Mission 1: When iron rusts, does it gain or lose electrons?

Four ways to spot oxidation

• The two always happen together: a redox reaction.
• The electron and oxidation-state definitions work for every redox reaction.

Oxidation states: the rules

• An uncombined element is 0: Fe, O₂, Cl₂.
• A monatomic ion's state is its charge: Fe³⁺ is +3.
• Usually O is −2 and H is +1. F is −1 in all its compounds.
• In compounds, group 1 metals are +1; group 2 metals, +2.
• The states add up to the charge on the species: 0 for a compound.
• Write the sign first: +3, not 3+.

Worked examples

• SO₄²⁻: x + 4(−2) = −2, so S is +6.
• H₂S: 2(+1) + x = 0, so S is −2.
• P₂O₇⁴⁻: 2x + 7(−2) = −4, so 2x = +10 and each P is +5.
• Sulfur runs from −2 to +6: most non-metals have several states.

Oxidation states in names

• A Roman numeral in a name gives the oxidation state of the element before it.
• FeCl₂ is iron(II) chloride; FeCl₃ is iron(III) chloride.
• Transition elements have variable states: Cu₂O is copper(I) oxide; CuO is copper(II) oxide.
• Oxyanions too: ClO₃⁻ is chlorate(V); SO₃²⁻ is sulfate(IV); SO₄²⁻ is sulfate(VI).

Oxidizing and reducing agents

• The oxidizing agent takes electrons, so it is itself reduced.
• The reducing agent gives electrons, so it is itself oxidized.
• Agents are reactants, never products.
• Zn + Cu²⁺ → Zn²⁺ + Cu. Zn: 0 → +2, oxidized: the reducing agent.
• Cu²⁺: +2 → 0, reduced: the oxidizing agent.

Half-equations

• A half-equation shows the oxidation or the reduction alone, with the electrons written in.
• Oxidation: electrons are products. Zn → Zn²⁺ + 2e⁻.
• Reduction: electrons are reactants. Cl₂ + 2e⁻ → 2Cl⁻.
• Both atoms and charge must balance.

Balancing in acid: worked example

Cr₂O₇²⁻ → Cr³⁺ in acidic solution:

• Balance Cr: Cr₂O₇²⁻ → 2Cr³⁺.
• Balance O with H₂O: add 7H₂O on the right.
• Balance H with H⁺: add 14H⁺ on the left.
• Balance charge with e⁻: left +12, right +6, so add 6e⁻ on the left.
• Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O. Check: 2 Cr each fall by 3, so 6e⁻.

Combining half-equations

• Scale so the electrons lost equal the electrons gained, then add and cancel them.
• Oxidation: 2I⁻ → I₂ + 2e⁻. Multiply by 3 to give 6e⁻.
• Add to the dichromate half-equation; cancel 6e⁻:
• Cr₂O₇²⁻ + 14H⁺ + 6I⁻ → 2Cr³⁺ + 7H₂O + 3I₂
• Check charge: left −2 + 14 − 6 = +6; right 2(+3) = +6.

Mission 2: Why does zinc fizz in acid when a gold ring doesn't?

Metals: the trend in ease of oxidation

• Metals react by losing electrons: they are oxidized, so they act as reducing agents.
• Down a group, oxidation gets easier: outer electrons are further out and more shielded.
• Towards the left of a period it gets easier too: Na > Mg > Al.
• So K is oxidized more readily than Na, and Na more readily than Mg.

Halogens: the opposite trend

• Halogens react by gaining electrons: X₂ + 2e⁻ → 2X⁻. They are oxidizing agents.
• Reduction gets easier up group 17: F₂ > Cl₂ > Br₂ > I₂.
• A smaller atom holds an added electron closer to the nucleus, with less shielding.
• A halogen oxidizes the halide ions below it: Cl₂ + 2I⁻ → 2Cl⁻ + I₂.
• Br₂ with Cl⁻: no reaction.

Metal displacement reactions

• A metal oxidized more readily displaces a less readily oxidized metal from its ions.
• Zn + Cu²⁺ → Zn²⁺ + Cu: copper coats the zinc; the blue colour fades.
• Cu + Zn²⁺: no reaction.
• Exams give you the data. Rank the metals by which displaces which.
• The ion of the least readily oxidized metal is the strongest oxidizing agent in the set.

Reactive metals with dilute acids

• A reactive metal reduces H⁺ to hydrogen gas and is itself oxidized.
• Metal + acid → salt + hydrogen. This is redox, not neutralization.
• Mg + 2HCl → MgCl₂ + H₂. Ionic: Mg + 2H⁺ → Mg²⁺ + H₂.
• Zn + H₂SO₄ → ZnSO₄ + H₂.
• Each H₂ needs 2e⁻: count the electrons the metal loses.
• Metals oxidized less readily than hydrogen, such as Cu, Ag and Au, do not react.

Mission 3: Why can you recharge a phone battery but not a standard AA battery?

Anode and cathode

• Anode: the electrode where oxidation happens.
• Cathode: the electrode where reduction happens.
• This holds in every electrochemical cell. Their signs depend on the type of cell.

Building a primary (voltaic) cell

• A voltaic cell turns the chemical energy of a spontaneous redox reaction into electrical energy.
• Two half-cells: zinc in Zn²⁺(aq); copper in Cu²⁺(aq).
• Joined by an external circuit (wire) and a salt bridge, e.g. KNO₃(aq).
• Zinc is oxidized more readily. Anode: Zn → Zn²⁺ + 2e⁻.
• Cathode: Cu²⁺ + 2e⁻ → Cu.

Which way do electrons and ions go?

