chapter 12 oxidation-reduction reactions

Chapter 12Oxidation-Reduction Reactions

In chapter 5, we learned how to recognize an oxidation-reduction reaction. In this chapter we introduce electrode potentials and the Nernst equation. These allow for a quantitative treatment of electrochemistry.

Common Oxidation-Reduction Reactions

● Oxidation-reduction reactions used for heat or work:

Combustion

Metabolic ● Corrosion● Photosynthesis


● Oxidation-reduction reactions involve the transfer of electrons.

Element or compound that gains electrons undergoes reduction.

Element or compound that loses electrons undergoes oxidation.


● Consider this reaction:

2Na + Cl2 → 2NaCl

● The Na has been oxidized.● The Cl2 has been reduced.


● This reaction, 2Na + Cl2 → 2NaCl, can be written as the sum of two half-reactions:

2Na ⇄ 2Na+ + 2e- oxidation

Cl2 + 2e- ⇄ 2Cl- reduction


● The addition of oxygen atoms or hydrogen atoms to an element or compound is also classified as an oxidation-reduction reaction.


● CO2 + H2 ⇄ CO + H2O

The H2 is oxidized and the CO2 reduced.

● C2H4 + H2 ⇄ C2H6

The C2H4 is reduced and the H2 oxidized.


● In order to determine that an oxidation-reduction reaction has occurred, we must be able to assign oxidation numbers or oxidation states.

Determining Oxidation Numbers

● Introduced in Chapter 5, section 16● Two methods:

One based on Lewis structure– Good for organic compounds

Other based on set of rules– e. g., elements = 0, monatomic ions = charge– Review section 16, chapter 5

Recognizing Oxidation-Reduction Reactions

● After all the oxidation numbers in a chemical reaction have been determined, look for changes.

Oxidation occurs when the oxidation number of an atom increases.

Reduction occurs when the oxidation number of an atom decreases.


● In an oxidation-reduction reaction, both oxidation and reduction must occur.

If one species is being oxidized, another must be reduced.

● In biochemical reactions, often only the oxidation or reduction reaction is shown explicitly.


Figure 12.1


● Organic reactions can be classified by examining the Lewis structures.

If the number of C-H bonds decreases, the molecule is being oxidized.

If the number of C-O bonds increases, the molecule is being oxidized.

Conversely, if the number of C-H bonds increases, the molecule is being reduced.

Voltaic Cells

● Also known as galvanic cells● Physically separate the half-reactions● Force electrons to travel through an

external circuit connecting the two half-reactions

Battery!

Voltaic Cells

Figure 12.2

Voltaic Cells

● As H+ ions leave the solution on the right, K+ ions fill in to keep the solution electrically neutral.

Salt Bridge● The voltage required to prevent the flow

of electrons is measured with a voltmeter.This voltage is called the cell potential.

Voltaic Cells

● Oxidation takes place at the anode.● Reduction takes place at the cathode.● In Figure 12.2, the half-reactions involve

two electrons.The half-reactions are added to produce the overall reaction.

Voltaic Cells

● What if the half reactions do not have the same number of electrons?

Figure 12.3

Voltaic Cells

2+ -

+ -

+ 2+

Cu( ) Cu ( ) + 2e

2× Ag ( ) + e Ag( )

_______________________________

Cu( ) + 2Ag ( ) Cu ( ) + 2Ag( )

s aq

aq s

s aq aq s

Standard Cell Potentials

● The relative half-reactions from Figure 12.2 are

Zn ⇄ Zn+2 +2e- E° =+0.7628

2H+ + 2e- ⇄ H2 E° =+0.0000

Figure 12.2

Standard Cell Potentials

● The overall standard cell potential, E°, for the cell is the sum of the two half-reaction E°.

● Expect the reaction to go as written if the overall E° >0.

Oxidizing and Reducing Agents

● Reducing agent donates electrons: its oxidation number increases.

● Oxidizing agent accepts electrons: its oxidation number decreases.


● In Figure 12.2, zinc metal is the reducing agent.

● Hydrogen ions are the oxidizing agent.

Figure 12.2


● As with acids and bases, there are conjugate oxidizing and reducing agents.

When Zn is oxidized to Zn+2, Zn+2 becomes the conjugate oxidizing agent because its oxidation number drops in the reverse reaction:

Zn ⇄ Zn+2 + 2e-


● Strong reducing agents produce weak conjugate oxidizing agents.

