Chapter 14 Oxidation-Reduction Reactions Malone and Dolter - Basic Concepts of Chemistry 9e2 Setting...

Chapter 14Chapter 14

Oxidation-Reduction ReactionsOxidation-Reduction Reactions

Malone and Dolter - Basic Concepts of Chemistry 9e

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Setting the Stage – Electron Flow

The flow of electrons is common. It occurs naturally (lightning) and as a

result of human activity (electricity). Early electrical experiments involved

chemically generated electricity (batteries).

This electron flow results from one reactant having a greater affinity for electrons than the other.


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Setting a Goal – Part ARedox Reactions – The Exchange of Electrons

You will learn of an important classification of chemical reactions that involves an exchange of electrons.


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Objective for Section 14-1

Calculate oxidation states and determine the species oxidized, the species reduced, the oxidizing agent, and the reducing agent in an electron exchange reaction.


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Professor’s Little Jokes

An electron is setting in jail with her room mate,who asks, “Why did they put you in prison?”The electron replies, “For making a forbidden transition”.What do two parting dipoles say to each other?Answer: “Debye”!A male polar bear and a female brown bear are sitting

at a bar. The polar bear says, “Sorry babe, I just don't think The chemistry is right”. Question: "If H2O is water, what is H2O4?" Answer:

“Drinking and washing, of course! Are you awake yet?”


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14-1 The Nature of Oxidation and Reduction and Oxidation States

Na reacts with Cl2 vigorously, such that the piece of Na glows white hot with the heat of the reaction.

The process forms ordinary table salt:2 Na(s) + Cl2(g) 2 NaCl(s)

Na + Na+ Cl_

Cl

electron transfer


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Half Reactions

Electron exchange reactions can be viewed as the sum of two half reactions.

Half reactions represent either the loss of electrons or the gain of electrons as a separate balanced equation.

Sodium half reaction:Na Na+ + e-

Na loses an electron to form a sodium ion. A substance that loses electrons is said to be

oxidized.


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Half Reactions

Chlorine half reaction2 e- + Cl2 2 Cl-

The neutral chlorine molecule has gained two electrons to form chloride ions.

A substance that gains electrons is said to be reduced.


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Terminology

Redox or oxidation-reduction reactions are reactions involving the exchange of electrons.

Oxidizing agent – the species that accepts electrons.

Reducing agent – the species that donates electrons.


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Redox Reactions

Electrons are conserved in redox reactions.

For example, electrons gained in the reduction process must equal the electrons lost in the oxidation process.

A common, and unfortunate, redox reaction is the rusting of iron:4 Fe(s) + 3 O2 (g) 2 Fe2O3 (s)


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Oxidation States or Numbers

The charge that an atom in a molecule or ion would have if all atoms were present as monatomic ions.

Similar to formal charge, but in this case, both electrons in the bond are assigned to the more electronegative atom sharing the electrons.


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Rules for Assigning Oxidation State

1. OS of an element in its free, natural state is zero [Cu(s), Cl2 (g), Hg(l)]

2. The OS of a monoatomic ion is the same as the charge of the ionalkali metals +1 same as group #alkaline earths +2 same as group #oxygen O2-

aluminum Al3+


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Rules for Assigning Oxidation State…..Cont’d

3. Halogens are in a -1 oxidation state in binary compounds whether ionic or covalent when bound to a less electronegative element (e.g. as in HCl).

4. Oxygen is O2- except in peroxides and superoxides. Oxygen is positive when bound to F.


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Rules for Assigning Oxidation State…..Cont’d

5. H is usually +1. When bound to a less electronegative atom (usually a metal) it is -1 (as in LiH).

6. The sum of the OS of all of the atoms in a compound must equal the charge of the compound or ion.e. g. FeO has Fe2+ and O2-


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Rules for Assigning Oxidation State…..Cont’d. Examples

Determine the oxidation state (number) of the bold atom ineach of the following.(a)MnO4

- (b) Cr2O72- (c) N2O4 (d) H2CO

Solutions

(a) Let OS of Mn be y, then y + (-2 x 4) = -1; y = 7 (Mn(VII))

(b) 2y + (-2 x 7) = -2; y = 6 (Cr(VI))

(c) 2y + (-2 x 4) = 0; y = 4

(d) y = (+1 x 2) + -2 = 0; y = 0


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Oxidation State Changes

Oxidation is the loss of electrons, and therefore results in an increase in the oxidation state. Reduction results in the decrease in the oxidation state.

Generally, only one element changes oxidation state in a compound, but this is not always true.


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Balance redox reactions by the bridge method.


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14-2 Balancing Redox Equations: Oxidation State Method

The oxidation state method focuses on the atoms of the elements undergoing a change in oxidation state.

