Electron transfer reactions — IB Chemistry HL Revision Notes & Topical Questions

Section 1 of 72Oxidising and reducing agents

Oxidising agent

• Oxidising agent oxidises another species by taking its electrons;
  • it is itself reduced and its oxidation number decreases.

Reducing agent

• Reducing agent reduces another species by giving electrons;
  • it is itself oxidised and its oxidation number increases.

Redox

• Redox needs both an oxidising and a reducing agent.

Dual role

• Dual role some species, such as H₂O₂, can act as either depending on what they react with.

Identify

• Identify assign oxidation numbers on both sides of the equation.

Stock notation

• Stock notation the Roman numeral in a name shows the metal's oxidation state, for example iron(II) sulfate contains Fe²⁺.

Section 2 of 72Balancing redox equations and redox titrations

1

Write the unbalanced equation and identify the atoms that change oxidation number.

2

Deduce the oxidation number changes.

3

Balance the oxidation number changes.

4

Balance the charges.

5

Balance the atoms.

Half-equations in acid

• Half-equations in acid balance atoms other than H and O, add H₂O for O, add H⁺ for H, add e⁻ for charge, then combine.

Manganate(VII)

• MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O; self-indicating as the purple colour disappears at the end point.

Iodine–thiosulfate

• 2S₂O₃²⁻ + I₂ → 2I⁻ + S₄O₆²⁻; starch is added near the end point and the blue-black colour disappears.

Uses

• Uses iron tablets with manganate(VII) and chlorine in bleach with thiosulfate.

Section 3 of 72Relative ease of oxidation and reduction

Relative ease of oxidation and reduction

Metals in groups

• Metals in groups ease of oxidation increases down Group 1 and Group 2.

Halogens

• Halogens oxidising power decreases down Group 17, with F₂ the strongest and I₂ the weakest.

Halogen displacement

• Halogen displacement a more reactive halogen displaces a less reactive one from its halide solution.

Reactivity series

• Reactivity series a more reactive metal displaces a less reactive metal ion, acting as a reducing agent.

Metals with acids

• Metals with acids acid + metal → salt + hydrogen, as 2H⁺ + Zn → Zn²⁺ + H₂ with H⁺ reduced and Zn oxidised.

Rate

• Rate more reactive metals and stronger acids react more vigorously; copper does not react with dilute acids.

Reducing power order

• Mg, Al, Zn, Fe, Pb, H, Cu, Ag, from strongest to weakest.

Section 4 of 72Voltaic cells

1

Voltaic cell

• Voltaic cell generates electricity from a spontaneous redox reaction by separating the two half-cells.

2

Half-cell

• Half-cell a metal dipped in a solution of its own ions, with an equilibrium setting up an electrode potential.

3

EMF

• EMF the potential difference between two half-cells; only differences can be measured.

4

Anode

• Anode oxidation, negative; cathode — reduction, positive.

5

Electron flow

• Electron flow through the external circuit from anode to cathode.

6

Salt bridge

• Salt bridge an inert electrolyte such as KNO₃ that completes the circuit, with anions moving to the anode and cations to the cathode.

7

Cell notation

• Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s), with a single line for a phase boundary and a double line for the salt bridge.

Section 5 of 72Primary, secondary and fuel cells

Primary cell

• Primary cell the reaction is not reversible and the cell is discarded when used up.

Secondary cell

• Secondary cell the reaction is reversed by applying a voltage greater than the cell voltage.

Lead–acid

• Pb and PbO₂ in H₂SO₄, about 2 V per cell and 12 V from six cells; heavy, with toxic lead and corrosive acid.

NiCad

• Cd and NiO(OH), about 1.2 V; cadmium is toxic and the cell has a memory effect.

Lithium-ion

• Lithium-ion lithium cobalt oxide and carbon, 3.5–4.0 V;
  • light with no memory effect, but cells degrade, lithium is limited and there is a fire risk.

Fuel cell

• Fuel cell fuel and oxygen are supplied continuously;
  • the hydrogen–oxygen cell gives E = +1.23 V with only water as product.

Section 6 of 72Electrolytic cells

1Electrolysis

Electrolysis splitting a compound using electrical energy, driving a non-spontaneous reaction.

2Electrolyte

• Electrolyte molten or dissolved ionic compound.

3Anode

• Anode positive, where anions are oxidised.

4Cathode

• Cathode negative, where cations are reduced ("RED CAT").

5Ion movement

• Ion movement cations move to the cathode and anions to the anode.

6Molten binary compounds

• Molten binary compounds metals form at the cathode and non-metals at the anode, as for PbBr₂.

7Comparison

• Comparison in a voltaic cell the anode is negative and the cathode positive;
  • in an electrolytic cell the anode is positive and the cathode negative.

8Uses

• Uses purifying copper, plating, extracting aluminium and making chlorine and sodium hydroxide.

Section 7 of 72Organic redox reactions

Primary alcohol

• Primary alcohol oxidised to an aldehyde and then to a carboxylic acid (reflux with excess oxidant).
  • (distil off as formed)

Secondary alcohol

• Secondary alcohol oxidised to a ketone only.

Tertiary alcohol

• Tertiary alcohol not oxidised, because the functional group carbon has no hydrogen.

Oxidant

• Oxidant acidified potassium dichromate(VI), orange Cr₂O₇²⁻ turning green Cr³⁺.

Reduction

• Reduction carboxylic acids to primary alcohols with LiAlH₄ in dry ether;
  • aldehydes to primary alcohols and ketones to secondary alcohols with LiAlH₄ or NaBH₄.

Reagents

• NaBH₄ cannot reduce carboxylic acids; both give hydride H⁻.

Alkenes and alkynes

• Alkenes and alkynes hydrogenation with a nickel catalyst, about 200 °C and 1000 kPa, as in making margarine.

Section 8 of 72Standard electrode potentials (HL)

Standard electrode potentials (HL)

Standard hydrogen electrode

• H₂ at 100 kPa, 1.00 mol dm⁻³ H⁺ and platinum, 2H⁺ + 2e⁻ ⇌ H₂, E° = 0.00 V.

Standard electrode potential

• Standard electrode potential the EMF with the SHE on the left and the half-cell on the right at 1.0 mol dm⁻³, 100 kPa and 298 K.

Reduction potentials

• Reduction potentials half-reactions are written as reductions, as in Section 19 of the data booklet.

More negative E°

• More negative E° stronger reducing agent, equilibrium to the left.

More positive E°

• More positive E° stronger oxidising agent, equilibrium to the right.

Cell EMF

• E°_cell = E°_right − E°_left = E°_reduction − E°_oxidation.
  What the symbols mean: E_cell = standard cell potential (V) · E_right = electrode potential of the positive (reduction) half-cell · E_left = potential of the negative half-cell (V)

Spontaneity

• Spontaneity a positive E°_cell means the reaction is spontaneous, but thermodynamics predicts and kinetics controls, so a reaction may not be seen.

Section 9 of 72Gibbs energy and cell potential (HL)

Equation

• ΔG° = −nFE°, where n is the moles of electrons and F = 9.65 × 10⁴ C mol⁻¹.

Positive E°

• Positive E° negative ΔG°, spontaneous.

Negative E°

• Negative E° positive ΔG°, so the forward reaction is not spontaneous and the reverse is.

Equilibrium

• ΔG = 0 and E = 0, as when a voltaic cell is exhausted.

Data booklet

• Data booklet the ΔG° = −nFE° equation and F are given.

Section 10 of 72Electrolysis of aqueous solutions and electroplating (HL)

Electrolysis of aqueous solutions and electroplating (HL)

Competing species

• Competing species water can be reduced to H₂ or oxidised to O₂, so it competes with the ions.

Factors

• Factors standard electrode potentials, concentration of the ions and the identity of the electrodes.

Dilute H₂SO₄ or NaOH

• Dilute H₂SO₄ or NaOH water is electrolysed to hydrogen and oxygen.

