Chapter 19 Oxidation–Reduction (Redox) Reactions

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Mark S. CracoliceEdward I. Peters

Mark S. Cracolice • The University of Montana

Chapter 19Oxidation–Reduction

(Redox) Reactions

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Electron Transfer Reactions

The reaction between zinc and copper(II) ions is spontaneous.


Zinc is said to be oxidized to zinc ion and copper(II) ions is said to be reduced to copper.


The same reaction between zinc and copper(II) ions can be performed using the following arrangement.

Electron Transfer ReactionsWhen the two metallic rods are connected, zinc is oxidized to zinc

ions and releases electrons at the zinc electrode and copper (II) ions collect electrons from the copper electrode to become copper.


The overall oxidation-reduction reaction can be separated into two half reactions.

Zn (s) Zn2+ (aq) + 2e- ( at the zinc rod)

Cu2+ (aq) + 2e- Cu (s) ( at the copper rod)_________________________________________

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


The process of loosing electrons is called oxidation.

The chemical change that occurs at the zinc electrode is oxidation, zinc is oxidized to Zn2+.

Zn(s) Zn2+(aq) + 2 e–

The process of gaining electrons is called reduction.

Copper (II) cations are reduced to Cu at the copper rod.Cu2+(aq) + 2 e– Cu(s)

Electron Transfer ReactionsIf half-reaction equations are added algebraically, the result is

the net ionic equation for the oxidation–reduction reaction:

Zn(s) Zn2+(aq) + 2 e–

Cu2+(aq) + 2 e– Cu(s)

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

The chemical change is an electron transfer reaction; electrons have been transferred from zinc atoms to copper(II) ions.

Voltaic (Galvanic) Cells

Voltaic (Galvanic) CellA cell in which a spontaneous chemical reaction is used

to produce electricity.

Voltaic (Galvanic) CellsAnode: Electrode at which oxidation occurs. The species being

oxidized releases electrons into the electrode.Cathode: Electrode at which reduction occurs. The species

being reduced collect electrons from the electrode

Commercial Galvanic CellsAlkaline battery One electrode is zinc in contact with a gel of potassium

hydroxide.One electrode is graphite in contact with a mixture of manganese

(IV) oxide, graphite and potassium hydroxide.

Reaction at the anode:Zn (s) + 2 OH- (aq) ZnO (s) + H2O (l) + 2e

Reaction at the cathode:2 MnO2 (s) + H2O (l) + 2e Mn2O3 (s) + 2 OH-

Commercial Galvanic CellsLead battery One electrode is lead in contact with a sulfuric acid solution One electrode is lead in contact with lead (IV) oxide and sulfuric

acid

Reaction at the anode:Pb (s) + SO4

2- (aq) PbSO4 (s) + 2e

Reaction at the cathode:PbO2 (s) + 4 H+ + SO4

2- (aq) + 2e PbSO4 (s) + 2 H2O

Lead battery is rechargeable.

Electrolytic CellsElectrolytic Cell

In an electrolytic cell, current supplied by an external source is used to drive a non-spontaneous redox reaction

ComponentsAn electrolyte and two electrodes.

When the electrodes are connected to an outside source of electricity, they become charged, one positively and one

negatively. The ions in the electrolyte move to the oppositely charged electrode, where chemical reactions occur. The

movement of ions is an electric current in solution.

Electrolytic CellsElectrolysis of liquid sodium chloride. Two electrodes are connected to an external electric source and dipped in liquid sodium chloride.

Electrolytic CellsAt Negative Electrode: sodium ions Na+ are attracted to the negative electrode, accept electrons from the electrode and are reduced to sodium. The reaction is reduction so the electrode is called cathode.

Na+ + e- Na

At Positive Electrode: Chloride ions Cl- are attracted to the positive electrode, where they are oxidized to chlorine gas.

The reaction is oxidation, so the electrode is an anode.

2 Cl- + 2 e- Cl2

Oxidation Numbers & Redox Reactions

Electron-transfer reaction is called redox reaction.

The reactant that gains electrons is called the oxidizing agent.

The reactant that loses electrons is called the reducing agent.

In the reaction, the oxidizing agent is reduced, and the reducing agent oxidized.

To determine which species has gained and which has lost electrons, it is useful to determine the oxidation states, or

oxidation numbers of reactants and products.

