basic principles of chemistry online southeast missouri state university cape girardeau, mo...

Basic Principles of Chemistry OnlineSoutheast Missouri State University

Cape Girardeau, MO

Introductory Chemistry, 3rd EditionNivaldo Tro

Chapter 16Oxidation and

Reduction

2009, Prentice Hall

Tro's Introductory Chemistry, Chapter 16

2

Oxidation–Reduction Reactions• Oxidation–reduction reactions are also called redox

reactions.• All redox reactions involve the transfer of electrons

from one atom to another.• Spontaneous redox reactions are generally

exothermic, and we can use their released energy as a source of energy for other applications.Convert the heat of combustion into mechanical energy to

move our cars.Use electrical energy in a car battery to start our car engine.


3

Combustion Reactions• Combustion reactions are always exothermic.

• In combustion reactions, O2 combines with all the elements in another reactant to make the products.

4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) + energy

CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) + energy


4

Reverse of Combustion Reactions• Since combustion reactions are exothermic, their

reverse reactions are endothermic.

• The reverse of a combustion reaction involves the production of O2.

energy + 2 Fe2O3(s) → 4 Fe(s) + 3 O2(g)

energy + CO2(g) + 2 H2O(g) → CH4(g) + 2 O2(g)

• Reactions in which O2 is gained or lost are redox reactions.


5

Oxidation and Reduction:One Definition

• When an element attaches to an oxygen during the course of a reaction it is generally being oxidized. In CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g), C is being

oxidized in this reaction, but H is not.

• When an element loses an attachment to oxygen during the course of a reaction, it is generally being reduced. In 2 Fe2O3(s) → 4 Fe(s) + 3 O2(g), the Fe is being reduced.

• One definition of redox is the gain or loss of O, but it is not the best.


6

Another Oxidation–Reduction

• Consider the following reactions:

4 Na(s) + O2(g) → 2 Na2O(s)

2 Na(s) + Cl2(g) → 2 NaCl(s)

• The reaction involves a metal reacting with a nonmetal.• In addition, both reactions involve the conversion of

free elements into ions.

4 Na(s) + O2(g) → 2 Na+2O–(s)

2 Na(s) + Cl2(g) → 2 Na+Cl–(s)


7

Oxidation and Reduction:Another Definition

• In order to convert a free element into an ion, the atoms must gain or lose electrons.Of course, if one atom loses electrons, another must accept

them.

• Reactions where electrons are transferred from one atom to another are redox reactions.

• Atoms that lose electrons are being oxidized, atoms that gain electrons are being reduced.

2 Na(s) + Cl2(g) → 2 Na+Cl–(s)Na → Na+ + 1 e– (oxidation)Cl2 + 2 e– → 2 Cl– (reduction)

Leo

Ger


8

Practice—Identify the Element Being Oxidized and the Element Being

Reduced.

• 2 C(s) + O2(g) → 2 CO(g)

• Mg(s) + Cl2(g) → MgCl2(s)

• Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s)


9

Practice—Identify the Element Being Oxidized and the Element Being

Reduced, Continued.

• 2 C(s) + O2(g) → 2 CO(g)


• Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s)

C is oxidized because it is gaining an attachment to O. O is reduced; there has to be reduction and it’s the only other element.

Mg is oxidized because it is becoming a cation by losing electrons. Cl is reduced because it is becoming an anion by gaining electrons.

0 0 2+ −

Mg is oxidized because it is becoming a cation by losing electrons. Fe2+ is reduced because it is gaining electrons to become neutral.


10

Oxidation–Reduction• Oxidation and reduction must occur simultaneously.

If an atom loses electrons, another atom must take them.

• The reactant that reduces an element in another reactant is called the reducing agent.The reducing agent contains the element that is oxidized.

• The reactant that oxidizes an element in another reactant is called the oxidizing agent.The oxidizing agent contains the element that is reduced.

2 Na(s) + Cl2(g) → 2 Na+Cl–(s)Na is oxidized, Cl is reduced.

Na is the reducing agent, Cl2 is the oxidizing agent.


11

Practice—Identify the Oxidizing and Reducing Agents.

• 2 C(s) + O2(g) → 2 CO(g)


• Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s)

C is oxidized because it is gaining attachment to O. O is reduced; there has to be reduction and it’s the only other element.

Mg is oxidized because it is becoming a cation by losing electrons. Cl is reduced because it is becoming an anion by gaining electrons.