NaCl(aq)

• NaCl(aq) hydrogen forms at the cathode; at high chloride concentration chlorine forms at the anode, giving NaOH(aq).
  • (above about 25 %)

CuSO₄(aq) with inert electrodes

• CuSO₄(aq) with inert electrodes copper forms at the cathode and oxygen at the anode.

CuSO₄(aq) with copper electrodes

• CuSO₄(aq) with copper electrodes the anode dissolves, which purifies copper, with impurities falling as anode slime.

Electroplating

• Electroplating the object is the cathode and the anode is the plating metal, which keeps the ion concentration stable.

Thickness

• Thickness controlled by current, time and electrolyte concentration.

Plastics

• Plastics they are etched and given embedded metal particles to conduct before plating.

Section 11 of 72Ease of oxidation in Groups 1 and 2

1

A metal is oxidised when it loses electrons.

2

Lithium reacts with water: 2Li(s) + 2H₂O(l) → 2LiOH(aq) + H₂(g).

• Lithium is oxidised from 0 to +1.

3

The reaction becomes more vigorous down Group 1, so the ease of oxidation increases down the group.

4

The same trend holds in Group 2.

• Mg(s) + 2H₂O(l) → Mg(OH)₂(s) + H₂(g) is very slow.
• Ca(s) + 2H₂O(l) → Ca(OH)₂(s) + H₂(g) is very quick.

Section 12 of 72Halogens as oxidising agents

1

Halogens react with metals by accepting an electron to form a 1− ion, as in Ca + Cl₂ → Ca²⁺(Cl⁻)₂.

2

A halogen removes electrons from the metal, so the metal's oxidation number increases.

3

The halogen is reduced, so its oxidation number decreases.

4

Oxidising power decreases down Group 17.

• F₂ is the strongest oxidising agent because it is the most readily reduced.
• I₂ is the weakest oxidising agent because it is the least readily reduced.

Section 13 of 72Halogen displacement reactions

Halogen displacement reactions

A more reactive halogen displaces a less reactive halogen from an aqueous solution of its halide.

Reactivity of the non-metals in Group 17 increases up the group, so Cl₂ > Br₂ > I₂.

Colours in aqueous solution:

Halogen	Colour in aqueous solution
Chlorine	Very pale green, usually appears colourless as it is very dilute
Bromine	Orange, yellow when very dilute
Iodine	Brown

Halogen displacement reactions

Chlorine with bromide: 2KBr(aq) + Cl₂(aq) → 2KCl(aq) + Br₂(aq), and the solution turns orange.

Bromine with iodide: Br₂(l) + 2NaI(aq) → 2NaBr(aq) + I₂(aq).

A halogen lower in the group cannot displace one higher up, so iodine cannot displace chlorine from NaCl.

Section 14 of 72Metal displacement and the reactivity series

1

Metals are ranked in the reactivity series (activity series).

2

A more reactive metal displaces a less reactive metal from solutions of its ions or from its oxides.

3

Mg(s) + CuSO₄(aq) → MgSO₄(aq) + Cu(s) splits into two half-equations.

• Mg → Mg²⁺ + 2e⁻ is oxidation (loss of electrons).
• Cu²⁺ + 2e⁻ → Cu is reduction (gain of electrons).

4

The more reactive metal acts as the reducing agent, because it pushes electrons onto the ions of a less reactive metal.

5

Order of reducing power, strongest first: Mg, Al, Zn, Fe, Pb, H (reference), Cu, Ag.

• Metals at the top lose electrons most easily and are readily oxidised.
• Metals at the bottom lose electrons less easily and are weak reducing agents.

Section 15 of 72Acids with reactive metals

Acids with reactive metals

General reaction: acid + metal → salt + hydrogen.

• 2HCl(aq) + Zn(s) → ZnCl₂(aq) + H₂(g)
• H₂SO₄(aq) + Fe(s) → FeSO₄(aq) + H₂(g)

A salt is an ionic compound formed when the hydrogen of an acid is replaced by a metal or another positive ion.

The extent of reaction depends on the reactivity of the metal and the strength of the acid.

• Very reactive metals react dangerously, so these reactions are not usually carried out.
• Copper does not react with dilute acids.
• Stronger acids react more vigorously, with more effervescence and faster dissolving of the metal.

Section 16 of 72Ionic equation for a metal with an acid

1

Write the full equation: 2HCl(aq) + Zn(s) → ZnCl₂(aq) + H₂(g).

2

Split the aqueous substances into ions: 2H⁺ + 2Cl⁻ + Zn → Zn²⁺ + 2Cl⁻ + H₂.

3

Remove the spectator ions (Cl⁻) to get 2H⁺(aq) + Zn(s) → Zn²⁺(aq) + H₂(g).

4

Write the half-equations: 2H⁺ + 2e⁻ → H₂ (reduction) and Zn → Zn²⁺ + 2e⁻ (oxidation).

• H⁺ is the oxidising agent (+1 to 0) and Zn is the reducing agent (0 to +2).

Section 17 of 72More from the specification

1

Standard Electrode Potentials in Electrochemical Cells

• The standard electrode potential (E°) measures the tendency of a chemical species to be reduced at standard conditions, measured relative to the standard hydrogen electrode (SHE, 0.00 V).
  • (298 K, 1 mol dm^-3, 100 kPa)

2

Redox Processes at Electrodes

• Anode vs. Cathode — In both galvanic and electrolytic cells, oxidation occurs at the anode and reduction occurs at the cathode (gain of electrons), driving electron flow through the external circuit.
  • (loss of electrons)

3

Standard Cell Potential and Spontaneity

• The overall cell potential of an electrochemical cell is calculated using E°_cell = E°_cathode - E°_anode;
  • a positive E°_cell indicates a thermodynamically feasible redox reaction under standard conditions.

4

Electrolysis vs. Galvanic Energy Conversion

• Galvanic cells convert spontaneous chemical energy into electrical potential difference,
• whereas electrolysis uses an external electric current to drive non-spontaneous redox reactions, discharging ionic species at the anode and cathode.

5

Electrolytic cell

• A cell using electricity to drive a non-spontaneous reaction.

6

Carboxylic acid

• A compound containing –COOH with a name ending in -oic acid.

7

Primary alcohol

• An alcohol in which the C–OH carbon is attached to one other carbon.

8

Spontaneous reaction

• A reaction with ΔG < 0.

9

Secondary alcohol

• An alcohol whose –OH carbon is bonded to two other carbons.

10

Sodium chloride

• Compound of sodium and chlorine; common salt.

11

Alcohol

• A compound containing –OH with a name ending in -ol.

12

Cathode

• The negative electrode, where reduction occurs.

13

Electrolysis

• The decomposition of a molten or aqueous ionic compound by an electric current.

14

Half-cell

• One electrode and its solution, in which a redox equilibrium is set up.

15

Electrode potential

• A measure of how easily a species is reduced, relative to a reference half-cell.

Section 18 of 72Types of electrochemical cell

1

Redox reactions involve electrons flowing from the reducing agent to the oxidising agent, and the direction depends on the reactivity of the species.

2

Voltaic (galvanic) cells generate electricity from spontaneous redox reactions.

3

Electrolytic cells use electrical energy to drive non-spontaneous reactions.

4

A voltaic cell produces a potential difference called the electromotive force (EMF), also called the cell potential, with symbol E.

5

Only differences in potential can be measured, not the absolute value of one cell.

Section 19 of 72Half-cells and electrode potential

1

A half-cell is a metal rod (electrode) dipped into a solution of its own ions, and an equilibrium is set up between the metal and its ions.

• Zn(s) ⇌ Zn²⁺(aq) + 2e⁻

2

If zinc atoms leave electrons on the rod and enter solution, the rod becomes negative.

3

If Zn²⁺ ions take electrons from the rod and deposit as atoms, the rod becomes positive.

4

In both cases a potential difference exists between the rod and the solution, called the electrode potential.

5

A copper half-cell works in the same way: Cu²⁺(aq) + 2e⁻ ⇌ Cu(s).