● Strong oxidizing agents produce weak conjugate reducing agents.

Relative Strengths of Oxidizing Agents and Reducing Agents

● Oxidation-reduction reactions should occur when they convert the stronger of a pair of oxidizing agents and the stronger of a pair of reducing agents into a weaker oxidizing agent and a weaker reducing agent.


Table 12.1


● Standard electrode potentials, E°red

● Half-reactions written as reductions● Standard means gases at 1 bar, solutions at

1 M● When written as oxidations, the sign on

E°red is reversed.

Batteries

● Alkaline dry cells, ubiquitous● Lead-Acid, cars● NiCd, rechargeable● NiMH, hybrids● Lithium ion, compact● Fuel cells, hydrogen

Batteries

● Lead-Aciddischarge+ - -

2 4 4 2charge

discharge- + -4 4charge

PbO ( ) + 3H ( ) + HSO ( ) + 2e PbSO ( ) + 2H O( )

Pb( ) + HSO ( ) PbSO ( ) + H ( ) + 2e

s aq aq s l

s aq s aq

Batteries

● NiMH

discharge- -2charge

discharge- -2 2charge

MH + OH M + H O + e

NiO(OH) + H O + e Ni(OH) OH

Electrochemical Cells at Nonstandard Conditions: The Nernst Equation● To determine E when a cell is not at

standard conditions, the Nernst Equation is used.

cQnF

RTEE ln0

Electrochemical Cells at Nonstandard Conditions: The Nernst Equation

● n is the number of electrons transferred.

● Qc is the reaction quotient.

Notice that if all concentrations are 1 M, E=E°.

● F is the Faraday constant.

cQnF

RTEE ln0

Electrochemical Cells at Nonstandard Conditions: The Nernst Equation

Zn(s) + Cu+2(aq) → Zn+2(aq) + Cu(s)

Figure 12.8

Electrochemical Cells at Nonstandard Conditions: The Nernst Equation● At equilibrium Qc = K and E = 0.

● This provides an alternate equation for expressing equilibrium.

RTnFEc eK /0

Electrolysis and Faraday’s Law

● Voltaic cells operate spontaneously.● Electrolytic cells require an external power

supply.e. g. electroplating


Figure 12.9


● The amount of a substance consumed or produced at one of the electrodes in an electrolytic cell is directly proportional to the amount of electricity that passes through the cell.


● Amps × time (in secs) = Coulombs, C● F = 96,485 C/mol of e-

● C/F = mol of e- passed● Grams of silver plated out can be

determined from

[Ag(CN)2]-(aq) + e- ⇄ Ag(s) + 2CN-(aq)

Electrolysis of Molten NaCl

Figure 12.11

+ -

- -2

Na + e Na

2Cl Cl + 2e

Electrolysis of Molten NaCl

● CaCl2 added to the NaCl to lower the melting point. No effect on half reactions.

● Na(l) less dense than NaCl(l).● Cl2(g) and Na(l) kept apart. Why?

Electrolysis of Aqueous NaCl

Figure 12.13

- -2 2

- -2

2H O( ) + 2e H ( ) + 2OH ( )

2Cl Cl ( ) + 2e

l g aq

g

Electrolysis of Aqueous NaCl

● Chloride is oxidized instead of water.● Water is reduced, not sodium ion.● Hydrogen gas and NaOH(aq) are produced

and sold.

Electrolysis of Water

Figure 12.15

- -2 2

+ -2 2

2H O + 2e H 2OH

2H O O + 4H + 4e

Electrolysis of Water

● A salt which resists electrolysis is added to improve conductivity.

● Similar (but not exactly) half reactions running in reverse describe a fuel cell.

● If the gases were collected, what would their volume ratio be?

The Hydrogen Economy

● Using hydrogen gas as a common fuel.Solar energy for electrolysis of water.Fuel cells to generate electricity from hydrogen and reproduce the water.

– No drain on fossil fuels.– No carbon emissions.– Water (seawater) already has the electrolyte for

improved conductivity added!– Plenty of solar radiation.

The Hydrogen Economy

● Not a new idea.● Big challenges.

Economical production of hydrogen.

Storage.

Distribution.

Better fuels cells.– Cheaper.– More robust.

chapter 12 oxidation-reduction reactions

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