Redox & Electrochemistry

Oxidation numbers (also called oxidation states) are a way of keeping track of electron distribution in compounds. They represent the hypothetical charge an atom would have if all bonds were ionic. Rules for assigning oxidation numbers: 1. The oxidation number of an element in its uncombined state is 0 (e.g., O₂ = 0, Na = 0). 2. The oxidation number of a monatomic ion is equal to its charge (e.g., Na⁺ = +1, Cl⁻ = -1). 3. In compounds, Group 1 metals are +1, Group 2 metals are +2. 4. Fluorine is always -1 in compounds. 5. Oxygen is usually -2 in compounds, except in peroxides (e.g., H₂O₂) where it is -1, and with fluorine (e.g., OF₂) where it is +2. 6. Hydrogen is usually +1 in compounds, except in metal hydrides (e.g., NaH) where it is -1. 7. The sum of the oxidation numbers in a neutral compound is 0. 8. The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion.

Oxidation is the loss of electrons by a species, resulting in an increase in its oxidation number. The species that undergoes oxidation is called the reducing agent because it causes another species to be reduced.

Reduction is the gain of electrons by a species, resulting in a decrease in its oxidation number. The species that undergoes reduction is called the oxidising agent because it causes another species to be oxidised.

A redox reaction is a chemical reaction in which both oxidation and reduction occur simultaneously. Electrons are transferred from one species to another.

Redox reactions can be split into two half-equations: one representing the oxidation process and the other representing the reduction process. These half-equations show the electron transfer explicitly. To balance half-equations in acidic solution: 1. Balance atoms other than oxygen and hydrogen. 2. Balance oxygen atoms by adding H₂O molecules. 3. Balance hydrogen atoms by adding H⁺ ions. 4. Balance the charge by adding electrons (e⁻) to the more positive side.

Electrolysis is the process of using electrical energy to drive a non-spontaneous chemical reaction. It involves passing an electric current through an electrolyte (a molten ionic compound or an ionic solution) to cause chemical decomposition. - **Anode**: The positive electrode where oxidation occurs (anions lose electrons). - **Cathode**: The negative electrode where reduction occurs (cations gain electrons). - **Electrolyte**: The substance containing free ions that conducts electricity. Applications include electroplating, extraction of reactive metals (e.g., aluminium from bauxite), and purification of metals (e.g., copper).

Electrochemical cells are devices that convert chemical energy into electrical energy through spontaneous redox reactions. They consist of two half-cells, each containing an electrode immersed in an electrolyte. - **Anode**: The electrode where oxidation occurs (negative terminal). - **Cathode**: The electrode where reduction occurs (positive terminal). - **Salt Bridge**: Connects the two half-cells, allowing ion flow to maintain electrical neutrality. - **External Circuit**: Wires connecting the electrodes, allowing electron flow. Standard electrode potential (E°) is the potential difference of a half-cell compared to the standard hydrogen electrode (SHE, defined as 0.00 V) under standard conditions (298 K, 1 atm, 1 M concentration). The cell potential (E°_cell) is calculated as E°_cell = E°_cathode - E°_anode. The more positive E° value corresponds to the species that will be reduced (cathode).

Corrosion is the deterioration of a metal due to its reaction with its environment, typically an electrochemical process. Rusting of iron is a common example of corrosion. - **Mechanism of Rusting**: Iron acts as the anode, oxidising to Fe²⁺ (Fe → Fe²⁺ + 2e⁻). Oxygen and water act as the cathode, reducing to hydroxide ions (O₂ + 2H₂O + 4e⁻ → 4OH⁻). The Fe²⁺ and OH⁻ ions then react to form iron(II) hydroxide, which is further oxidised by oxygen to hydrated iron(III) oxide (Fe₂O₃.nH₂O), which is rust. - **Galvanic Corrosion**: Occurs when two different metals are in electrical contact in the presence of an electrolyte. The more reactive metal (with a more negative standard electrode potential) acts as the anode and corrodes preferentially, protecting the less reactive metal (cathode). - **Prevention**: Methods include barrier protection (painting, oiling, plastic coating), galvanising (coating with zinc, which acts as a sacrificial anode), and cathodic protection (connecting the metal to a more reactive metal or an external power supply).