6

No net chemical reaction occurs in a single half-cell, only a potential difference.

Section 20 of 72The zinc–copper (Daniell) cell

1

In Zn(s) + CuSO₄(aq) → Cu(s) + ZnSO₄(aq), the two redox processes are separated into two half-cells instead of electrons passing directly between zinc and copper ions.

2

Electrons flow through the external circuit from the anode (oxidation) to the cathode (reduction).

3

The system then generates electrical energy.

4

Oxidation takes place at the negative electrode and reduction takes place at the positive electrode.

Section 21 of 72Salt bridge and voltmeter

Salt bridge

• Salt bridge completes the circuit by allowing ions to flow between the two solutions.
• It is often filter paper soaked in an inert electrolyte such as KNO₃(aq).
• KCl and KNO₃ are common choices because chlorides and nitrates are usually soluble, so no precipitates form to disturb the half-cell equilibria.
• Anions move towards the anode and cations towards the cathode to balance charge in each half-cell.

Voltmeter

• Voltmeter measures the EMF.
• It has a high resistance so that it diverts very little current from the circuit.

The two half-cells are in series, because the same current flows through both.

Section 22 of 72Conventional cell diagrams

1

A single vertical line (|) shows a phase boundary between a solid and a solution.

2

A double vertical line (||), or dashed lines, shows the salt bridge.

3

The species with the highest oxidation state in each half-cell is written next to the salt bridge.

4

The cell potential is given with the polarity of the right-hand electrode.

5

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s), E_cell = +1.10 V.

• The copper half-cell is more positive, so electrons flow from zinc to copper.

6

The same cell can be written Cu(s) | Cu²⁺(aq) || Zn²⁺(aq) | Zn(s), E_cell = −1.10 V.

• The right-hand electrode is negative, so electrons still flow from zinc to copper.

7

Phase boundaries may also be shown on the outside of the cell diagram.

Section 23 of 72Describing a voltaic cell fully

Describing a voltaic cell fully

Negative electrode (anode)

• Negative electrode (anode) oxidation occurs here and electrons are released.

Positive electrode (cathode)

• Positive electrode (cathode) reduction occurs here.

Electrons flow in the external wire from the negative to the positive electrode.

Ions flow in the salt bridge, with anions towards the anode and cations towards the cathode.

Removing the salt bridge stops the cell, because charge builds up in each half-cell and prevents further reaction.

Section 24 of 72More from the specification

More from the specification

Galvanic electrochemical cell group

• Galvanic electrochemical cell group non-rechargeable primary cells belong to a chemical battery group converting irreversible chemical reactions into electrical energy.

Standard Hydrogen Electrode (SHE)

• The reference standard assigned an electrode potential of 0.00 V, consisting of platinum in contact with 1.0 mol dm⁻³ H⁺ and hydrogen gas at 100 kPa.

Electrochemical Series and Reducing Agents

• More reactive metals with more negative standard electrode potentials act as stronger reducing agents, displacing less reactive metals from aqueous solutions.

Oxidation of Primary Alcohols

• Primary alcohols are oxidised in the presence of acidified dichromate to produce carboxylic acids and water, involving multi-step electron transfer.

Balancing Redox Equations

• In aqueous half-cell equations, oxygen atoms are balanced by adding water (H₂O), and hydrogen atoms are balanced by adding H⁺ ions.

Standard hydrogen electrode

• Standard hydrogen electrode acts as the universal reference electrode with an assigned potential of exactly 0.00 V under standard conditions.

Reactivity of metals

• Reactivity of metals more reactive metals have more negative electrode potentials and readily displace less reactive metal cations from aqueous solution.

Balancing redox equations

• Balancing redox equations write separate oxidation and reduction half-equations and equalize electron counts before adding water or hydrogen ions.

Carboxylic acid

• A compound containing –COOH with a name ending in -oic acid.

Primary alcohol

• An alcohol in which the C–OH carbon is attached to one other carbon.

Secondary alcohol

• An alcohol whose –OH carbon is bonded to two other carbons.

Sodium chloride

• Compound of sodium and chlorine; common salt.

Hydrogen

• A gas that gives a squeaky pop with a lighted splint.

Alcohol

• A compound containing –OH with a name ending in -ol.

Section 25 of 72The hydrogen–oxygen fuel cell

1

A fuel cell is an electrochemical cell in which a fuel donates electrons at one electrode and oxygen gains electrons at the other.

2

Fuel cells are becoming common in the automotive industry as a replacement for petrol and diesel engines.

3

The hydrogen–oxygen fuel cell has:

• A reaction chamber with separate inlets for hydrogen and oxygen gas.
• An outlet for the product, water.
• An electrolyte of aqueous sodium hydroxide.
• A semi-permeable membrane that separates the two gases.

4

Half-equations:

• 2H₂(g) + 4OH⁻(aq) → 4H₂O(l) + 4e⁻, E = −0.83 V.
• O₂(g) + 2H₂O(l) + 4e⁻ → 4OH⁻(aq), E = +0.40 V.

5

Overall reaction: 2H₂(g) + O₂(g) → 2H₂O(l), E = +1.23 V.

Section 26 of 72Advantages and problems of fuel cells

Advantages and problems of fuel cells

Advantages:

• Water is the only product, so fuel cells have clear environmental advantages.
• The reaction is the same as combustion of hydrogen but happens at room temperature,
  • so the energy becomes electrical energy rather than heat and light.
• No harmful oxides of nitrogen form, as they do in high-temperature combustion in air.
• Fuel cells have been used on spacecraft, where the water can be drunk by astronauts.

Problems:

• Hydrogen is highly flammable, so its production and storage carry safety hazards.
• Thick-walled cylinders are needed to store hydrogen, which has economic costs.
• Hydrogen is a by-product of the crude oil industry, so it relies on a finite, non-renewable resource.
• Hydrogen has high energy density per gram but low energy density per unit volume, so larger containers are needed than for liquid fuels.

A fuel cell works continuously while hydrogen and oxygen are supplied, and energy is not stored in the cell.

Section 27 of 72Voltaic and electrolytic cells compared

Voltaic and electrolytic cells compared

Voltaic cells generate electricity from a spontaneous chemical reaction that drives electrons around a circuit.

Electrolytic cells use electrical energy to drive a non-spontaneous reaction, reversing the normal direction of chemical change.

• Stable compounds such as sodium chloride or lead bromide can be broken down into their elements.

Oxidation takes place at the anode and reduction takes place at the cathode in both types of cell.

Cell	Negative electrode	Positive electrode
Voltaic	Anode (oxidation)	Cathode (reduction)
Electrolytic	Cathode (reduction)	Anode (oxidation)

• The memory aid RED CAT states that REDuction takes place at the CAThode.

Section 28 of 72Secondary cells and the lead–acid battery

1

Secondary (rechargeable) cells use reactions that can be reversed by applying a voltage greater than the cell voltage,

• which pushes electrons in the opposite direction.

2

Common examples are lead–acid batteries, nickel–cadmium (NiCad) cells and lithium cells.

3

A lead–acid battery has six cells in series.

• The negative electrode is lead and the positive electrode is lead(IV) oxide.
• The electrolyte is sulfuric acid.

4

Half-equations on discharge:

• Pb(s) + SO₄²⁻(aq) → PbSO₄(s) + 2e⁻, E = −0.36 V.
• PbO₂(s) + 4H⁺(aq) + SO₄²⁻(aq) + 2e⁻ → PbSO₄(s) + 2H₂O(l), E = +1.70 V.

5

Overall: PbO₂ + 4H⁺ + 2SO₄²⁻ + Pb → 2PbSO₄ + 2H₂O, E_cell = +2.06 V.

6

Each cell gives about 2 V, so six cells give about 12 V in a car.

7

While driving, the generator reverses the discharge reaction and regenerates the electrodes.