ConsiderHNO3 (aq) + H2S(aq) NO(g) + S(s) + H2O(l)


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Oxidation State Method….cont’d

1. Identify the atoms whose OS have changed.

2. Draw a bridge between the same atoms whose OS has changed, indicating the electrons gained or lost.

HNO3 H2S NO+ + S + H2O0+2-2+5

HNO3 H2S NO+ + S + H2O0+2-2+5

+3e-

-2e-


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3. Multiply the two numbers of electrons (in this case +3 and -2) by whole numbers that produce a common number:3 × 2 = 6 2 × 3 = 6Use these multipliers as coefficients of the respective compounds or elements.


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4. Balance the rest of the elements by inspection2 HNO3 + 3 H2S 2 NO + 3 S + 4 H2O

HNO3 H2S NO+ + S + H2O0+2-2+5

+3e- x 2 = +6e-

-2e- x 3 = -6e-


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Balance redox reactions by the ion-electron method.


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14-3 Balancing Redox Equations: Ion-Electron Method

Ion-electron or half reaction method:

1. Separate the total reaction into half reactions.

2. Balance the half reactions separately.

3. Methodology depends on whether the reactions are done in acidic or basic conditions.


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Half Reaction Method

ConsiderH+(aq) + Cl-(aq) + Cr2O7

2-(aq) Cr3+(aq) + Cl2(aq) + H2O(l)

1. Separate the molecule or ion that contains atoms of an element that have changed oxidation state and product containing the atoms of that element. (You don’t need to know the actual oxidation state).


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Half Reaction Method…..Cont’d

1. Cr2O72- Cr3+ (reduction)

2. Balance the atoms other than hydrogen or oxygen.Cr2O7

2- 2Cr3+

3. Balance the oxygen by adding H2O on the side missing the oxygen.


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4. Cr2O72- 2Cr3+

+ 7 H2O

5. Balance the hydrogens by adding H+ on the other side of the equation from the H2O14 H+ + Cr2O7

2- 2Cr3+ + 7 H2O

6. The atoms in the half reaction should now be balanced.


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7. Add e- to balance the charge on both sides of the equation:6 e- + 14 H+ + Cr2O7

2- 2Cr3+ + 7 H2O

8. Do the same for other half reaction:2 Cl- Cl2 + 2 e- (oxidation)

9. Multiply each equation by a coefficient so that the number of electrons is the same in both half reactions.


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10. 3 (2 Cl- Cl2 + 2 e-) 6 Cl- 3 Cl2 + 6 e-

11. 6 e- + 14 H+ + Cr2O72- 2Cr3+

+ 7H2O6 Cl- 3 Cl2 + 6 e-

_Add________________________________ 14 H+ + Cr2O7

2- + 6 Cl- 3 Cl2 + 2Cr3+

+ 7 H2O


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Half Reaction Method –Alkaline Solution

In alkaline solution, OH- is the predominant species (besides water).

The simplest way to adjust for basic solution is to balance the reaction as if it occurred in acid, then neutralize the H+ by adding OH- to both sides.

The H+ will combine with the OH-.


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Example

1. 2 e- + 2 H+ + ClO- Cl- + H2O

2. 2 e- + (2 H+ + 2OH-) + ClO- Cl- + H2O + 2 OH-

3. 2 e- + 2 H2O + ClO- Cl- + H2O + 2 OH-

4. 2 e- + H2O + ClO- Cl- + 2 OH-

Half equation for the reduction of hypochlorite to chloride in alkaline solution


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Spontaneous Redox Reactions

A spontaneous reaction occurs between the stronger oxidizing agent and the stronger reducing agent to form weaker oxidizing and reducing agents.

By testing reactions between pairs of atoms and molecules, we can determine a relative order of strength of oxidizing or reducing ability.


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Setting a Goal – Part BSpontaneous and Nonspontaneous Redox Reactions

You will understand the extensive practical applications of redox reactions.


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Using a table of relative strengths of oxidizing agents, determine whether a specific redox reaction is spontaneous.


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14-4 Predicting Spontaneous Redox Reactions

Conventionally, we write redox reactions as reductions in order of increasing reducing ability (Table 14-1; most reducing at the bottom)

We can get quantitative values for the comparative strength of a given reducing agent. These are known as reduction potentials. The species are compared in solution under standard conditions (1.00 M concentration and 1.00 atm).


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Activity Series Table

Table 14-1 is the activity series table. A favorable reaction occurs between an

element, ion or compound on the left (an oxidizing agent) with a species on the right (a reducing agent) that lies below it in the table.


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Example: Reaction between Cl2 and H2O

Cl2 lies above H2O in Table 14-1 so it is a stronger oxidizing agent.Cl2 + 2 e- 2 Cl-

O2 + 4 H+ + 4 e- 2 H2O2 Cl2 + 2 H2O O2 + 4 H+ + 4 Cl- (slow but favorable)


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Example: Reaction between Na and H2O

Na lies below H2O in Table 14-1, so water will oxidize Na.2 H2O + 2 e- H2 + 2OH-

Na+ + e- Na 2 H2O + 2 Na H2 + 2OH-

+ 2 Na+

Active metals are those that react with water and therefore lie below water in the activity series.