• Electrons flow through the wire, from the anode (Zn) to the cathode (Cu).
• So in a voltaic cell the anode is negative; the cathode is positive.
• Electrons never pass through the solutions or the salt bridge.
• Salt bridge: anions move to the anode half-cell, where Zn²⁺ is building up.
• Cations move to the cathode half-cell, where Cu²⁺ is being used up.
• This keeps both solutions neutral, so the current keeps flowing.

Primary, secondary and fuel cells

• Primary cell: single use. Its reaction cannot practically be reversed.
• Secondary cell: rechargeable. A charger's current drives the reaction backwards.
• Charging reverses each whole half-equation: products become reactants, electrons included.
• Fuel cell: fuel (e.g. H₂) and O₂ are fed in continuously; it never needs recharging.
• A cell stores chemical energy, not electricity.

Charging, worked: the lithium-ion cell

• Discharge, negative electrode (oxidation): LiC₆ → C₆ + Li⁺ + e⁻.
• Discharge, positive electrode (reduction, Co +4 → +3): CoO₂ + Li⁺ + e⁻ → LiCoO₂.
• Charging reverses both: C₆ + Li⁺ + e⁻ → LiC₆ and LiCoO₂ → CoO₂ + Li⁺ + e⁻.
• During charging, Li⁺ ions move back into the graphite.

Trade-offs

Mission 4: Why does a bulb light only once solid lead(II) bromide has melted?

The electrolytic cell

• An electrolytic cell uses electrical energy to drive a non-spontaneous redox reaction.
• A DC power supply is connected to two electrodes in an electrolyte.
• Anode: joined to the positive terminal. Cathode: joined to the negative terminal.
• Oxidation is still at the anode; reduction at the cathode.

Voltaic and electrolytic cells side by side

How the current is carried

• In the wires and electrodes: electrons.
• In the electrolyte: ions moving. Cations go to the cathode; anions go to the anode.
• In a solid ionic compound the ions are held in the lattice: no conduction.
• Molten (or dissolved), the ions are free to move. No electrons pass through the melt.

Products from a molten salt: worked example

Molten lead(II) bromide, PbBr₂(l), with inert electrodes:

• Cathode (−): Pb²⁺ + 2e⁻ → Pb(l). Molten lead collects.
• Anode (+): 2Br⁻ → Br₂(g) + 2e⁻. Brown vapour.
• Overall: PbBr₂(l) → Pb(l) + Br₂(g).
• A molten binary salt gives the metal at the cathode and the non-metal at the anode.

Mission 5: How does a breath tester detect alcohol?

Primary, secondary, tertiary alcohols

• Find the carbon bonded to the –OH. Count the carbons attached to it.
• Primary: one (methanol: none). Ethanol, CH₃CH₂OH.
• Secondary: two. Propan-2-ol, CH₃CH(OH)CH₃.
• Tertiary: three. 2-methylbutan-2-ol, (CH₃)₂C(OH)CH₂CH₃.
• Primary and secondary alcohols have an H on that carbon; tertiary alcohols have none.

What oxidation gives

• [O] stands for the oxidizing agent; you need not name it.
• Primary → aldehyde → carboxylic acid, in two steps:
• CH₃CH₂OH + [O] → CH₃CHO + H₂O, then CH₃CHO + [O] → CH₃COOH.
• Secondary → ketone: CH₃CH(OH)CH₃ + [O] → CH₃COCH₃ + H₂O.
• Tertiary: not oxidized under the same conditions; there is no H on the C–OH carbon.

Distillation or reflux?

• The aldehyde boils lowest: it has no O–H, so no hydrogen bonding between its molecules.
• Distillation: the aldehyde boils off as it forms, so it is not oxidized further.
• Reflux: a vertical condenser returns vapour to the flask, so the aldehyde stays and becomes the acid.
• For the aldehyde: distil, with the alcohol in excess. For the acid: reflux, with excess oxidizing agent.

Reduction: adding hydrogen to C=O

• Reducing a carbonyl compound or acid is a gain of hydrogen. [H] stands for the reducing agent.
• The reducing agent supplies hydride ions, H⁻: hydrogen with two electrons, which adds to the C=O carbon.
• Acid → aldehyde: CH₃COOH + 2[H] → CH₃CHO + H₂O.
• Aldehyde → primary alcohol: CH₃CHO + 2[H] → CH₃CH₂OH.
• Overall: CH₃COOH + 4[H] → CH₃CH₂OH + H₂O.
• Ketone → secondary alcohol: CH₃COCH₃ + 2[H] → CH₃CH(OH)CH₃.

Hydrogenation

• Hydrogenation: H₂ adds across C=C or C≡C, over a heated nickel catalyst.
• It lowers the degree of unsaturation: a reduction (gain of hydrogen).
• One H₂ per π bond: C=C takes 1 H₂; C≡C takes 2 H₂. The σ bond stays.
• CH₃CH=CH₂ + H₂ → CH₃CH₂CH₃.
• CH₃C≡CCH₃ + 2H₂ → CH₃CH₂CH₂CH₃, via but-2-ene.

The big ideas

• Oxidation is loss of electrons, a rise in oxidation state. The oxidizing agent is itself reduced.
• In acid, balance O with H₂O, H with H⁺, then charge with e⁻.
• Metals are oxidized more readily down a group; halogens are reduced more readily up group 17.
• Reactive metal + dilute acid → salt + hydrogen.
• Anode = oxidation. Voltaic: anode negative. Electrolytic: anode positive. Electrons travel only in wires.
• Molten salt: ions carry the current; metal at the cathode, non-metal at the anode.
• Primary alcohol → aldehyde → acid; secondary → ketone; tertiary is not oxidized. Reduction and hydrogenation add hydrogen.