Section 29 of 72Lead–acid battery uses and disadvantages

Lead–acid battery uses and disadvantages

It delivers a high current for a short time, which suits starting engines.

Disadvantages:

• Very heavy and bulky.
• Contains toxic lead and lead dioxide.
• Sulfuric acid is highly corrosive.
• Disposal and recycling are environmentally challenging.

Section 30 of 72Choosing a cell for an application

1

Primary cells are cheap, light and have a long shelf life, but cannot be recharged and suit only low-power devices.

2

Fuel cells give a steady but low current and need a continuous supply of fuel.

3

Lead–acid batteries suit high-current needs such as starter motors.

4

Rechargeable cells suit applications that need repeated use, such as power tools and vehicles.

Section 31 of 72More from the specification

More from the specification

Alkaline dry cell group

• Alkaline dry cell group zinc-manganese dioxide primary cells form a commercial cell group offering steady terminal voltage and reliable discharge capacity.

Standard Electrode Potentials (E°)

• The standard electrode potential (E°) of a half-cell is the potential difference measured under standard conditions relative to the standard hydrogen electrode (0.00 V).
  • (298 K, 100 kPa, 1.0 mol dm⁻³ ion concentrations)

Electrochemical Cell EMF Calculations

• The electromotive force (EMF) of a galvanic cell is calculated as E°_cell = E°_reduction(cathode) - E°_oxidation(anode),
• where a positive E°_cell indicates thermodynamic feasibility under standard conditions.

Oxidising and Reducing Agents

• A half-cell with a more positive standard electrode potential contains species that are more easily reduced, acting as powerful oxidising agents that oxidise species with more negative potentials.

Oxidation of Alcohols with Acidified Potassium Dichromate

• Primary alcohols are oxidised by acidified potassium dichromate(VI) to aldehydes under distillation,
• and further to carboxylic acids under reflux,
• while secondary alcohols are oxidised to ketones.

Cell potential calculation

• Cell potential calculation cell potential equals cathode reduction potential minus anode reduction potential measured under standard conditions.

Oxidising agents

• Oxidising agents species with higher positive reduction potentials,
• such as acidified dichromate, readily gain electrons from reducing metals or sulfite ions.

Metal displacement

• Metal displacement zinc metal immersed in copper(II) sulfate solution becomes oxidised, while copper cations deposit as copper metal.

Half-cell

• One electrode and its solution, in which a redox equilibrium is set up.

Oxidising agent

• A substance that oxidises another and is itself reduced.

Electrode potential

• A measure of how easily a species is reduced, relative to a reference half-cell.

Reducing agent

• A substance that reduces another and is itself oxidised.

Carboxylic acid

• A compound containing –COOH with a name ending in -oic acid.

Alcohol

• A compound containing –OH with a name ending in -ol.

Metal

• An element with a lattice of positive ions in a sea of delocalised electrons, usually shiny, conductive and malleable.

Section 32 of 72NiCad cells

1

NiCad cells come in many standard sizes and voltages, so they can replace traditional zinc–carbon cells.

2

They are more expensive, but they can be recharged hundreds of times, so they are commercially viable.

3

The negative electrode is cadmium and the positive electrode is a nickel(II) hydroxide–oxide system.

4

Half-equations on discharge:

• Cd(s) + 2OH⁻(aq) → Cd(OH)₂(s) + 2e⁻, E = −0.82 V.
• NiO(OH)(s) + H₂O(l) + e⁻ → Ni(OH)₂(s) + OH⁻(aq), E = +0.38 V.

5

Overall: 2NiO(OH) + 2H₂O + Cd → 2Ni(OH)₂ + Cd(OH)₂, E = +1.2 V.

6

Disadvantages:

• Cadmium is toxic, so disposal is hazardous and the cells must be recycled.
• The memory effect — repeated recharging without full discharge gradually lowers capacity and shortens life.

Section 33 of 72Lithium-ion cells

1

Lithium-ion cells are common in phones, laptops and other portable devices.

2

Lithium is chosen for its low density and high electrode potential.

3

Structure:

• Positive electrode — lithium cobalt oxide.
• Negative electrode — carbon.
• Electrolyte — a solid polymer membrane, which cannot leak and is safer than a liquid or paste.

4

Lithium ions move between the electrodes through the polymer electrolyte during charge and discharge.

5

The cell EMF is between 3.5 V and 4.0 V.

6

The 2019 Nobel Prize in Chemistry was awarded to Goodenough, Whittingham and Yoshino for developing this technology.

Section 34 of 72Comparing cell types

Comparing cell types

Lithium-ion advantages:

• Lightweight with a high voltage output.
• No toxic heavy metals such as lead or cadmium.
• No memory effect, so topping up the charge does not reduce capacity.

Lithium-ion disadvantages:

• Cells lose capacity with repeated cycles.
• Lithium supply is limited, and unrecycled lithium is a lost resource.
• Reactive lithium gives a fire risk if cells are damaged and overheat.

NiCad cells last longer than lead–acid batteries but give a lower voltage, contain toxic cadmium and are expensive.

Section 35 of 72Electrolysis of molten ionic compounds

1

An ionic compound conducts electricity when molten or in solution, and the current splits it into new substances.

2Electrolysis

"splitting by electricity".

• Uses include purifying copper, plating metals with silver and gold, extracting reactive metals such as aluminium and making chlorine, hydrogen and sodium hydroxide.

3

The electrolyte is the substance the current passes through, and it contains positive and negative ions.

4

Cations move to the cathode, where they gain electrons (reduction).

5

Anions move to the anode, where they lose electrons (oxidation).

6

Electrons flow through the external wires and electrodes, and ions move through the electrolyte.

7

Molten lead bromide:

• CathodePb²⁺ + 2e⁻ → Pb.
• Anode2Br⁻ → Br₂ + 2e⁻.

8

Metals form at the cathode and non-metals form at the anode in the electrolysis of a molten binary salt.

Section 36 of 72Classifying alcohols

Classifying alcohols

Alcohols contain at least one –OH group, with general formula C_nH_2n+1OH.

Primary alcohol

• Primary alcohol the carbon bonded to –OH is attached to one other carbon (or alkyl group).

Secondary alcohol

• Secondary alcohol the carbon bonded to –OH is attached to two other carbons.

Tertiary alcohol

• Tertiary alcohol the carbon bonded to –OH is attached to three other carbons.

Section 37 of 72Oxidation of alcohols

1

The oxidising agent is acidified potassium dichromate(VI), K₂Cr₂O₇, with sulfuric acid supplying H⁺.

2

Orange Cr₂O₇²⁻ is reduced to green Cr³⁺, so the colour change is orange to green.

3

Primary alcohols oxidise first to aldehydes and then to carboxylic acids.

4

Secondary alcohols oxidise to ketones only.

5

Tertiary alcohols do not oxidise, because there is no hydrogen on the carbon bonded to –OH.

6

The dichromate test distinguishes tertiary alcohols from primary and secondary alcohols.

Section 38 of 72Conditions for oxidising primary and secondary alcohols

1Carboxylic acid from a primary alcohol

• Carboxylic acid from a primary alcohol heat under reflux with excess oxidising agent.
• CH₃CH₂OH + 2[O] → CH₃COOH + H₂O
• Vapour condenses and drips back into the flask until all the alcohol is oxidised.

2Aldehyde from a primary alcohol

• Aldehyde from a primary alcohol warm with the oxidising agent and distil off the aldehyde as soon as it forms.
• CH₃CH₂OH + [O] → CH₃CHO + H₂O
• The aldehyde has a lower boiling point than the alcohol, so it can be distilled away before it oxidises further.

3Ketone from a secondary alcohol

• CH₃CH(OH)CH₃ + [O] → CH₃COCH₃ + H₂O.
• Ketones cannot be oxidised further, so the product need not be distilled off immediately.

4

Equations use [O] for the oxidising agent and must still be balanced.

Section 39 of 72More from the specification

More from the specification

Oxidising and Reducing Agents

• An oxidising agent gains electrons and is reduced, while a reducing agent loses electrons and is oxidised during a redox chemical process.

