1. Lecture 14 CM1001

2. Oxidation Numbers Positive or negative numbers decided using agreed rules. Help us work out whether a substance is oxidized or reduced. Rules : 1/ Oxidation number of an atom of a free element is zero e.g. N 2, Ca 2/ Oxidation number of an ion is equal to its charge e.g. Cu 2+, O 2- 3/ The oxidation state of Flourine is always -1 in a compound 4/ Group I elements have an oxidation state of +1 in a compound 5/ Group II elements have an oxidation state of +2 in a compound 6/ Group VII elements have an ox. state of -1, e.g. HCl

3. Oxidation State of Oxygen 7/ Oxygen usually has an ox. state of -2 but there are 3 exceptions (a) In compounds with Flourine, Oxygen has a positive ox. State e.g. Oxygen is +2 in OF 2 (b) In peroxides Oxygen has a ox. State of -1 e.g. H 2 O 2 (c) In compounds which contain superoxides (O 2 - ), Oxygen has a ox.state of -1/2 e.g. KO 2. 8/ Hydrogen has an ox.state of -1 in compounds with less electronegative elements and +1 in compounds with more electronegative elements e.g. NaH and CH 4 9/ The sum of the oxidation number of the atoms in a neutral compound is zero, e.g. Ca(OH)2 [Ca(2) + 2 O(-4) + 2H(2) =0]

4. ElementOxidation number K +, Na + +1 Mg 2+, Ca 2+ +2 Al 3+ +3 H + +1 F -, Cl - -1 O 2- -2 These elements almost always have these oxidation numbers. During a reaction, an increase in oxidation number indicates oxidation, a decrease in oxidation number indicates reduction Oxidation Numbers

5. Redox Reaction An oxidizing agent is a substance that takes up electrons during a reaction and so is reduced. A reducing reagent supplies the electrons and so is oxidized. This is a redox reaction. To write and balance a redox equation for a chemical reaction: 1.Identify the atoms that are oxidized and reduced, using the oxidation number 2.Balance the half reactions: 3.Make sure that there are the same number of atoms of the element that is oxidised or reduced on each side of the half equation 4.If there are any oxygen atoms present, balance them with water molecules on the other side of the half reaction 5.If there are any hydrogen atoms present, balance them with hydrogen ions 6.Make sure that the half-reactions have the same overall charge by adding electrons

6. Redox Reaction If a piece of zinc metal, Zn (s), is dipped in a solution of silver ions, Ag + (aq), a reaction occurs and silver metal is produced. Zn (s) + 2 Ag + (aq) → Zn 2+ + 2Ag (s) The half reactions are: Zn (s) → Zn 2+ + 2e - 2Ag + (aq) +2e - → 2 Ag (s) The fact that a reaction occurs means that the silver ions want to accept electrons more than the zinc ions want to keep them. Zinc metal dipped in magnesium ions – no reaction, the magnesium ions have a lower tendency to accept electrons than silver ions.

7. Redox reactions have many applications in the modern world e.g. power supply units in cars, phones ipods and a myriad of consumer products. The transfer of electrons in redox reactions can be harnessed for energy by separating the reactants into separate ‘half’cells. When the ‘half’ cells are paired they work as a battery and are known as galvanic of voltaic cells. Each half cell consists of a ‘redox couple’. Two members of the couple are the same element or compound, but of different oxidation states e.g., Ag + (aq) /Ag (s) and Mg 2+ (aq) /Mg (s) The Ag + (aq) /Ag (s) couple is a stronger oxidising agent than the Mg 2+ (aq) /Mg (s) couple. We can also say that the Mg 2+ (aq) /Mg (s) is a stronger reducing agent than the Ag + (aq) /Ag (s) couple. Electrochemistry

8. Standard Reduction Potential E° The relative oxidizing or reducing strength of redox couples are expressed in terms of their Standard Reduction Potential E (in Volts) The E° of a redox couple is measured with reference to the Standard Hydrogen Electrode (S.H.E.). The S.H.E. is assigned potential of 0V. This consists of a H 2 gas electrode, a piece of platinum dipped in a solution of hydrogen ions (1M) with hydrogen gas (1atm) bubbling over the surface of the platinum. To find the potential to any other electrode relative to S.H.E., construct a galvanic cell.

9. Zn (s) is added to an aqueous solution of HNO 3, Zn dissolves, H 2 bubbles off Zn (s) + 2H 3 O + (aq) +2NO 3 - → H 2(g) + Zn 2+ (aq) + 2NO 3 - (aq) + 2 H 2 O (l) oxidation./reduction reaction- can be expressed in 2 half reactions Zn (s) →Zn 2+ (aq) + 2e - Oxidation 2H 3 O + (aq) +2e - → H 2(g) + 2 H 2 O (l) Reduction Allow these reactions to proceed but in two separate beakers: Standard Reduction Potential E°

11. Salt Bridge: tube containing an electrolyte solution e.g.,(KCl). Ions migrate from the salt bridge to keep solutions neutral. The two solutions cannot mix. 2 half reactions in two separate beakers: e - transfer occurs across the wire as an electric current convert chemical energy →electrical energy. Electrical current measured across the wire is the cell potential. As S.H.E. = 0V E°cell = E° Mg/Mg 2+ - E° S.H.E. = E° Zn/Zn 2+ E° measure of the electron attracting power Metals with large negative reduction potentials are good reducing agent, most easily oxidised themselves Elements with large reduction potentials are good oxidising agents. Reduced form of any element reduces the oxidised form of any element below it Standard Reduction Potential E°