IB Chem topic 9: Redox processes

Reduction reactions

reactants gain electrons

Oxidation reactions

reactants lose electrons

Lavoisier’s explanation of redox reactions

Oxidation = gain of oxygen, loss of hydrogen

Reduction = loss of oxygen, gain of hydrogen

Oxidation states

• measure of the electron control/possession an atom has relative to the atom in the pure element state

• assigned to each atom to keep track of electrons

• signifies electron control gained or lost

Oxidation state of atoms in free element form

ZERO

Oxidation state of simple ions

same as the charge on the ion

Oxidation states of all atoms in a neutral compound add to __

ZERO

Oxidation states of all atoms in a polyatomic ion add to ______

the charge of the polyatomic ion

usual oxidation state for an element is

the same as the charge on its most common ion (many exceptions exist; most main group non metals, elements at bottom of group 14, and transition metals have variable oxidation states)

Li, Na, K usual oxidation state

+1

Mg, Ca usual oxidation state

+2

F usual oxidation state

-1

O usual oxidation state (and exceptions)

-2, in peroxides (H₂O₂) it is -1, in OF₂ it is +2

H usual oxidation state (and exceptions)

+1, -1 in metal hydrides such as NaH

Cl usual oxidation state

-1, different when combined with O or F

PRACTISE :) Assign oxidation states to each atom

(Co(NH₃)₆)³⁺

Co: +3

N: -3

H: +1

Greater oxidation states indicate…

the more the atom has lost control over electrons, the more oxidized it is

More negative oxidation states indicate…

the more the atom gains control over electrons, the more reduced it is

Half reactions

• Splits overall redox into distinct reduction and oxidation reactions

• The two half reactions occur at the same time

• They must have an equal number of electrons on both half-reactions to cancel when added together

Steps for constructing equations with half reactions

1. Assign oxidation states to determine which atoms are being oxidized and which are being reduced

2. Write half equations

   1. Balance atoms other than H and O

   2. Balance each half-reaction for O by adding H₂O

   3. Balance the H by adding H⁺

   4. Balance the charges by adding electrons to side with positive charge

3. Equalize the number of electrons by multiplying half reactions appropriately

4. Add two half-reactions together, cancelling same on both sides

Oxidizing agent

• the reactant that accepts electrons, as it brings about oxidation of other reactant

• It itself reduces

Reducing agent

• The reactant that supplies electrons, as it brings about reduction in the other reactanbt

• It itself oxidizes

Explanation of the activity series of metals in relation to reducing power

• By comparing displacement reactions between different combinations of metals and their ions, we build up a list of relative strengths of metals as reducing agents

• Allows us to predict if a particular redox reaction is feasible

• More reactive metals are stronger reducing agents, meaning they oxidize most readily

Will this reaction occur?

ZnCl₂ + 2Ag → 2AsCl + Zn

No twink! silver is a weaker reducing agent

STUDY BREAK

go make a snack or a drink

Explanation of the activity series of halogens in relation to oxidizing power

More reactive halogens (high EN) most readily become reduced and are strong oxidizing agents

Equation of redox titration using acidified potassium permanganate to analyze iron

5Fe²⁺ + MnO₄^- + 8H⁺ → 5Fe³⁺ + Mn²⁺ + 4H₂O

(the K is not considered)

Description of Iodine - Thiosulfate redox titration

• Uses an oxidant to react with excess iodine ions to form iodine

• 2I^- + Oxidizing agent → I₂ + Reduced product

• The liberated iodine, I₂ is then titrated with sodium thiosulfate (Na₂S₂O₃) using starch as an indicator

• 2Na₂SO₄ + I₂ → 2NaI + Na₂S₄O₆

Biological oxygen demand

• measure of the degree of pollution of a body of water

• the amount of oxygen used to decompose organic matter over a specified time period (usually 5 days)

Ratio of water to S₂O₃^2- in the Winkler method

1 mol: 4 mol

Description of Galvanic/Voltaic cells

• Generate electricity from spontaneous redox reactions —converts chemical energy to electrical energy

• Half reactions are separated into half-cells, and electrons are allowed to flower between them through an external circuit

• In a half cell, a strip of metal is placed into a solution of its own ions

Electrode potential in Galvanic cells

charge separation formed between the metal and its ions in solution

The connection between half cells

• an external wire allows electrons to spontaneously flow from the half cell where oxidation is occurring ANODE (eg. zinc) to half cell where reduction is occurring CATHODE (eg. copper)

• a salt bridge completes the circuit: a glass tube or strip of absorptive paper containing aqueous solutions of ions. Allows for the free flow of ions between half cells and neutralizes/balances charge

  • KNO₃ often used

Cell diagram convention of a zinc and copper galvanic cell

Zn_(s) | Zn²⁺_(aq) || Cu²⁺_(aq) | Cu_(s)

→ the direction of electron flow →

Description of electrolytic cells

• an external source of electricity drives a non-spontaneous redox reaction

• Electrolytic cells use an external source of electrical energy to bring about a redox reaction (ie. converts electrical energy to chemical energy

• Electrodes are immersed in an electrolyte (solution of ionic compounds) and connected to a power supply, they are described as inert, but must not touch

• Electric wires connect the electrodes to the power supply

Describe the electrolysis of a molten salt

• When the electrolyte is a molten salt, the only ions present are those from the compound itself as there is no solvent

• One ion migrates to each electrode