Standard Half-Cell Potential

• The potential difference generated by a standard half-cell coupled to the standard hydrogen electrode under standard conditions.
  • (298 K, 1 mol dm⁻³, 100 kPa)

Spontaneous Redox Reactions

• A spontaneous reaction occurs when the overall cell potential difference E°cell is positive, corresponding to a negative Gibbs free energy change.
  • (ΔG° = -nFE°cell)

Balancing Redox Equations with Water

• In acidic aqueous redox half-equations, oxygen atoms are balanced by adding water (H₂O) molecules, followed by adding H⁺ ions to balance hydrogen.

Oxidising and Reducing Agents

• An oxidising agent gains electrons and is reduced, while a reducing agent loses electrons and is oxidised during a redox chemical process.

Standard Half-Cell Potential

• The potential difference generated by a standard half-cell coupled to the standard hydrogen electrode under standard conditions.
  • (298 K, 1 mol dm⁻³, 100 kPa)

Spontaneous Redox Reactions

• A spontaneous reaction occurs when the overall cell potential difference E°cell is positive, corresponding to a negative Gibbs free energy change.
  • (ΔG° = -nFE°cell)

Balancing Redox Equations with Water

• In acidic aqueous redox half-equations, oxygen atoms are balanced by adding water (H₂O) molecules, followed by adding H⁺ ions to balance hydrogen.

Half-cell

• One electrode and its solution, in which a redox equilibrium is set up.

Reducing agent

• A substance that reduces another and is itself oxidised.

Spontaneous reaction

• A reaction with ΔG < 0.

Section 40 of 72Reduction of carbonyl compounds

Reduction of carbonyl compounds

Oxidation: primary alcohol → aldehyde → carboxylic acid, and secondary alcohol → ketone; tertiary alcohols do not react.

These reactions can be reversed by a suitable reducing agent.

• 1Carboxylic acid→2aldehyde→3primary alcohol
• Ketone → secondary alcohol.

Equations for reduction use [H] to represent the reducing agent.

Section 41 of 72Reducing agents

Reagent	Conditions	What it reduces
Lithium aluminium hydride, LiAlH4	Anhydrous (dry ether), followed by aqueous acid	Carboxylic acids, aldehydes and ketones
Sodium borohydride, NaBH4	Aqueous or alcoholic solution	Aldehydes and ketones only

Reducing agents

Both reagents supply the nucleophilic hydride ion, H⁻.

NaBH₄ is less reactive and safer to handle, but it cannot reduce carboxylic acids.

Typical reagents and conditions must be known, but precise temperatures are not needed.

Section 42 of 72Products of reduction

1

Carboxylic acid → primary alcohol, using LiAlH₄ under reflux in dry ether followed by dilute acid.

2

Aldehyde → primary alcohol, using LiAlH₄ or NaBH₄.

3

Ketone → secondary alcohol, using LiAlH₄ or NaBH₄.

4

An aldehyde cannot be isolated from the reduction of a carboxylic acid with LiAlH₄, because the reduction continues to the alcohol.

• To form an aldehyde from a carboxylic acid, reduce to the primary alcohol and then oxidise back to the aldehyde.

Section 43 of 72Reduction of unsaturated compounds

1

Hydrogenation (reduction) of an alkene with hydrogen needs a nickel catalyst, 200 °C and 1000 kPa.

• Ethene is reduced to ethane.

2

Alkynes (C_nH_2n−2) are reduced to alkenes (C_nH_2n) and then to alkanes (C_nH_2n+2).

3

The number of H₂ molecules needed to saturate a molecule depends on the number of multiple bonds.

• An alkene needs one H₂ to become an alkane, and an alkyne needs two.

4

Margarine is made by partly hydrogenating unsaturated vegetable oils.

• Hydrogenation raises the melting point because saturated chains pack closer and have stronger London dispersion forces.
• Controlling the conditions leaves some C=C bonds, which gives the right texture.

Section 44 of 72The standard hydrogen electrode

1

The absolute value of a half-cell potential cannot be measured, only differences between half-cells, so a reference electrode is needed.

2

The standard hydrogen electrode (SHE) consists of:

• Hydrogen gas at 100 kPa.
• H⁺ ions at 1.00 mol dm⁻³.
• An inert platinum electrode in contact with the gas and the ions.

3

The half-equation is 2H⁺(aq) + 2e⁻ ⇌ H₂(g), and E^⦵ is defined as 0.00 V.

4

Connecting another half-cell to the SHE lets its standard electrode potential be read on a high-resistance voltmeter.

5

The SHE is rarely used in practice.

• The electrode reaction is slow.
• The electrode is not easily portable.
• A constant pressure is hard to maintain.

6

Any half-cell with a known potential can serve as a secondary reference.

Section 45 of 72Measuring with the hydrogen electrode

1

The hydrogen electrode is always written on the left of the cell diagram.

2

The half-reaction is always shown as a reduction, so the values are standard reduction potentials.

3

Pt | H₂(g), 2H⁺(aq) || Zn²⁺(aq) | Zn(s) gives E = −0.76 V.

4

Pt | H₂(g), 2H⁺(aq) || Cu²⁺(aq) | Cu(s) gives E = +0.34 V.

5

Data booklet Section 19 lists standard reduction potentials at 298.15 K from most negative to most positive.

Section 46 of 72Interpreting electrode potentials

1

E^⦵ shows how readily a species gains or loses electrons in the half-equation.

2

A more negative E^⦵:

• The species loses electrons more readily.
• The equilibrium lies to the left and favours oxidation.
• It is a stronger reducing agent.

3

A more positive E^⦵:

• The species gains electrons more readily.
• The equilibrium lies to the right and favours reduction.
• It is a stronger oxidising agent.

4

A cell diagram follows the order reduced, oxidised, oxidised, reduced (ROOR), with oxidation on the left.

Section 47 of 72More from the specification

1Metal Reactivity Series and Redox Displacements

• A more reactive metal readily loses electrons to form cations, displacing ions of a less reactive metal (such as copper) from aqueous solutions in an exothermic redox reaction.
  • (such as magnesium or zinc)

2Electrolysis of Aqueous Solutions

• During electrolysis of aqueous copper(II) sulfate using inert electrodes, copper metal is deposited at the cathode,
  • (Cu²⁺ + 2 e- → Cu)
• while water molecules are oxidised at the anode to produce oxygen gas and H⁺ ions.

3Electrorefining of Impure Copper

• In industrial copper purification, an impure copper anode dissolves into electrolyte solution as Cu²⁺ ions,
• while pure copper metal plates out onto the cathode, driven by electrical energy.

4Competitive Discharge of Water vs. Metal Cations

• In aqueous electrolysis, if the metal in solution is more reactive than hydrogen, water is reduced at the cathode to produce hydrogen gas rather than the metal itself.
  • (such as Group 1 or 2 metals)

5Sodium chloride

• Compound of sodium and chlorine; common salt.

6Metal

• An element with a lattice of positive ions in a sea of delocalised electrons, usually shiny, conductive and malleable.

7Cathode

• The negative electrode, where reduction occurs.

8Electrolysis

• The decomposition of a molten or aqueous ionic compound by an electric current.

9Electrolytic cell

• A cell using electricity to drive a non-spontaneous reaction.

10Spontaneous reaction

• A reaction with ΔG < 0.

Section 48 of 72EMF = E(right) − E(left)

EMF = E(right) − E(left)

A voltmeter measures the potential of the right-hand electrode minus that of the left-hand electrode: EMF = E_right − E_left.

A memory aid is "knives and forks": the knife is in the right hand and the fork in the left, so EMF is right minus left.

With the standard hydrogen electrode on the left, Eleft

0 and the EMF equals the electrode potential of the right-hand electrode.

• Zn²⁺ + 2e⁻ ⇌ Zn gives −0.76 V.
• Cu²⁺ + 2e⁻ ⇌ Cu gives +0.34 V.

