Chem U3

• In redox reactions, something is oxidised and something else is reduced
• Oxidation is the loss of electrons
• Reduction is the gain of electrons
• An oxidising agent accepts electrons and helps oxidation, becoming reduced itself
• A reducing agent donates electrons and helps reduction, becoming oxidised itself
• Rules for assigning oxidation numbers in a compound:

  • All elements have an oxidation number of zero
  • Hydrogen is +1 unless with a Group 1 metal, then it's -1
  • Oxygen is -2 (except in peroxides when it's -1 or with fluorine when it's +2)
  • Group 1 and 2 elements are +1 and +2 respectively
  • Oxidation numbers in a compound must add up to zero or the ion's charge
  • The most electronegative element is given the negative oxidation number
• Changes in oxidation number indicate if an element has been oxidised or reduced:

  • If the oxidation number increases, the element is oxidised
  • If the oxidation number decreases, the element is reduced
• Equations for redox reactions can be split into two half-equations, one for oxidation and one for reduction
• Electrochemical cells have half-cells for oxidation and reduction, joined to form a complete circuit
• Three types of half-cells:

  1. Metal/Metal ion

  2. Gas/Non-metal ion solution

  3. Solution of a metal in two different oxidation states
• Cells can be represented using cell diagrams:

  • The conducting metal of the left-hand cell goes first
  • Change of state is shown with a single vertical line
  • Same physical state uses commas
  • Salt bridge is represented by a double vertical line
  • The more positive electrode is on the right-hand side
• Standard electrode potentials are measured against the standard hydrogen electrode, which is taken as 0.0 V
• The standard hydrogen electrode setup:

  • Platinum electrode coated with fine platinum grains
  • Standard conditions: 1 atm pressure for H2 gas, H+ concentration of 1 mol dm^-3, temperature of 298K
• Using standard electrode potentials helps determine charge on each electrode, direction of electron flow, and likelihood of a reaction
• The Electrochemical Series orders reducing power of different half-cells:

  • Most reactive metals have the most negative Eθ
  • Least reactive metals/most reactive non-metals have the most positive Eθ
• The zinc electrode is negative and the copper one is positive because electrons are being produced in this reaction
• Zinc metal is being oxidised to zinc ions
• Copper ions are being reduced to copper atoms
• Zinc metal is acting as a reducing agent because it’s supplying electrons
• Copper ions are acting as an oxidising agent because they are accepting electrons
• The red arrow on the diagram refers to the ions on the left, and the blue arrow refers to the reduced species on the right
• EMF values for cells must always be positive
• The EMF of the cell can be calculated using the standard electrode potentials of the two half-cells
• For the given cell: Eθ = +0.34 – (-0.76) = +1.10V
• Eθ values are positive or negative depending on whether the species has a more positive or negative potential compared to hydrogen
• If we compare zinc ions with hydrogen ions, then zinc ions lose electrons more easily than hydrogen ions
• The reaction feasibility can be predicted using standard electrode potential values
• For a reaction to go ahead, its EMF must be positive
• The oxidation half-reaction is the one with the most negative Eθ value
• The reduction half-reaction is the one with the most positive Eθ value
• When calculating the EMF of a reaction, determine which half-reaction is the most positive (reduction) and which is the most negative (oxidation)
• In redox reactions, if there is a change in oxidation numbers, then it is a redox reaction
• Redox reactions are common in chemistry and biology, seen in processes like respiration and photosynthesis
• Organic chemistry in Unit 4 includes examples of redox reactions such as oxidation of alcohols and reduction of nitriles
• Fuel cells convert chemical energy stored in fuel into electrical energy efficiently
• Advantages of fuel cells include producing only water and being more energy efficient
• Disadvantages of fuel cells include the flammability of hydrogen gas and energy loss in hydrogen production from fossil fuels
• You need to construct and combine ion/electron half-equations for redox reactions
• Example 1: Chlorine gas oxidises iron(II) ions to iron(III) ions, and chlorine is reduced to chloride ions
• Example 2: Manganate (VII) ions can oxidise hydrogen peroxide to oxygen gas, and manganate (VII) ions are reduced to manganese (II) ions