Electrode potentials and redox

• An oxidising agent itself gets reduced – gains electrons.
• Some substances can act both as oxidising and reducing agents.
• In a titration, the concentration of a solution is determined by titrating with a solution of known concentration.
• The nature of a substance as an oxidising or reducing agent is dependent upon what it is reacting with and the reaction conditions.
• In redox titrations, an oxidizing agent is titrated against a reducing agent.
• Electrons are transferred from one species to the other.
• Indicators are sometimes used to show the endpoint of the titration.
• Most transition metal ions naturally change colour when changing oxidation state.
• There are two common redox titrations you should know about manganate(VII) titrations and iodine-thiosulfate titrations.
• In manganate(VII) titrations, the manganate(VII) is the oxidising agent and is reduced to Mn2+(aq).
• The iron in manganate(VII) titrations is the reducing agent and is oxidised to Fe3+(aq).
• In manganate(VII) titrations, the reaction mixture must be acidified, to excess acid is added to the iron(II) ions before the reaction begins.
• The acid used in manganate(VII) titrations must not react with the manganate(VII) ions, so the acid normally used is dilute sulfuric acid.
• Dilute sulfuric acid does not oxidise under these conditions and does not react with the manganate(VII) ions.
• Standard electrode potential (E°) is measured at pH = 0, pressure = 1 atm, concentration = 1 M, temperature = 298 K.
• The standard reduction potential is the tendency of an element to gain electrons.
• A positive value indicates that the species will accept electrons and act as an oxidising agent.
• Reduction half-reactions are written on the left side of the equation, while oxidation half-reactions are written on the right side.
• Potassium permanganate acts as its own indicator, as the purple potassium permanganate solution is added to the titration flask from the burette and reacts rapidly with the Fe2+(aq).
• The manganese(II) ions, Mn2+(aq), have a very pale pink colour but they are present in such a low concentration that the solution looks colourless.
• The burette used in this practical should be one with white numbering not black, as you would struggle to read the values for your titres against the purple colour of the potassium permanganate if black numbering was used.
• As soon as all of the iron(II), Fe2+(aq), ions have reacted with the added manganate(VII) ions, Mn7+(aq), a pale pink tinge appears in the flask due to an excess of manganate(VII) ions, Mn7+(aq).
• The position of equilibrium and therefore the electrode potential depends on factors such as temperature, pressure of gases, and concentration of reagents.
• To be able to compare the electrode potentials of different species, they all have to be measured against a common reference or standard.
• Standard conditions for electrode potentials include an ion concentration of 1.00 mol dm-3, a temperature of 298 K, and a pressure of 100 kPa.
• Standard measurements for electrode potentials are made using a high resistance voltmeter so that no current flows and the maximum potential difference is achieved.
• The electrode potentials are measured relative to a standard hydrogen electrode, which is given a value of 0.00 V, and all other electrode potentials are compared to this standard.
• The standard electrode potential (Eθ) is the potential difference (sometimes called voltage) produced when a standard half-cell is connected to a standard hydrogen cell under standard conditions.
• Br2(l) + 2e– ⇌ 2Br–(aq) Eθ = +1.09 V
• 2H+(aq) + 2e– ⇌ H2(g) Eθ = 0.00 V
• Na+ (aq) + e– ⇌ Na(s) Eθ = -2.71 V
• 2H+ (aq) + 2e– ⇌ H2(g) Eθ = 0.00 V
• Chemists use a type of shorthand convention to represent electrochemical cells.
• In this convention, a solid vertical (or slanted) line shows a phase boundary, which is an interface between a solid and a solution.
• A double vertical line (sometimes shown as dashed vertical lines) represents a salt bridge.
• A salt bridge has mobile ions that complete the circuit.
• Potassium chloride and potassium nitrate are commonly used to make the salt bridge as chlorides and nitrates are usually soluble.
• This should ensure that no precipitates form which can affect the equilibrium position of the half cells.
• The substance with the highest oxidation state in each half cell is drawn next to the salt bridge.
• The cell potential difference is shown with the polarity of the right hand electrode.