Section 49 of 72Standard conditions

Standard conditions

A standard electrode potential is the EMF of a cell with the standard hydrogen electrode on the left and the standard half-cell on the right.

Standard conditions are:

• 1.0 mol dm⁻³ concentration of ions.
• 100 kPa pressure.
• 298.15 K.

Section 50 of 72Two uses of the equation

1Calculating an unknown standard electrode potential

• Calculating an unknown standard electrode potential if the EMF and one electrode potential are known, rearrange EMF = E_RHS − E_LHS.

2Calculating a cell EMF

• Calculating a cell EMF if both standard electrode potentials are known and the sides are specified, subtract E_LHS from E_RHS.

Section 51 of 72Predicting the direction of spontaneous change

Predicting the direction of spontaneous change

A positive EMF means the reaction as written is spontaneous.

• It is obtained when the more negative half-cell is on the left and the more positive on the right.

Oxidation takes place on the left, so an alternative form is EMF

E_reduction − E_oxidation.

Method:

1

Find the two half-equations and their E^⦵ values in Section 19 of the data booklet.

2

The half-cell with the more negative E^⦵ is oxidised (left), and the one with the more positive E^⦵ is reduced (right).

3

Calculate EMF = E(reduction) − E(oxidation).

4

If the EMF is positive, the written reaction is spontaneous.

5

If the EMF is negative, the reverse reaction is the spontaneous one.

Section 52 of 72The electrochemical series

1

The electrochemical series lists half-reactions in order of standard electrode potential, from most positive to most negative.

2

A more positive E^⦵:

• The species is more easily reduced.
• It is a stronger oxidising agent.

3

A more negative E^⦵:

• The species is more easily oxidised.
• It is a stronger reducing agent.

4

Metals with the most negative values are the strongest reducing agents.

5

Non-metals with the most positive values are the strongest oxidising agents.

6

"More" and "less" can refer to direction from zero, so always check the numerical value and context.

Section 53 of 72Limits of predictions

Limits of predictions

A positive EMF shows a reaction is energetically feasible but says nothing about its rate.

A reaction with a very high activation energy may be too slow to see at room temperature.

Thermodynamics predicts, and kinetics controls.

Section 54 of 72More from the specification

1

Redox titrations

• Redox titrations determine the unknown concentration of a reducing solution by titrating with a standard oxidising agent such as acidified manganate(VII).

2

Electrolysis electrodes

• Electrolysis electrodes oxidation occurs at the positive anode where anions lose electrons, while reduction occurs at the cathode.

3

Copper refining

• Copper refining during copper electrolysis, impure copper dissolves at the anode and pure copper deposits on the cathode.

4

Voltaic vs electrolytic cells

• Primary voltaic cells convert chemical energy into electrical energy via spontaneous redox reactions,
• whereas electrolysis uses an external power supply to drive non-spontaneous redox reactions in an electrolyte solution.

5

Anode and cathode electrode processes

• In all electrochemical cells, oxidation takes place at the anode and reduction takes place at the cathode (gain of electrons).
  • (loss of electrons)

6

Electrolysis of copper sulfate solution

• During the electrolytic refining of copper using an aqueous copper(II) sulfate electrolyte solution, the impure copper anode dissolves (Cu → Cu²⁺ + 2e-) and pure copper deposits at the cathode.

7

Redox transformations of functional groups

• Primary and secondary alcohols containing the hydroxyl group are transformed through multi-stage redox reactions into aldehydes, ketones,
• or carboxylic acids by electron-transfer oxidants.

8

Electrolytic cell

• A cell using electricity to drive a non-spontaneous reaction.

9

Carboxylic acid

• A compound containing –COOH with a name ending in -oic acid.

10

Primary alcohol

• An alcohol in which the C–OH carbon is attached to one other carbon.

11

Secondary alcohol

• An alcohol whose –OH carbon is bonded to two other carbons.

12

Sodium chloride

• Compound of sodium and chlorine; common salt.

13

Alcohol

• A compound containing –OH with a name ending in -ol.

14

Cathode

• The negative electrode, where reduction occurs.

15

Electrolysis

• The decomposition of a molten or aqueous ionic compound by an electric current.

Section 55 of 72Free energy and cell potential

1

Free energy G measures the energy available to do useful work, and it takes account of the entropy change and the enthalpy change.

2

A reaction is spontaneous when ΔG is negative.

3

By convention the more negative half-cell is on the left, so a spontaneous reaction gives a positive voltmeter reading.

• The more negative electrode pushes electrons to the more positive electrode.
• A digital multimeter showing a negative sign means the terminals are reversed.

4

If E is positive, the reaction is spontaneous as written.

5

If E is negative, the forward reaction is non-spontaneous but the reverse reaction is spontaneous.

Section 56 of 72The relationship ΔG = −nFE

1ΔG⦵

−nFE^⦵.

• n is the number of electrons transferred.
• F is the Faraday constant, 9.65 × 10⁴ C mol⁻¹.

2

Both the equation (Section 1) and the Faraday constant (Section 2) are in the data booklet.

3The units work out as C V mol⁻¹

J mol⁻¹, and the result is often converted to kJ mol⁻¹.

4

Summary:

ΔG	E	Meaning
negative	positive	Reaction is spontaneous
positive	negative	Forward reaction is non-spontaneous and the reverse is spontaneous
zero	zero	System is at equilibrium with no net electron flow

• At equilibrium ΔG 0 and E = 0, which happens when the reactants in a voltaic cell are used up and electron flow stops.

Section 57 of 72Electrolysis of aqueous solutions

1

In molten binary compounds, cations are reduced to the metal at the cathode and anions are oxidised to the non-metal at the anode.

2

In aqueous solutions, water adds more species that can be oxidised or reduced.

• Oxidation2H₂O(l) → 4H⁺(aq) + O₂(g) + 4e⁻
• Reduction2H₂O(l) + 2e⁻ → H₂(g) + 2OH⁻(aq)

3

At the cathode, either the metal ion or water is reduced, and at the anode, either the anion or water is oxidised.

4

The species discharged depends on:

• The relative values of the standard electrode potentials.
• The concentration of the ions.
• The identity of the electrode.

Section 58 of 72Products of specified electrolytes

Electrolyte	Cathode product	Anode product
Water	Hydrogen	Oxygen
Sodium chloride solution	Hydrogen	Oxygen or chlorine
Copper sulfate solution, inert electrodes	Copper	Oxygen
Copper sulfate solution, copper electrodes	Copper	Copper(II) ions, Cu2+(aq)

Section 59 of 72Influence of electrode potentials

1

Pure water electrolyses slowly because of its low ion concentration, so acid or base is added to increase the ion availability.

2

The overall products are the same in acidic and alkaline conditions, but the half-equations differ.

3

Acidic conditions (dilute sulfuric acid):

• At the cathode H⁺ (E = 0.00 V) is more easily reduced than water (E = −0.83 V), so hydrogen is discharged.
• At the anode sulfate cannot be oxidised because sulfur is already in its maximum oxidation state (+6), so water is oxidised to oxygen.

4

Alkaline conditions (dilute sodium hydroxide):

• At the cathode water (E = −0.83 V) is more easily reduced than Na⁺ (E = −2.71 V), so hydrogen is discharged.
• At the anode OH⁻ is oxidised in preference to water, so oxygen is discharged: 4OH⁻ → 2H₂O + O₂ + 4e⁻.

5

In both cases the overall reaction is 2H₂O(l) → 2H₂(g) + O₂(g).

Section 60 of 72Influence of ion concentration

1

In sodium chloride solution, the cathode product is hydrogen, from H⁺ or water.

2

At the anode the E^⦵ values for 2Cl⁻ → Cl₂ (−1.36 V) and water oxidation (−1.23 V) are similar.

3

Either chlorine or oxygen can be formed, depending on conditions.

4

At high chloride concentration (above about 25%), chlorine is the main product.