0 0 2+ −

Mg is oxidized because it is becoming a cation by losing electrons. Fe2+ is reduced because it is gaining electrons to become neutral.


12

Practice—Identify the Oxidizing and Reducing Agents, Continued.

• 2 C(s) + O2(g) → 2 CO(g)


• Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s)

C is the reducing agent because it contains the element that is oxidized. O is the oxidizing agent because it contains the element that is reduced.

0 0 2+ −

Mg is the reducing agent because it contains the element that is oxidized. Cl2 is the oxidizing agent because it contains the element that is reduced.

Mg is the reducing agent because it contains the element that is oxidized. Fe2+ is the oxidizing agent because it contains the element that is reduced.


13

Electron Bookkeeping• For reactions that are not metal + nonmetal, or do

not involve O2, we need a method for determining how the electrons are transferred.

• Chemists assign a number to each element in a reaction called an oxidation state that allows them to determine the electron flow in the reaction.Although they look like them, oxidation states are not

ion charges!Oxidation states are imaginary charges assigned based on

a set of rules. Ion charges are real, measurable charges.


14

Rules for Assigning Oxidation States

• Rules are in order of priority.

1. Free elements have an oxidation state = 0. Na(s) = 0 and Cl2(g) = 0 in 2 Na(s) + Cl2(g)

2 NaCl(s).

2. Monoatomic ions have an oxidation state equal to their charge.

Na = +1 and Cl = -1 in NaCl(s).

3. a. The sum of the oxidation states of all the atoms or ions in a compound is 0.

Na = +1 and Cl = -1 in NaCl, and (+1) + (-1) = 0.


15

Rules for Assigning Oxidation States, Continued

3. b. The sum of the oxidation states of all the atoms in a polyatomic ion equals the charge on the ion.

N = +5 and O = -2 in NO3–, (+5) + 3(-2) = -1.

4. a. Group I metals have an oxidation state of +1 in all their compounds.

Na = +1 in NaCl.

b. Group II metals have an oxidation state of +2 in all their compounds.

Mg = +2 in MgCl2.


16

Rules for Assigning Oxidation States, Continued

5. In their compounds, nonmetals have oxidation states according to the table below.

Nonmetals higher on the table take priority.Nonmetal Oxidation state Example

F -1 CF4

H +1 CH4

O -2 CO2

Group 7A -1 CCl4

Group 6A -2 CS2

Group 5A -3 NH3


17

Practice—Assign an Oxidation State to Each Element in the Following:

• F2

• Mg2+

• KCl

• SO2

• PO43−

• BaO2


18

Practice—Assign an Oxidation State to Each Element in the Following,

Continued:

• F2 F = 0 (Rule 1)

• Mg2+ Mg = +2 (Rule 2)

• KCl K = +1 (Rule 4a) and Cl = -1 (Rule 5)

• SO2 O = -2 (Rule 5) and S = +4 (Rule 3a)

• PO43− O = -2 (Rule 5) and P = +5 (Rule 3b)

• BaO2 Ba = +2 (Rule 4b) and O = -1 (Rule 3a)


19

Oxidation and Reduction:A Better Definition

• Oxidation occurs when an atom’s oxidation state increases during a reaction.

• Reduction occurs when an atom’s oxidation state decreases during a reaction.

CH4 + 2 O2 → CO2 + 2 H2O-4 +1 0 +4 –2 +1 -2

oxidationreduction


20

Practice—Assign Oxidation States and Identify the Oxidizing and Reducing Agents

in Each of the Following:

• 3 H2S + 2 NO3– + 2 H+ S + 2 NO + 4 H2O

• MnO2 + 4 HBr MnBr2 + Br2 + 2 H2O


21

• 3 H2S + 2 NO3– + 2 H+ S + 2 NO + 4 H2O

• MnO2 + 4 HBr MnBr2 + Br2 + 2 H2O

+1 -2 +5 -2 +1 0 +2 -2 +1 -2

oxidizing agentreducing agent

+4 -2 +1 -1 +2 -1 0 +1 -2

oxidationreduction

oxidation

reduction

reducing agentOxidizing agent

Practice—Assign Oxidation States and Identify the Oxidizing and Reducing Agents

in Each of the Following, Continued:


22

Will a Reaction Take Place?

• Reactions that are energetically favorable are said to be spontaneous.They can happen, but the activation energy may

be so large that the rate is very slow.

