Redox Reactions

Year 12 Chemistry

Chapter 6.1, 6.2 and 6.3 Redox Reactions

• Teacher: Mrs. Pickert

Metal Activity Series

• The metal activity series is a tool for predicting:

  • Which metals will corrode.

  • Methods for preventing corrosion.

  • Extraction methods for metals from ores.

• Typical metal activity series ranking (from most reactive to least reactive):

  • Sodium

  • Magnesium

  • Aluminium

  • Chromium (carbon)

  • Zinc

  • Iron

  • Nickel

  • Tin

  • Lead (hydrogen)

  • Copper

  • Silver

  • Gold

Electrochemistry: Types of Redox Reactions

• Types of redox reactions include:

  • Respiration reactions (energy source for living organisms)

  • Photosynthesis in green plants

  • Combustion of fuels (e.g., cars)

  • Burnt coal in electricity generation

  • Chlorine use for swimming pool disinfection

  • Manufacturing of explosives

  • Electrolysis for producing chemicals

  • Production and use of fertilizers

Types of REDOX Reactions

• Displacement Reactions:

  • Involve a metal element, often referred to as single displacement reactions.

  • Practical example: Practical 6.1

• Combustion Reactions:

  • Reactions with oxygen.

• Corrosion Reactions:

  • Occur between metals and atmospheric conditions.

  • More stable as ions than as solids (based on metal reactivity series).

Reduction and Oxidation (OIL RIG)

• Relates to Bohr’s model of electrons.

• Redox reactions involve the transfer of electrons:

  • Oxidation: Reactant loses an electron.

  • Reduction: Reactant gains an electron.

  • Memory key: OIL RIG (Oxidation Is Loss, Reduction Is Gain).

Understanding OIL RIG

• Historically, reactions were noted where products had lighter masses, which linked to reduction.

• Scientific understanding evolved, revealing the electron transfer dynamics.

Redox Reaction Dynamics

• When electrons transfer, they create REDOX reactions:

  • "Red" signifies reduction.

  • "Ox" signifies oxidation.

  • All redox processes occur concurrently; oxidation of one substance correlates with reduction of another.

Identifying Agents in Redox Reactions

• An oxidised reactant is termed a reducing agent (reductant) and causes reduction in another.

• A reduced reactant is termed an oxidising agent (oxidant) and causes oxidation in another.

Predicting Redox Reactions via the Periodic Table

• Elements likely to undergo oxidation or reduction can be predicted:

  • Low Ionisation Energy means an element is likely to be oxidised as it loses valence electrons easily.

  • High ionisation energy indicates a higher likelihood to be reduced.

• Metals are typically stronger reducing agents due to low ionisation energy.

Additional Prediction Factors

• An element with high electronegativity is more likely to be reduced, gaining electrons to achieve a stable state.

• Non-metals can be stronger oxidising agents.

Study Activities

• Take a brief break.

• Write summary notes for this section and complete the 6.1 worksheets attached to Toddle.

• Review 6.2 slides, write summary notes, and complete those worksheets as well.

6.2 - Oxidation Numbers in Redox Reactions

• Oxidation numbers help simplify understanding oxidation vs. reduction:

  • Reflect total electrons gained or released.

  • An increase in oxidation state indicates oxidation.

  • A decrease indicates reduction.

Rules for Assigning Oxidation Numbers

1. Any atom in its elemental form has an oxidation number of zero (e.g., Na, O₂, Pb, S, Ne = 0).

2. The oxidation number of a monatomic ion equals its charge (e.g., Fe³⁺ = +3, S²⁻ = -2).

Group Oxidation States

• Group I: +1 (None)

• Group II: +2 (None)

3. For neutral compounds, total oxidation number equals zero (e.g., H₂SO₄: H = +1, O = -2, S = +6).

4. For polyatomic ions, total oxidation number equals the ion's charge (e.g., CO₃²⁻: C = +4).

6.3 - Redox Half Equations

• Half equations represent oxidation or reduction. They can be combined to create a balanced redox equation.

Steps for Balancing Redox Equations

1. Identify reactants and products.

2. Balance all atoms other than O and H.

3. Add H₂O to balance for O.

4. Add H⁺ to balance for H.

5. Add electrons (e⁻) to balance charge (to the more positive side).

6. Equalize the number of electrons in both half equations (multiplication required if necessary).

7. Add the half equations, cancel electrons and common species.