5

Overall: 2NaCl(aq) + 2H₂O(l) → 2NaOH(aq) + H₂(g) + Cl₂(g).

Section 61 of 72Influence of the electrodes

1

Active electrodes take part in the redox reactions.

2

Passive (inert) electrodes, such as platinum and carbon, do not take part.

3

Copper sulfate with passive electrodes:

• At the cathode Cu²⁺ (+0.34 V) is reduced in preference to water, so copper is deposited.
• At the anode water is oxidised to oxygen.
• Overall2CuSO₄ + 2H₂O → 2Cu + O₂ + 2H₂SO₄.

4

Copper sulfate with copper (active) electrodes:

• Copper is deposited at the cathode.
• At the anode copper dissolvesCu(s) → Cu²⁺(aq) + 2e⁻.

Section 62 of 72Purification of copper

1

The anode is impure copper and the cathode is pure copper.

2

Copper from the anode dissolves as Cu²⁺ ions, which are deposited on the cathode as pure copper.

3

Impurities fall to the bottom of the cell as anode slime, which may contain precious metals.

4

The process produces very high-grade copper for electrical wires.

Section 63 of 72More from the specification

1Controlled Oxidation of Primary Alcohols

• A primary alcohol with a terminal -CH2OH group is oxidised by a strong oxidising agent to an aldehyde under distillation,
  • (such as acidified potassium dichromate, K₂Cr₂O₇)
• or fully oxidised to a carboxylic acid under reflux.

2Oxidation of Secondary Alcohols to Ketones

• Secondary alcohols, containing an -OH group attached to a carbon bonded to two alkyl groups, are oxidised by acidified dichromate oxidising agents to form stable ketones,
• causing an orange-to-green Cr(VI) to Cr(III) colour change.

3

Electrochemical Cell Potentials and Thermodynamic Feasibility — In electrochemical cells, standard cell potential (E°_cell = E°_reduction - E°_oxidation) determines spontaneity; a positive cell potential corresponds to a negative Gibbs free energy change (ΔG° = -n F E°_cell).

4Relative Strength of Oxidising and Reducing Agents

• In redox chemistry, species with the most positive standard electrode potentials act as the most powerful oxidising agents, readily accepting electrons to undergo reduction.

5Oxidising agent definition

• Oxidising agent definition an oxidising agent causes oxidation in another substance by accepting electrons while being reduced itself.

6Halogen oxidizing power

• Halogen oxidizing power oxidising ability increases going up Group 17, enabling chlorine to displace bromide ions from aqueous solution.

7Oxidation of organic alcohols

• Oxidation of organic alcohols potassium dichromate acts as an oxidising agent to convert primary alcohols into carboxylic acids.

8Secondary alcohol

• An alcohol whose –OH carbon is bonded to two other carbons.

9Alcohol

• A compound containing –OH with a name ending in -ol.

10Oxidising agent

• A substance that oxidises another and is itself reduced.

11Reducing agent

• A substance that reduces another and is itself oxidised.

12Electrolytic cell

• A cell using electricity to drive a non-spontaneous reaction.

13Carboxylic acid

• A compound containing –COOH with a name ending in -oic acid.

14Primary alcohol

• An alcohol in which the C–OH carbon is attached to one other carbon.

Section 64 of 72Purpose of electroplating

Electroplating

Electroplating the electrolytic coating of an object with a very thin layer of metal.

Decoration

• Decoration for example gold-plated jewellery.

Corrosion prevention

• Corrosion prevention for example zinc-coated (galvanised) iron resists rusting.

Section 65 of 72Electrode set-up

1

Electroplating uses electrolysis with active electrodes.

2

The object to be plated is the cathode, where the metal is deposited.

3

The anode is usually made of the metal being deposited.

• It replenishes the metal ions in the electrolyte.
• It keeps the metal ion concentration stable.

4

The electrolyte must contain ions of the coating metal.

Section 66 of 72Reactions at the electrodes

Reactions at the electrodes

At the cathode, metal ions gain electrons and form metal atoms on the surface of the object (reduction): Mⁿ⁺ + ne⁻ → M(s).

At the soluble metal anode, metal atoms lose electrons and enter the solution (oxidation): M(s) → Mⁿ⁺ + ne⁻.

Section 67 of 72Silver-plating a spoon

1

The spoon is the cathode and impure silver is the anode.

2

A common electrolyte is sodium silver cyanide, Na[Ag(CN)₂], also called sodium dicyanoargentate(I).

3

At the anode (oxidation): Ag(s) + 2CN⁻(aq) → [Ag(CN)₂]⁻(aq) + e⁻.

4

At the cathode (reduction): [Ag(CN)₂]⁻(aq) + e⁻ → Ag(s) + 2CN⁻(aq).

Section 68 of 72Controlling the coating

Controlling the coating

The rate of deposition must be controlled to give a smooth, even layer.

The thickness of the plating is controlled by adjusting:

• The electrical current.
• The time.
• The electrolyte concentration.

More charge passed deposits more metal, but a rate of deposition that is too fast can give a rough coating.

Section 69 of 72Electroplating plastics

1

A non-conductive object such as a plastic cannot act as an electrode at first.

2

The surface is first etched or pitted.

3

Metallic particles are embedded into the surface so that it conducts electricity.

4

The conducting plastic can then be electroplated like a metal object.

Section 70 of 72More from the specification

More from the specification

Oxidising agents and reducing agents

• Oxidising agents and reducing agents chemical reducing agents donate electrons to electropositive transition metals,
  • while powerful oxidising agents gain electrons.

Carboxylic acids and functional group reactivity

• Carboxylic acids and functional group reactivity homologous carboxylic acids possess a terminal carboxyl functional group that donates protons in aqueous solutions.

Standard Hydrogen Electrode (SHE) Reference

• The standard hydrogen electrode consists of platinum foil coated in platinum black in contact with 1.0 mol dm⁻³ H⁺(aq) and hydrogen gas at 100 kPa, assigned an arbitrary standard potential of 0.00 V.

Electrochemical vs Electrolytic Cells

• Voltaic cells use spontaneous redox reactions to generate electrical power,
  • (positive E°_cell)
• whereas electrolytic cells consume external electrical energy to drive non-spontaneous chemical reactions.

Electrolytic Copper Purification and Electroplating

• In the electroplating or purification of copper, the object to be plated is attached to the negative cathode where Cu²⁺ ions are reduced,
  • (Cu²⁺ + 2e⁻ → Cu)
• while an impure copper anode undergoes oxidation.

Redox Terminology and Electrode Polarity

• In electrolysis, oxidation occurs at the positive anode and reduction occurs at the negative cathode;
  • electrons flow from the external power supply into the cathode and exit through the anode.

Standard electrode potential

• Standard electrode potential measured under standard conditions of 298 K and 100 kPa relative to the standard hydrogen electrode.

Copper and silver cell

• Copper and silver cell linking a standard copper half-cell to a silver electrode yields a positive cell potential.

Oxidising agent

• Oxidising agent a substance that gains electrons and is reduced, while causing another species to be oxidised.

Secondary alcohol

• An alcohol whose –OH carbon is bonded to two other carbons.

Sodium chloride

• Compound of sodium and chlorine; common salt.

Half-cell

• One electrode and its solution, in which a redox equilibrium is set up.

Oxidising agent

• A substance that oxidises another and is itself reduced.

Electrode potential

• A measure of how easily a species is reduced, relative to a reference half-cell.

Reducing agent

• A substance that reduces another and is itself oxidised.

Carboxylic acid

• A compound containing –COOH with a name ending in -oic acid.

Primary alcohol

• An alcohol in which the C–OH carbon is attached to one other carbon.

Spontaneous reaction

• A reaction with ΔG < 0.

Hydrogen

• A gas that gives a squeaky pop with a lighted splint.

Alcohol

• A compound containing –OH with a name ending in -ol.

