Oxidation and Reduction (chpt 14)

Practise balancing redox equations and finding oxidation numbers

Examples of oxidation and reduction

• 2Mg + O₂ → 2Mg²⁺O²⁻ - Mg loses 2 electrons and so is oxidised

• Cu²⁺O²⁻ + H₂ → Cu + H₂O - Cu gains two electrons and so is reduced

• Zn + Cu²⁺ → Zn²⁺ + Cu - Zn loses two electrons and Cu gains two electrons, therefore Zn is oxidised and Cu is reduced

Oxidation

When an element loses electrons (OIL)

Reduction

When an element gains electrons (RIG)

Oxidising agent

Substance that brings about oxidation in other substance. As a characteristic, the oxidising agent is always reduced as the ion has already accepted some electrons

Reducing agent

Substance that brings about reduction in other substances. As a characteristic, the reducing agent is always oxidised as the ion has already given away some electrons

Solving oxidation reduction equations

1. Assign oxidation numbers to all atoms

2. Note which elements change oxidation number (ie identify compounds involved in reaction)

3. Show number of electrons lost and gained in reaction. This becomes our ratio

4. Multiply by ratio (redox by oxidation and vice versa)

5. Balance remaining elements and double check for atoms and their charges

Oxidation Numbers

The charge that an atom has or appears to have when electrons are distributed according to certain rules

Changes: reduces (decreases)

Oxidation-Reduction Reactions - Halogens as Oxidising Agents (Experiment)

1. Potassium bromide solution + chlorine water (equal amounts) → colourless to yellow → Br₂ formed as Cl₂ oxidised 2Br⁻ into Br₂: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂

2. Potassium iodide solution + chlorine water → Colourless to reddish-brown → I₂ formed as Cl₂ oxidised 2I⁻ to I₂: Cl₂ + 2I⁻ → 2Cl⁻ + I₂

3. Potassium iodide solution + bromine water → colourless to reddish-brown → I₂ formed as Br₂ oxidised: Br₂ + 2I⁻ → 2Br⁻ + I₂

4. Sodium sulfite + chlorine water → tested for presence of sulfate solution by adding barium chloride solution → + dilute HCl solution → white perecipitate when barium added, does not dissolve in dilute HCl solution (sulfate ion present) → Cl oxidised by sulfite ion: Cl₂ + SO₃²⁻ + H₂O → 2Cl⁻ + SO₄²⁻ + 2H⁺

5. Iron (II) sulfate solution + chlorine water + dilute sodium hydroxide solution → no change w chlorine water, greenish-brown precipitate w sodium hydroxide → precipitate is Fe₂O₃ and confirms Fe³⁺ ions: Cl₂ + 2Fe²⁺ → 2Fe³⁺ + 2Cl⁻

Oxidation, Oxidising agent

Oxidation-Reduction Reactions - Displacement Reactions of Metals (Experiment)

1. Granulated/powdered zinc + copper sulfate solution (left to stand) → shiny zinc metal turns dark brown → zinc coated w copper metal as Zn has greater tendency to lose electrons than copper so zinc forces Cu²⁺ ions to accept electrons: Zn + Cu²⁺ → Zn²⁺ + Cu↓

2. Magnesium ribbon cleared w sandpaper until shiny + copper sulfate solution → magnesium metal coated w dark brown substance, blue solution lighter in colour, eventually colourless → Magnesium displaces copper metal out of solution, as megnesium has greater tendency to lose electrons than copper and forces Cu²⁺ ions to accept electrons: Mg + Cu²⁺ → Mg²⁺ + Cu

Oxidation, Reduction

Rules for Assigning Oxidation Numbers

1. Oxidation number of any uncombined element is zero, eg. Zn, Fe, O₂, Cl₂

2. Oxidation number of an ion of an element is the same as its charge, eg. Fe²⁺ is +2, Cl⁻ is -1

3. Sum of oxidation numbers of all elements in a compound must add to zero

4. Oxygen always has an oxidation number of -2 except in peroxides (ON is -1) and in compound OF₂ (ON is +2), eg. Mg²⁺O²⁻ ON is -2, but in H₂O₂ it is -2 (this is to match with hydrogens and fluorines)

5. Hydrogen is always +1 in compounds except in metal hydrides (-1), eg. NH₃ ON is +1, but in NaH ON is -1 to match sodium

6. Halogens always have an oxidation number of -1 in compounds unless bonded to a more electronegative atom, eg. fluorine (most electronegative) always -1, chlorine -1 in Cl₂O but is +7 in Cl₂O₇ as O is -2 (more electronegative is usually written on write, eg. OF₂)

7. The sum of the oxidation numbers of all elements in a complex ion must equal the charge on the ion, eg. NO₃⁻ must equal to -1, not 0

8. Oxidation number of an element can be in fractions or even 0 (valency is always a whole number)

Oxidation and reduction in terms of oxidation numbers

• When oxidised: oxidation number increases

• When reduced: oxidation number decreases