Rules for Determination of Oxidation NumbersThe oxidation number of any elemental substance is 0The oxidation number of a monatomic ion is the same as the

charge on the ionThe oxidation number of combined oxygen is -2, except in

peroxide(-1), superoxide(-1/2), and OF2 (+2)The oxidation number of combined hydrogen is +1(except

hydride ion , H- ).In any molecular or ionic species the sum of the oxidation

numbers of all atoms in a molecule or polyatomic ion is equal to the charge on the species.

Oxidation is an increase in oxidation number.Reduction is a decrease in oxidation number.•

Oxidizing and Reducing AgentsOxidizing Agent (Oxidizer)

Species that removes electrons in a redox reaction(the species that is reduced).

Reducing Agent (Reducer)Species from which electrons are removed

(the species that is oxidized).

Strengths of Ox and Red AgentsStrong oxidizing agent

A species with a strong attraction for electrons.

Weak oxidizing agentA species that attracts electrons only slightly.

Strong reducing agentA species that releases electrons readily.

Weak reducing agentA species that does not easily give up its electrons.

Strengths of Oxidizing and Reducing Agents

Predicting Redox Reactions

Writing Redox EquationsWriting Redox Equations for

Half-Reactions in Acidic Solutions

1. After identifying the element oxidized or reduced, write a skeleton half-reaction equation with the element in its original form (element, monatomic ion, or part of a polyatomic ion or compound) on the left and its final form on the right.

2. Balance the element oxidized or reduced.

3. Balance elements other than hydrogen or oxygen, if any.

Writing Redox EquationsWriting Redox Equations for

Half-Reactions in Acidic Solutions

4. Balance oxygen by adding water molecules where necessary.

5. Balance hydrogen by adding H+ where necessary.

6. Balance charge by adding electrons to the more positive side.

7. Recheck the equation to be sure it is balanced in both number of atoms and charge.

Writing Redox EquationsExample:Iron(II) ion solution and a permanganate ion (MnO4

–) solution are combined in an acidic solution. Manganese(II) ion and iron(III) ion are produced. Write a balanced net ionic equation for the reaction.

Solution:First, we summarize the reaction description:Reactants: Fe2+(aq) and MnO4

–(aq)The reaction is in acidic solution, so H+(aq) is also a reactantProducts: Fe3+(aq) and Mn2+(aq)

Writing Redox EquationsReactants: Fe2+(aq), MnO4

–(aq), H+(aq) Products: Fe3+(aq), Mn2+(aq)

Solution:Step 1: Write skeleton half-reaction equations

Fe2+(aq) Fe3+(aq)MnO4

–(aq) Mn2+(aq)



Solution:Step 2: Balance element oxidized/reduced

Fe2+(aq) Fe3+(aq) MnO4–(aq) Mn2+(aq)

Both Fe and Mn are already balanced



Solution:Step 3: Balance elements other than H, O

Fe2+(aq) Fe3+(aq) MnO4–(aq) Mn2+(aq)

There are none



Solution:Step 4: Balance O by adding H2O

Fe2+(aq) Fe3+(aq) MnO4–(aq) Mn2+(aq) + 4 H2O(l)



Solution:Step 5: Balance H by adding H+

Fe2+(aq) Fe3+(aq) 8 H+ + MnO4–(aq) Mn2+(aq) + 4 H2O(l)



Solution:Step 6: Balance charge by adding electrons

Fe2+(aq) Fe3+(aq) + e–

5 e– + 8 H+ + MnO4–(aq) Mn2+(aq) + 4 H2O(l)



Solution:Step 7: Check for charge and atom balance

Fe2+(aq) Fe3+(aq) + e–

2+ each side, 1 Fe each side5 e– + 8 H+ + MnO4

–(aq) Mn2+(aq) + 4 H2O(l)2+ each side, 8 H, 1 Mn, 4 O each side



Solution:Add the half reactions to yield the complete net ionic equation

5 × [Fe2+(aq) Fe3+(aq) + e–]5 e– + 8 H+ + MnO4

–(aq) Mn2+(aq) + 4 H2O(l)

5 Fe2+(aq) + 8 H+ + MnO4–(aq) 5 Fe3+(aq) + Mn2+(aq) + 4 H2O(l)

Homework• Homework: 3, 9, 11, 17, 19, 21, 23,25, 29, 49, 59• Write net equations for all reactions of the Experiment 13

Chapter 19 Oxidation–Reduction (Redox) Reactions

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