Section 71 of 72Key terms

Term	Meaning
Oxidising agent	A species that accepts electrons and is reduced.
Reducing agent	A species that donates electrons and is oxidised.
Voltaic cell	A cell generating electricity from a spontaneous redox reaction.
Electrolytic cell	A cell using electricity to drive a non-spontaneous reaction.
Salt bridge	A link of inert electrolyte completing the circuit.
Standard hydrogen electrode	The reference half-cell with E° = 0.00 V.
Electroplating	Coating an object with a metal by electrolysis.

Term	Meaning
Oxidising agent	A species that oxidises another by accepting electrons and is itself reduced.
Reducing agent	A species that reduces another by donating electrons and is itself oxidised.
Displacement reaction	A reaction in which a more reactive element displaces a less reactive one from its compound.
Reactivity series	A ranking of metals by how easily they lose electrons.
Spectator ion	An ion that is unchanged in a reaction.
Salt	An ionic compound formed when the hydrogen of an acid is replaced by a metal or other positive ion.

Term	Meaning
Voltaic cell	A cell in which a spontaneous redox reaction produces an electric current.
Half-cell	A metal electrode in a solution of its own ions.
Electrode potential	The potential difference between an electrode and its solution.
Salt bridge	A link containing mobile ions that completes the circuit.
EMF	Electromotive force, the potential difference generated by a cell.
Phase boundary	The interface between a solid electrode and its solution, shown by a single vertical line.

Term	Meaning
Fuel cell	An electrochemical cell in which a fuel is oxidised at one electrode and oxygen reduced at the other.
Electrolytic cell	A cell that uses electrical energy to drive a non-spontaneous reaction.
Secondary cell	A rechargeable cell whose reaction can be reversed by an applied voltage.
Lead–acid battery	A rechargeable battery with Pb and PbO2 electrodes in sulfuric acid.
Anode	The electrode where oxidation takes place.
Cathode	The electrode where reduction takes place.

Term	Meaning
Memory effect	Gradual loss of capacity when a NiCad cell is recharged repeatedly without being fully discharged.
Electrolysis	The splitting of an ionic compound by an electric current.
Electrolyte	The substance that conducts the current and is split during electrolysis.
Primary alcohol	An alcohol whose –OH carbon is bonded to one other carbon.
Secondary alcohol	An alcohol whose –OH carbon is bonded to two other carbons.
Tertiary alcohol	An alcohol whose –OH carbon is bonded to three other carbons.

Term	Meaning
LiAlH4	A strong reducing agent used in dry ether that reduces carboxylic acids, aldehydes and ketones.
NaBH4	A milder reducing agent that reduces aldehydes and ketones but not carboxylic acids.
Hydrogenation	The addition of hydrogen across a C=C or C≡C bond using a nickel catalyst.
Standard hydrogen electrode	The reference half-cell of H2 at 100 kPa and H+ at 1.00 mol dm−3 on platinum, with E = 0.00 V.
Standard electrode potential	The EMF of a cell with the SHE on the left and the standard half-cell on the right.
Hydride ion	The H− ion supplied by LiAlH4 and NaBH4, which acts as a nucleophile.

Term	Meaning
EMF	The potential difference of a cell, E(right) − E(left).
Standard electrode potential	The EMF of a half-cell measured against the standard hydrogen electrode under standard conditions.
Standard conditions	1.0 mol dm−3, 100 kPa and 298.15 K.
Electrochemical series	Half-reactions listed by standard electrode potential from most positive to most negative.
Spontaneous reaction	A reaction with a positive EMF under the conditions considered.

Term	Meaning
Gibbs energy change (ΔG)	The free energy change of a reaction, negative when spontaneous.
Faraday constant	9.65 × 104 C mol−1, the charge carried by one mole of electrons.
Inert (passive) electrode	An electrode, such as platinum or carbon, that does not take part in the reaction.
Active electrode	An electrode that takes part in the redox reactions.
Anode slime	The impurities that fall to the bottom of a copper refining cell.

Term	Meaning
Electroplating	Coating an object with a thin metal layer by electrolysis.
Cathode (in plating)	The object being plated, where metal ions are reduced.
Active anode	An anode made of the plating metal that dissolves to replace the ions removed.
Galvanised iron	Iron coated with zinc to resist rusting.
Dicyanoargentate(I)	The complex ion [Ag(CN)2]−, present in the silver plating electrolyte.

Section 72 of 72Exam technique

1

Putting oxidation at the cathode — oxidation is always at the anode.

2

Saying the anode is always negative — it is negative in voltaic cells and positive in electrolytic cells.

3

Writing E°_cell as left minus right.

4

Using ΔG° in kJ and F in C mol⁻¹ without converting to the same energy unit.

5

Assuming a reaction with positive E° will always be observed.

1

Mixing up oxidising and reducing agents — the reducing agent is oxidised and the oxidising agent is reduced.

2

Saying halogen reactivity increases down the group — oxidising power decreases down Group 17.

3

Including spectator ions in ionic equations.

4

Expecting copper to react with dilute acid — it is below hydrogen in the series.

5

Reading displacement results backwards — the free metal must be the more reactive one.

1

Putting oxidation at the positive electrode in a voltaic cell — oxidation is at the negative electrode (anode).

2

Forgetting that electrons flow through the wire while ions flow through the salt bridge.

3

Choosing salt bridge ions that form a precipitate with the half-cell solutions.

4

Writing the species of higher oxidation state next to the electrode instead of next to the salt bridge.

5

Believing the absolute value of a half-cell potential can be measured.

1

Placing oxidation at the cathode — oxidation is always at the anode and reduction is always at the cathode.

2

Saying the anode is always negative — it is negative in a voltaic cell but positive in an electrolytic cell.

3

Thinking a fuel cell stores energy — it converts the energy of a continuously supplied fuel.

4

Forgetting that the hydrogen fuel cell equation depends on the electrolyte — in alkali the half-equations use OH⁻.

1

Forgetting that primary alcohols oxidise to acids under reflux but give aldehydes only when the product is distilled off.

2

Saying tertiary alcohols oxidise — they have no hydrogen on the carbon bonded to –OH.

3

Saying the NiCad memory effect is a loss of voltage — it is a gradual loss of capacity.

4

Predicting a metal at the anode in the electrolysis of a molten binary salt — metals form at the cathode.

1

Using NaBH₄ to reduce a carboxylic acid — only LiAlH₄ is strong enough.

2

Trying to stop the LiAlH₄ reduction of a carboxylic acid at the aldehyde — it continues to the primary alcohol.

3

Writing electrode half-equations as oxidations when quoting standard potentials — they are always reductions.

4

Saying a more negative E means a stronger oxidising agent — it means a stronger reducing agent.

5

Believing the SHE is easy to use routinely — it is slow, not portable and hard to keep at constant pressure.

1

Subtracting in the wrong order — EMF is E(right) − E(left), or E(reduction) − E(oxidation).

2

Multiplying E by the coefficient when a half-equation is scaled — E does not depend on the amount of substance.

3

Saying a positive E means the reaction will visibly happen — it shows only feasibility, not rate.

4

Using non-standard conditions and still calling the value a standard electrode potential.

5

Placing the more positive half-cell on the left — the more negative half-cell (oxidation) is on the left.

1

Using ΔG = nFE without the minus sign.

2

Forgetting to convert ΔG from J mol⁻¹ to kJ mol⁻¹, or using the wrong units for F.

3

Assuming the metal ion is always discharged at the cathode — in aqueous solution the more easily reduced species is discharged, so hydrogen forms for reactive metals.

4

Saying sulfate ions are oxidised at the anode — sulfur is already at +6, so water is oxidised.

5

Treating active and inert electrodes as the same — an active copper anode dissolves.

1

Putting the object to be plated at the anode — it is the cathode.

2

Using an inert anode and expecting the electrolyte concentration to stay constant — an anode of the plating metal replenishes the ions.

3

Forgetting that the silver cathode reaction involves the complex [Ag(CN)₂]− in this electrolyte.

4

Plating a plastic directly — its surface must first be made conductive.