Lecture 9. Chemistry of Oxidation-Reduction Processes Prepared by PhD Halina Falfushynska

Oxidation-Reduction Reactions Often called “redox” reactions Electrons are transferred between the reactants – One substance is oxidized, loses electrons Reducing agent – Another substance is reduced, gains electrons Oxidizing agent Oxidation numbers change during the reaction

LEO says GER LEO says GER Lose Electrons = Oxidation Sodium is oxidized Gain Electrons = Reduction Chlorine is reduced

Rules for assigning oxidation numbers 1.Elements (uncombined) are 0. Al, N 2, He, Zn, Ag, Br 2, O 2, O 3 2. Oxidation numbers must sum to the overall charge of the species. SO 4 2  =  2 (O is usually  2 so….)‏ ? + 4(  2) =  2 Solve: ?  8 =  2 ? = + 6 (S)

Guidelines for Assigning Oxidation Numbers  is  1 and for KO 2 is  ½.

Assign oxidation numbers for all elements in each species MgBr 2 Mg +2, Br  1 ClO 2  Cl +1, O  2

Copyright McGraw-Hill Oxidation Numbers on the Periodic Table (most common in red)

Displacement reactions – A common reaction: active metal replaces (displaces) a metal ion from a solution Mg(s) + CuCl 2 (aq)  Cu(s) + MgCl 2 (aq) – The activity series of metals is useful in order to predict the outcome of the reaction.

Balancing redox reactions – Electrons (charge) must be balanced as well as number and types of atoms – Consider this net ionic reaction: Al(s) + Ni 2+ (aq)  Al 3+ (aq) + Ni(s) – The reaction appears balanced as far as number and type of atoms are concerned, but look closely at the charge on each side.

Al(s) + Ni 2+ (aq)  Al 3+ (aq) + Ni(s) – Divide reaction into two half-reactions Al(s)  Al 3+ (aq) + 3e  Ni 2+ (aq) + 2e   Ni(s) – Multiply by a common factor to equalize electrons (the number of electrons lost must equal number of electrons gained)‏ 2 [Al(s)  Al 3+ (aq) + 3e  ] 3 [Ni 2+ (aq) + 2e   Ni(s) ]

– Cancel electrons and write balanced net ionic reaction 2Al(s)  2Al 3+ (aq) + 6e  3Ni 2+ (aq) + 6e   3Ni(s) 2Al(s) + 3Ni 2+ (aq)  2Al 3+ (aq) + 3Ni(s) ‏

Predict whether each of the following will occur. For the reactions that do occur, write a balanced net ionic reaction for each. - Copper metal is placed into a solution of silver nitrate -A gold ring is accidentally dropped into a solution of hydrochloric acid No reaction occurs, gold is below hydrogen on the activity series.

Combination Reactions – Many combination reactions may also be classified as redox reactions – Consider: Hydrogen gas reacts with oxygen gas 2H 2 (g) + O 2 (g)  2H 2 O(l) Identify the substance oxidized and the substance reduced.

15 Decomposition reactions – Many decomposition reactions may also be classified as redox reactions – Consider: Potassium chlorate is strongly heated 2KClO 3 (s)  2KCl(s) + 3O 2 (g) Identify substances oxidized and reduced.

16 Disproportionation reactions – One element undergoes both oxidation and reduction – Consider:

17 Combustion reactions – Common example, hydrocarbon fuel reacts with oxygen to produce carbon dioxide and water – Consider:

Reaction of Cu and Zn 2+ ions

Gets Smaller -> <- Gets Larger

Cell Notation Zn (s) + Cu 2+ (aq) Cu (s) + Zn 2+ (aq) [Cu 2+ ] = 1 M & [Zn 2+ ] = 1 M Zn (s) | Zn 2+ (1 M) || Cu 2+ (1 M) | Cu (s) anode cathode Zn (s)| Zn +2 (aq, 1M)|K(NO 3 ) (satur)|Cu +2 (aq, 1M)|Cu(s) anodecathode Salt bridge

K(NO 3 ) Zn (s) + 2 H + (aq) -> H 2 (g) + Zn +2 (aq) Zn(s)| Zn +2 |KNO 3 |H + (aq)|H 2 (g)|Pt

Electrochemical Cells The difference in electrical potential between the anode and cathode is called: cell voltage electromotive force (emf) cell potential

Standard Electrode Potentials Standard reduction potential (E 0 ) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm.    V Standard hydrogen electrode (SHE)  e       atm  Reduction Reaction

Determining if Redox Reaction is Spontaneous + E° CELL ; spontaneous reaction E° CELL = 0; equilibrium - E° CELL ; nonspontaneous reaction More positive E° CELL ; stronger oxidizing agent or more likely to be reduced

Relating E 0 Cell to  G 0 Units work, Joule charge, Coulomb E cell ; Volts Faraday, F; charge on 1 mole e- F = C/mole  G = -nFEcell

Relating   CELL to the Equilibrium Constant, K  G 0 = -RT ln K  G 0 = -nFE 0 cell -RT ln K = -nFE 0 cell

Corrosion – Deterioration of Metals by Electrochemical Process

Corrosion Damage done to metal is costly to prevent and repair Iron, a common construction metal often used in forming steel alloys, corrodes by being oxidized to ions of iron by oxygen. This corrosion is even faster in the presence of salts and acids, because these materials make electrically conductive solutions that make electron transfer easy

Corrosion Luckily, not all metals corrode easily Gold and platinum are called noble metals because they are resistant to losing their electrons by corrosion Other metals may lose their electrons easily, but are protected from corrosion by the oxide coating on their surface, such as aluminum – Figure 20.7, page 636 Iron has an oxide coating, but it is not tightly packed, so water and air can penetrate it easily

Corrosion Serious problems can result if bridges, storage tanks, or hulls of ships corrode Can be prevented by a coating of oil, paint, plastic, or another metal If this surface is scratched or worn away, the protection is lost Other methods of prevention involve the “sacrifice” of one metal to save the second Magnesium, chromium, or even zinc (called galvanized) coatings can be applied