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Example: Reaction between Fe and Acid solution

2 H+ + 2 e- H2

Fe2+ + 2 e- Fe 2 H+ + Fe H2 + Fe2+

Higher concentrations of acid make this reaction more favorable.

This is one of the problems with acid rain in that it accelerates the rusting of many metal structures.


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Objectives for Section 14-5

Describe the parts of a voltaic cell and how it generates electricity.

List the types of batteries discussed and a particular use or advantage of each type.


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14-5 Voltaic Cells

Use a favorable redox reaction to generate electrical energy through an external circuit.

First practical example was the Danielle cell.Zn(s) + Cu2+(aq) Zn2+ (aq) + Cu(s)It was used to power telegraphs and other early electrical devices.

To generate electricity, we must separate the half cells physically.


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The Daniell Cell

A Zn strip is immersed in a Zn2+ solution and in a separate compartment, a Cu strip is immersed in a Cu2+ solution.

Electrodes are the surfaces in a cell at which the reactions take place.

The Zn and Cu strips serve as electrodes. Refer to Figure 14-4.


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Figure 4.4 The Daniell Cell


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Electrodes

Anode – the electrode at which the oxidation takes place.

Cathode – the electrode at which the reduction takes place.

Active electrodes, such as the Zn and Cu strips, participate in the redox reaction.

Inert electrodes merely provide a surface at which the reaction occurs.


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Salt Bridge

Since negative charge is flowing through the external circuit, a means must be provided to keep charge balanced.

The salt bridge allows inert ions to pass between the two compartments while preventing the species involved in the redox reaction from mixing.


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Common Voltaic Cells

Dry cells (flashlight batteries) Lead-acid car batteries (keep in mind that

a battery is a collection of one or more separate cells joined together in one unit)


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Dry Cells

Not rechargeable, but cheap and portable. Zn casing serves as the anode and a

graphite rod serves as an inert cathode. In between the two is an aqueous paste

containing NH4Cl, MnO2 and carbon.anode: Zn(s) Zn2+(aq) + 2 e-

cathode: 2NH4+(aq) + 2MnO2(s) + 2 e-

Mn2O3(s) + 2 NH3(aq) + H2O(l)


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The Dry Cell

The disadvantage of the dry cell is that it creates an acidic solution.

This acidic solution slowly dissolves the zinc, so the shelf life of the battery is only a few months.

Alkaline batteries use NaOH or KOH in place of NH4Cl, which is more expensive but the shelf life is much greater. The anode reaction becomes:Zn(s) + 2 OH-(aq) ZnO(s) + H2O(l) + 2 e-


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The Lead Acid Car Battery

Anode:Pb(s) + H2SO4(aq) PbSO4(s) + 2H+(aq) + 2 e-

Cathode:2 e- + 2 H+(aq) + PbO2(s) + H2SO4(aq) PbSO4(s) + 2 H2O(l)

Total reaction: Pb(s) + PbO2(s) + 2 H2SO4(aq)

2 PbSO4(s) + 2 H2O(l)


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Other Batteries

Silver:Zn(s) + Ag2O(s) ZnO(s) + 2Ag(s)

Mercury:Zn(s) + HgO(s) ZnO(s) + Hg(l)

Ni/Cd:Cd(s) + NiO2(s) + 2 H2O(l) Ni(OH)2(s) + Cd(OH)2(s) (rechargeable)


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Examples of Batteries: Dry Cells

Figures 14-6 and 14-7


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Examples of Batteries: Lead-Acid

Fig. 14-5


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Fuel Cells: the Power Source of the Future

Uses the direct reaction of chemicals such as hydrogen and oxygen to produce electricity.

Must be used in applications where the electrical flow is continuous (no recharging).

Can also be turned on and off by controlling the flow of reagents.

Very expensive, but much research is going into changing that.


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A Fuel Cell

Fig. 14-8


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Objectives for Section 14-6

Write the reaction that occurs when a salt is electrolyzed.

Describe the various applications for electrolytic cells.


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14-6 Electrolytic Cells

Convert electrical energy into chemical energy

Electrolysis of H2O (using K2SO4) to yield H2 and O2

2 H2O(l) 2 H2(g) + O2(g) Electroplating is another electrolytic cell

that uses a base metal object (a spoon) as the cathode and a Ag bar as anode.

See Figures 14-9 and 14-10


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A Simple Electrolytic Cell

Fig. 14-9


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Electroplating

Fig. 14-10


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Electrosynthesis

Electrical energy is used to synthesize a wide range of commercial products

Na and Mg are common Al and Cl2 are probably the most

important


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Electrosynthesis of Sodium

Fig. 14-11

Chapter 14 Oxidation-Reduction Reactions Malone and Dolter - Basic Concepts of Chemistry 9e2 Setting...

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