Electrochemistry

• we can use redox reactions to generate electrons that we can then make flow
• we can use half equations as electron producers and acceptors
• all half equations have the ability to release electrons. this ability varies depending on the half equation
• when we place a metal in a solution of its own salt, the metal atoms lose their electrons. the electrons stay on the metal strip, making it negatively charged compared the solution. the metal ions in the solution then pick up the electrons to reform the metal atoms. this creates potential difference between the metal and the solution
• the more reactive the metal, the greater its ability to release electrons and the greater potential difference it creates between the metal strip and the solution
• an electrochemical cell consists of two half cells connected by a wire. a half cell is made up of a solid metal electrode dipped in an electrolyte of its metal ions.
• the salt bridge connects the two solutions in the electrochemical cell and is made of a saturated solution of potassium nitrate (K+ and NO3-). it completes the circuit and neutralises the charge in the solutions by allowing ion from itself into the solutions
• the ions in the salt bridge must be inert so they can't produce any by-products with any of the ions in the system
• the more reactive electrode is the negative electrode as it is more able to release electrons. the negative electrode loses electrons which flow through the wires, lighting up the bulb, and flow into the beaker of the other metal solution. the less reactive metal ions in the solution pick up the electrons to form metal atoms
• the negative electrode undergoes an oxidation reaction as it loses electrons and the positive electrode undergoes a reduction reaction as it gains electrons
• any half equation can be a half cell. if it doesn't have a solid metal in the half equation e.g. 2H+ + 2e- --> H2 , insert platinum to act as an artificial electrode. platinum is a good artificial electrode as it is a good conductor and is inert
• the electrochemical series is a list of reduction half equations in order of their ability to release electrons
• the availability of electrons from a half equation is indicated by their standard eledtrode potential. each half cell has its own electrode potential value in the electrochemical series
• the standard electrode potential of a half equation is measured against the standard hydrogen electrode (SHE). SHE is used to measure the standard electrode potential of all the other half cells under standard conditions
• the standard hydrogen electrode half cell involves the equation

  H+(aq) + e- --> 1/2 H2(g) . one mol dm-3 HCl is used as the source of H+ ions and a supply of hydrogen gas at 100kPa
• an inert platinum electrode is used to allow the hydrogen half equation to taqke place as it does not already have its own solid electrode
• the standard hydrogen electrode is used to find the standard electrode potential of another half equation by connecting the electrodes with a high resistance voltmeter and a salt bridge
• the standard electrode potential of the hydrogen electrode is always set to 0V BY DEFINITION which means that the reading on the voltmeter will be due to the other half cell
• the standard hydrogen electrode is set up at 298K and in absence of air as it prevents oxygen from oxidising any of the electrodes
• the more positive the standard electrode potential (E standard), the less electron-releasing it is, and less reactive
• we use a high resistance voltmeter in SHE as there is zero current flow
• for a pair of half equations, the one with the more positive E standard value will be reduction and the one with the more negative E standard value will be oxidation
• we can use the electrochemical series to write cell representations
• rules for cell representations

  • oxidation half cell always on the left of the salt bridge and the reduction is on the right, except if you are using SHE which will always be on the left
• rules for cell representation

  • write the species in order as they appear in the half equations and include state symbols
• rules for cell representation

  • if the half cell doesn't have its own electrode, add Pt at the start or end as required
• rules for cell representation

  • if there is no change of state in the half equation, use commas instead of phase boundaries
• how to write cell representations

  • separate species in the same half equation with phase boundaries, which are represented by |
  • separate the two half equations with a salt bridge, represented by ||
• we can use the electrochemical series to predict redox reactions
• we can calculate the E cell value for a reaction using the E standard values for the half equations on the electrochemical series.
• the E cell value tells us if a reaction is feasible or not. if E cell is positive, the reaction is feasible
• E cell = E reduction - E oxidation
• predicting redox reactions

  • write out the half equations from the electrochemical series with their respective E standard values
  • more positive is reduction, more negative is oxidation
  • combine the half equations to make the cell reaction (redox reaction)
  • calculate E cell which is the forward reaction
  • UNDERLINE the species in the question and species in the cell reaction to answer question
• does Cu2+ get reduced by Zn metal

  • cell reaction:
  • Cu2+ + Zn --> Cu + Zn2+ forward reaction is +1.1V
  • Yes, E cell is positive (+1.1V) when Zn reacts with Cu2+
  • E standard of Cu2+/Cu is greater than E standard than Zn2+/Zn
  • copper ions can be reduced by zinc, which is oxidised
• we can use the electrochemical series to select the strongest and weakest oxidising agents and reducing agents
• how to find the strongest oxidising agent

  • oxidising agents are found on the left hand side of the series
  • found in the half equation that is always reduced (most positive E standard)
• how to find the strongest reducing agent

  • reducing agents are found on the right hand side of the series
  • found in the half equation that is always oxidised (most negative E standard)
• the strongest oxidising agents have the most potential to be reduced
• the strongest reducing agents have the most potential to be oxidised
• explain why HCl should not be used to acidify the solution containing iron (II) ions

  • MnO4-/Mn2+ has a greater E standard than Fe2+/Fe3+ therefore Fe2+ is oxidised and MnO4- is reduced
  • Cl- from HCl is a reducing agent so can reduce MnO4-
  • titre volume increases are there is now more moles of reducing agent