are more likely to form compounds in which they have the lower oxidation state so lower down the group you go, the more likely it is that the elements forms a compound with the lower oxidation state

Inert pair effect

the tendency of the outermost s2 pair of electrons in an atom to remain unshared in compounds, leading to a lower oxidation state.

Trend in inert pair effect

Occurs in groups 3, 4 and 5

Tendency increases/becomes more significant down the group

Cause of trend in inert pair effect down the group

The two outer electrons in the outer s orbital are less likely to take part in bonding

Octet expansion

the ability of some atoms to use d-orbitals to have more than 8 electrons in their valence (outer) shell.