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

eNaNa10

Lose Electrons = Oxidation

Sodium is oxidized

Gain Electrons = Reduction

10 CleCl Chlorine is reduced

• Rules for assigning oxidation numbers 1. Elements (uncombined) are 0.

Al, N2, He, Zn, Ag, Br2, O2, O3

2. Oxidation numbers must sum to the overall charge of the species.

SO42 = 2 (O is usually 2 so….)

? + 4(2) = 2 Solve: ? 8 = 2 ? = + 6 (S)

Guidelines for Assigning Oxidation Numbers

is 1 and for KO2 is ½.

Assign oxidation numbers for all elementsin each species MgBr2

Mg +2, Br 1

ClO2

Cl +1 , O 2

Copyright McGraw-Hill 2009 7

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) + CuCl2(aq) Cu(s) + MgCl2(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) + Ni2+(aq) Al3+(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) + Ni2+(aq) Al3+(aq) + Ni(s)

– Divide reaction into two half-reactions Al(s) Al3+(aq) + 3e

Ni2+(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) Al3+(aq) + 3e ]

3 [Ni2+(aq) + 2e Ni(s) ]

– Cancel electrons and write balanced net ionic reaction

2Al(s) 2Al3+(aq) + 6e

3Ni2+(aq) + 6e 3Ni(s)

2Al(s) + 3Ni2+(aq) 2Al3+(aq) + 3Ni(s)

Predict whether each of the following willoccur. 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.

Cu (s) + 2 Ag (aq) Cu 2+ (aq) + 2 Ag(s)

• Combination Reactions– Many combination reactions may also

be classified as redox reactions – Consider:

Hydrogen gas reacts with oxygen gas

2H2(g) + O2(g) 2H2O(l)

Identify the substance oxidized and the substance reduced.

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• Decomposition reactions– Many decomposition reactions may also

be classified as redox reactions – Consider:

Potassium chlorate is strongly heated 2KClO3(s) 2KCl(s) + 3O2(g)

Identify substances oxidized and reduced.

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• Disproportionation reactions– One element undergoes both oxidation

and reduction – Consider:

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• Combustion reactions– Common example, hydrocarbon fuel

reacts with oxygen to produce carbon dioxide and water

– Consider:

Reaction of Cu and Zn2+ ions

Gets Smaller -> <- Gets Larger

Cell NotationCell Notation

Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)

[Cu2+] = 1 M & [Zn2+] = 1 M

Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)anode cathode

Zn (s)| Zn+2 (aq, 1M)|K(NO3) (satur)|Cu+2(aq, 1M)|Cu(s)anode cathodeSalt bridge

K(NO3)

Zn (s) + 2 H+(aq) -> H2 (g) + Zn+2 (aq)

Zn(s)| Zn+2|KNO3|H+(aq)|H2(g)|Pt

Electrochemical Cells

The difference in electrical potential between the anode and cathode is called:

000reductionoxidationCell EEE

•cell voltage

• electromotive force (emf)

• cell potential

Standard Electrode Potentials

Standard reduction potential (E0) 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)

eatm

Reduction Reaction

Determining if Redox Reaction is Spontaneous

• + E°CELL ; spontaneous reaction

• E°CELL = 0; equilibrium• - E°CELL;

nonspontaneous reactionMore positive E°CELL ; stronger oxidizing agent ormore likely to be reduced

Relating E0Cell to G0

ech

workECell arg

Unitswork, Joulecharge, CoulombEcell; Volts

Faraday, F; charge on 1 mole e-F = 96485 C/mole

G = -nFEcell

Relating CELL to the

Equilibrium Constant, K

G0 = -RT ln K

G0 = -nFE0cell

-RT ln K = -nFE0cell

K

nF

RTECell ln0

0257.0

96485

29831.8

moleC

KmolKJ

F

RT

Kn

Kn

ECell log0592.0

ln0257.00

Corrosion – Deterioration of Metals by Electrochemical Process

Corrosion – Deterioration of Metals by Electrochemical Process

Corrosion – Deterioration of Metals by Electrochemical Process

CorrosionCorrosion

•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

CorrosionCorrosion

•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

CorrosionCorrosion

•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