Electrode Potentials

• Define reduction


  Gain of electrons OR decrease in oxidation number
• Define oxidation


  Loss of electrons OR increase in oxidation number
• What does an oxidising agent do?


  Accepts electrons and contains species that are reduced
• What does a reducing agent do?


  Donates electrons to species being reduces. Therefore reducing agents are oxidised
• What do electrochemical cells do?


  Converts chemical energy to electrical energy through redox reactions.
• What is a half cell?


  A half-cell contains the chemical species present in a redox half-equation. An example would be Zn2+(aq)/Zn(s) and Cu2+(aq)/Cu(s).

  
• What are the two key components that link half cells?


  -> An external circuit: A wire that allows the flow of electrons from one electrode to another, creating an electric current. The current is measured with a voltmeter

  -> A salt bridge: This bridge allows ions to transfer between the half-cells, maintaining charge balance. It is often a filter paper, soaked in aqueous potassium nitrate
• Why are the chemicals in the 2 half cells kept apart?

  If the chemicals in the 2 half-cells were allowed to mix, then electrons will flow in an uncontrolled way and heat energy would be released rather than electrical energy
• How is the equilibrium of a half-cell written?


  The forward reaction shows reduction and the reverse shows oxidation
• What does the direction of flow of electrons depend on when the two half-cells are connected?


  The relative tendency of each electrode to release electrons
• What are the two types of half-cells?


  Metal/metal ion half-cells

  Ion/ion half-cells
• What do metal/metal ion half-cells consist of?


  A metal electrode dipped into a solution of its aqueous metal ion. For example, a zinc electrode immersed in a solution of Zn2+ ions.

  -> The metal can either be oxidised to its ions or the ions can be reduced to the metal, depending on the metal's reactivity

  
• What do Ion/ion half-cells consist of?


  Consists of two ions of the same element in different oxidation states, in contact with an inert platinum electrode. For example, Fe2+ and Fe3+ ions in a solution of a platinum electrode

  
• Why is an inert electrode used?


  To provide an inert surface for the electrode transfer between ions of different oxidation states

  
• Where do the two half-reactions occur?


  Oxidation happens at the anode (positive electrode)

  Reduction happens at the cathode (negative electrode)
• Metals that oxidise more readily...?


  Have more negative electrode potentials
• Metals that oxidise less readily...?


  Have less negative or even positive electrode potentials
• Explain what happens at the zinc-copper electrochemical cell


  1) At the zinc anode, zinc atoms are oxidised to Zn2+ ions releasing electrons into the external circuit.

  2) At the copper cathode, Cu2+ ions are reduced to copper atoms gaining electrons from the external circuit

  
• Define standard electrode potential (E°)

  The tendency to be reduced and gain electrons.
• What is the standard hydrogen electrode and what does it feature?


  A universal reference to measuring electrode potentials.

  -> It contains a half-cell containing hydrogen gas

  -> A solution containing H+(aq) ions

  -> An inert platinum electrode

  Standard conditions are:-

  -> Solutions have a concentration of exactly 1 moldm-3

  -> Temperature of 298K (25°)

  -> Pressure of 100kPa

  
• What E° value are metals more likely to have?


  Negative E° values and lose electrons (become oxidised).

  -> The more negative the E° value, the greater the reactivity of a metal in losing electrons
• What are E° value of non-metals more likely to have?


  Positive E° values and gain electrons (become reduced)

  -> The more positive the E° value, the greater the reactivity of a non-metal in gaining electrons
• How do you calculate the standard cell potential from standard electrode potentials?


  E(cell) = E(reduced) - E(oxidised)

  
• In an electrochemical series, where is the strongest reducing agent?


  The top right
• In an electrochemical series, where is the strongest oxidising agent?


  The bottom left
• How do you predict reaction feasibility using electrode potentials?


  If the E° cell is positive, the reaction is feasible under standard condition. The more positive the value of the E° cell, the more feasible the redox reaction is
• What are the limitations of predictions using E° values?


  -> Reaction rates: A reaction predicted to be feasible based on the E° ΔG lies with reactions may proceed very slowly if it has a very large activation energy

  -> Concentration: If the concentration of a solution is not 1moldm-3, then the value of the electrode potential will be different from the standard value.

  -> Non-aqueous reactions: Standard electrode potentials are determined for aqueous equilibria. They may not accurately predict feasibility reactions in non-aqueous solvents
• What are fuel cells?


  A fuel cell uses the energy from the reaction of a fuel with oxygen to create a voltage. Fuel cells can operate continuously and generate electricity as long as fuel and oxygen are provided.

  -> Fuel cells do not need to be recharged
• Give an example of a fuel cell


  Hydrogen fuel cells:- Seen as a promising source for electric vehicles.

  -> In a hydrogen fuel cell, hydrogen (the fuel) and oxygen react to produce electricity, with water as the only by-product
• Describe what happens in a hydrogen fuel cell


  1) At the anode, hydrogen is split into protons (H+) and electrons (e-) with the help of a platinum catalyst

  2) The protons (H+) move through the polymer electrolyte membrane, which only allows protons to pass. This forces the electrons to travel through the external circuit to get to the cathode

  3) The electrons flowing through the external circuit generates an electrical current that can power devices

  4) Oxygen at the cathode combines with the protons from the anode and the electrons from the circuit to produce water (as a waste product)

  
• What are primary cells?


  Primary cells are non-rechargeable and are designed to be used once only. Electrical energy is produced by oxidation and reduction at the electrodes and the reactions cannot be reversed.

  Examples of primary cells:-

  -> Wall clocks

  -> Smoke detectors

  Most modern primary cells are alkaline based on zinc and manganese dioxide
• What are secondary cells?


  Secondary cells are rechargeable as the reaction producing electrical energy can be reversed (chemicals in the cell are regenerated and can be used again)

  Examples of secondary cells:-

  -> Lead-acid batteries used in car batteries

  -> Lithium-ion cells used in our modern appliances (e.g. laptops, tablets, mobile phones)