Chapter 11 - Redox

Redox Reaction

A chemical process where electrons move from one substance to another, involving oxidation (the loss of electrons) and reduction (the gain of electrons)

Oxidation

A chemical process where an atom, ion, or molecule loses electrons, gains oxygen, or loses hydrogen, causing its oxidation number to increase

Reduction

A process where an atom, ion, or molecule gains electrons, causing its oxidation number to decrease

Reducing agents (Reductants)

A substance that donates electrons and gets oxidized during a chemical redox reaction

Oxidising agents (Oxidants)

A chemical substance that gains electrons, removes hydrogen, or transfers oxygen atoms to another substance

Oxidation Number / Oxidation State

A number assigned to an atom representing its hypothetical charge if all chemical bonds were completely ionic. It serves as a tool to track electron movement in redox reactions, helping to identify oxidation (loss of electrons, shown by an increase in the number) and reduction (gain of electrons, shown by a decrease in the number).

Oxidation Number Rules

Elements: Oxidation Number = 0
Ions: Oxidation Number = Charge of Individual Ion (adding the ONs of each ion equals the overall charge of the compound)
Hydrogen Covalent Compounds: Oxidation Number (Hydrogen) = +1 (adding the ONs of each atom equals zero)
Oxygen Covalent Compounds: Oxidation Number (Oxygen) = -2 (adding the ONs of each atom equals zero)
Other Covalent Compound: Oxidation Number for the most electronegative element is the charge that will form if it becomes an ion. (adding the ONs of each atom equals zero)

Half Equation

A chemical equation that shows either oxidation or reduction and shows movement of electrons

Half Equations Steps

1. Balance the key element/s.

2. Balance charge by adding electrons to the more positive side of the equation

3. Add states of matter (note, electrons do not have states)

Combining Half Equations Steps

1. Ensure that both half-reactions have the same number of electrons (e-). This may require multiplying a half equation.

2. Add all the reactants on the left-hand side.

3. Add all the products on the right-hand side.

4. Cancel the e- as they should be the same on both sides.

5. Remove spectator ions if present and combine remaining ions to make ionic substances.

6. Add states of matter.

Complex Half Equations Steps (KOHES)

K = Balance the key element (undergoing oxidation or reduction).

O = Balance oxygen by adding H2O to the opposite side of the equation.

H = Balance hydrogen by adding H+ to the opposite side of the equation.

E = Balance the charge on both sides of the equation by adding electrons to the more positive side.

S = Add states.

Combining KOHES Half Equations Steps

1. Ensure that both half reactions have the same number of electrons (e-).

2. Combine the half reactions and cancel out species which are the same on both sides.

3. The electrons cancel out, though note that water (H2O) and hydrogen ions (H+) are on both sides so can also be simplified.

Reactivity Series of Metals

An empirical ranking of metals listed in order of their general chemical reactivity from highest to lowest

• Strongest oxidising agent and weakest reducing agent is at the top

• Weakest oxidising agent and strongest reducing agent is at the bottom

Electrochemical Series

An arrangement of elements (metals and non-metals) based on their Standard Electrode Potential measured in volts under standard conditions relative to a standard hydrogen electrode.

• Strongest oxidising agent and weakest reducing agent is at the top

• Weakest oxidising agent and strongest reducing agent is at the bottom

Metal Displacement Reactions

A chemical process where a more reactive metal takes the place of a less reactive metal from its compound or salt solution

• Electrons are transferred from the metal solid (oxidation) to the metal cation (reduction).

Galvanic Cell / Voltaic Cell

An electrochemical device that converts stored chemical energy into usable electrical energy through a spontaneous oxidation-reduction (redox) reaction where where oxidation occurs at the anode and reduction occurs at the cathode, causing electrons to flow through an external circuit to generate electricity

Half-cell

One of the two separate compartments in a galvanic cell where an independent oxidation or reduction reaction occurs

Anode

The negatively charged site of oxidation in a galvanic cell

Cathode

The positively charged site of reduction in a galvanic cell

Wire

External metallic conductor which connects the anode and the cathode to allow electrons from the anode to the cathode to flow, completing the electrical circuit and generating electricity

Salt Bridge

A laboratory device containing an electrolyte solution that links the oxidation and reduction half-cells in a galvanic cell to maintain electrical neutrality and complete the internal circuit

Conjugate Redox Pairs

An electron-donating species and its corresponding electron-accepting species

Corrosion

The natural process where materials (mostly metals) break down through a chemical or electrochemical reaction with substances in their environment, such as oxygen, water, or acids

Dry Corrosion

The direct deterioration and oxidation of a metal through a chemical reaction with atmospheric gases (such as oxygen, sulfur dioxide, or chlorine) or anhydrous inorganic liquids in the complete absence of moisture or water

Wet Corrosion

The degradation and oxidation of a metal through an electrochemical electron-transfer process occurring in the presence of water containing dissolved salts and gases and oxygen

Sacrificial Metals

Highly reactive metals attached to structural or foundational metals to protect those metals

Rust

A reddish-brown chemical compound called iron oxide, formed when iron reacts with oxygen and water

Half-Equation

Reduction or oxidation equation of a complete redox equation

Oxidise

Cause a chemical species to lose an electron or electrons, increasing its oxidation number

Reduce

Cause a chemical species to gain an electron or electrons, decreasing its oxidation number

Polyatomic Ion

An ion made up of a group of atoms covalently bonded together

Oxidiation State of a Free Element

Zero

Free Element

An element by itself or bonded to itself

Oxidation Number of a Simple Ion

The charge of the ion

Sum of Oxidiation States in a Neutral Compound

Zero

The Sum of Oxidation States in a POlyatomic Ion

Charge of the Ion

Oxidation State of Hydrogen in a Compound

+1 with exceptions explained in unit 3/4

Oxidation State of Oxygen in a Compound

-2 with exceptions explained in unit 3/4

Spontaneous Redox Reaction

Feasible redox reaction that occurs without the addition of any extra heat or energy

Electrode

Electrically conductive medium

Sacrificial Anode

A readily oxidised anode used to coat other metals for protection against corrosion

Sacrificial Protection

The use of a protectie layer that is readily oxidised

Electrolysis

A type of redox reaction which occurs by an electrical current as it is a non-spontaneous reaction

Complex Species

An ion with a metal in the centre with a number of other molecules or ions surrounding it

Overall Redox Reaction

Combination of the oxidation and reduction half-equations

Examples of Redox Reactions

• Rust

• Corrosion

• Electrolysis

• Combustion

• Respiration

• Photosynthesis

• Bleaching

Spontaneously

Occurring naturally on its own under specific conditions without needing a continuous input of outside energy

Displacement Reaction

A chemical process where a more reactive element replaces a less reactive element from a compound

Types of Corrosion

• Wet corrosion

• Dry corrosion

Chemical Energy

Energy stored in the bonds of the compound

Electrical Energy

A type of kinetic energy caused by moving electric charges

Electrolyte

A substance that breaks apart into free ions when melted or dissolved in water