QCEChemistry
Redox reactions
recognise that a range of reactions, including displacement reactions of metals, combustion, corrosion and electrochemical processes, can be modelled as redox reactions involving oxidation of one substance and reduction of another substance
- Redox is a chemical reaction involving the transfer of electrons from one reactant to another
○ Oxidation is the loss of electrons
○ Reduction is the gain of electrons (OIL RIG)
Displacement reactions of metals
- Single displacement is a chemical reaction in which a more reactive metal ion replaces a less reactive metal ion in a molecule
- CuSO4 (aq) + Zn (s) ZnSO4 (aq) + Cu (s)
Copper is a cation in the solution - gains electrons to form copper metal copper undergoes reduction
Combustion reactions
- Combustion is a chemical reaction with oxygen to form a metal oxide, a covalent compound, or carbon dioxide and water
Metal + Oxygen Metal Oxide
- 2Cu (s) + O2 (g) 2CuO (s)
○ Copper metal loses electrons to form a positive metal cation copper undergoes oxidation
○ Oxygen gas gains electrons to form a negative non-metal anion oxygen undergoes reduction
Corrosion
- Corrosion is the degradation of a metal to form a more stable metal oxide when exposed to gases and liquid
- Wet corrosion can occur in moist air or by direct immersion in water, and accelerates the rate of corrosion
▪ Fe (s) Fe+2 + 2e-
▪ O2 (aq) + 2H2O (l) + 4e- 4OH- (aq)
▪ Overall equation: 2Fe (s) +O2 (aq) + 2H2O (l) 2Fe+2 (aq) + 4OH- (aq)
▪ Fe+2 (aq) + 2OH- (aq) Fe(OH)2 (s)
Step 3: further oxidation of iron (II) hydroxide occurs in the presence of oxygen and water to produce iron (III) hydroxide, a red-brown precipitate
▪ 4Fe(OH)2 (s) + O2 (aq) + 2H2O (l) 4Fe(OH)3 (s)
Step 4: In air, the iron (III) hydroxide loses water to form hydrated iron (III) oxide (Fe2O3 xH2O)
understand that the ability of an atom to gain or lose electrons can be predicted from the atom's position in the periodic table, and explained with reference to valence electrons, consideration of energy and the overall stability of the atom
- Group 1 and 2
○ Lose valence electrons to form positive cations easily
○ By losing electrons, they are oxidised and act as reducing agents
○ Strong reducing agents
- Group 17
○ Gain valence electrons to form negative anions easily
○ By gaining electrons, they are reduced and act as oxidising agents
○ Strong oxidising agents
- Ionisation energy
○ Elements in bottom left of the table have the lowest ionisation energy, meaning they lose their valence electrons readily
○ These elements are strong reducing agents
- Electronegativity
○ Elements in top right-hand side of periodic table have strongest electronegativities, meaning they accept electrons readily
○ These elements are strong oxidising agents
identify the species oxidised and reduced, and the oxidising agent and reducing agent, in redox reactions
- Oxidise means to lose electrons. Reduce means to gain electrons
Typically metals are oxidised and non-metals are reduced
- Oxidising agent is a reactant that causes another reactant to be oxidised and is itself reduced.
- Reducing agent is a reactant that causes another reactant to gain electrons and be reduced and is itself oxidised.
understand that oxidation can be modelled as the loss of electrons from a chemical species, and reduction can be modelled as the gain of electrons by a chemical species; these processes can be represented using balanced half-equations and redox equations (acidic conditions only)
Important: go through notebook for examples of balancing equation examples
deduce the oxidation state of an atom in an ion or compound and name transitional metal compounds from a given formula by applying oxidation numbers represented as roman numerals
use appropriate representations, including half-equations and oxidation numbers, to communicate conceptual understanding, solve problems and make predictions
understand that electrochemical cells, including galvanic and electrolytic cells, consist of oxidation and reduction half-reactions connected via an external circuit that allows electrons to move from the anode (oxidation reaction) to the cathode (reduction reaction).
- Electrochemical cells is a device in which chemical energy is converted into electrical energy, or vice versa
understand that galvanic cells, including fuel cells, generate an electrical potential difference from a spontaneous redox reaction which can be represented as cell diagrams including anode and cathode half-equations
- Galvanic cells are an electrochemical cell in which the reduction and oxidation half-equations are separated and connected through a circuit to generate electricity
recognise that oxidation occurs at the negative electrode (anode) and reduction occurs at the positive electrode (cathode) and explain how two half- cells can be connected by a salt bridge to create a voltaic cell (examples of half-cells are Mg, Zn, Fe and Cu and their solutions of ions)
- Oxidation occurs at the negative electrode (anode) and reduction occurs at the positive electrode (cathode)
- The salt bridge is an electrical connection between the two half-cells of an electrochemical cell and allows for the flow of charge by moving ions - cations to the cathode and anions to the anode
▪ Causes build-up of positive charge in the oxidation half-cell. Negative ions in the salt bridge migrate towards the oxidation half-cell to balance the charge
▪ Build-up of negative spectator ions that are left over in the solution. Positive ions in the salt bridge migrate towards reduction half-cell to balance the charge
describe, using a diagram, the essential components of a galvanic cell; including the oxidation and reduction half-cells, the positive and negative electrodes and their solutions of their ions, the flow of electrons and the movement of ions, and the salt bridge.
Galvanic cell
Fuel Cell
- Fuel cells are a galvanic cell that produces electricity by using a constant supply of reactants (often hydrogen and oxygen) and inert electrodes that do not break down