• The relative reactivity of metals can be used to determine if some redox reactions are spontaneous.


23

Single Displacement Reactions• Also known as single replacement reactions.• A more active free element displaces a less active

element in a compound.Metals displace metals or H.

Cu + 2 AgNO3 Cu(NO3)2 + 2 Ag

Mg + 2 HCl MgCl2 + H2

Nonmetals displace nonmetals.

2 KI + Br2 2 KBr + I2

Carbon displaces metals from oxides.

3 C + Fe2O3 3 CO + 2 Fe

• Always redox.

Tendency to Lose Electrons• Some metals have a greater tendency to lose

electrons than others.Metallic-free elements are always oxidized.The greater the tendency of a metal to lose electrons,

the easier it is to oxidize.The greater the tendency of a metal to lose electrons,

the harder it is to reduce its cations.

• If Metal A has a greater tendency to lose electrons than Metal B, then:

A(s) + B+(aq) A+(aq) + B(s),but: A+(aq) + B(s) no reaction.


24

25

KBaSrCaNaMgAlMnZnCrFeCdCoNiSnPbHSbAsBiCuHgAgPdPtAu

displace H2

fromcoldH2O

fromsteam

fromacids

reac

t wit

h O

2 in

the

air

to m

ake

oxid

es

Fe is above Cu, so Cu metalwill not displace Fe2+


displace H2

fromcoldH2O

fromsteam

fromacids

reac

t wit

h O

2 in

the

air

to m

ake

oxid

es

Gold is at thebottom, so it isvery unreactive.


displace H2

fromcoldH2O

fromsteam

fromacids

reac

t wit

h O

2 in

the

air

to m

ake

oxid

es

Zn is above H,so Zn will react with acids

Zn + Fe2+ Fe + Zn2+

Activity Series of Metals• Listing of metals by

reactivity.• Free metal higher on the

list displaces metal cation lower on the list.

• Metals above H will dissolve in acid:

Cu + Fe2+ no reactionZn + 2 H+ H2 + Zn2+


displace H2

fromcoldH2O

fromsteam

fromacids

reac

t wit

h O

2 in

the

air

to m

ake

oxid

es

Fe is below Zn, so Zn metalwill displace Fe2+.


26

Mg is aboveCu on theactivity series.Mg will react with

Cu2+ to form Mg2+

and Cu metal.

Cu will not react with Mg2+.

Table of Oxidation Half-Reactions

27

Table of Oxidation Half-Reactions, Continued

• Any oxidation half-reaction that is higher on the list will give a spontaneous reaction when combined with the reverse of a half-reaction that is lower on the list.The reverse of an oxidation half-reaction is a

reduction half-reaction.

• Metals will dissolve in acid if their oxidation half-reaction is above H2 2H++ 2e−.


28

Electrical Current• When we talk about the current

of a liquid in a stream, we are discussing the amount of water that passes by in a given period of time.

• When we discuss electric current, we are discussing the amount of electric charge that passes a point in a given period of time.Whether as electrons flowing

through a wire or ions flowing through a solution.


29

Redox Reactions and Current• Redox reactions involve the transfer of

electrons from one substance to another.• Therefore, redox reactions have the

potential to generate an electric current.• In order to use that current, we need to

separate the place where oxidation is occurring from the place that reduction is occurring.


30

Electric Current Flowing Directly Between Atoms

31

Electric Current Flowing Indirectly Between Atoms


32


33

Electrochemical Cells• Electrochemistry is the study of redox reactions that

produce or require an electric current.• The conversion between chemical energy and electrical

energy is carried out in an electrochemical cell.• Spontaneous redox reactions take place in a voltaic cell.

Also known as galvanic cells.Batteries are voltaic cells.

• Nonspontaneous redox reactions can be made to occur in an electrolytic cell by the addition of electrical energy.


34

Electrochemical Cells, Continued• Oxidation and reduction reactions kept separate.

Half-cells.

• Electron flow through a wire, along with ion flow through a solution, constitutes an electric circuit.

• Requires a conductive solid (metal or graphite) electrode to allow the transfer of electrons.Through external circuit.

• Ion exchange between the two halves of the system.Electrolyte.


35

Electrodes• Anode

Electrode where oxidation occurs.Anions attracted to it.Connected to positive end of battery in

electrolytic cell.Loses weight in electrolytic cell.

• CathodeElectrode where reduction occurs.Cations attracted to it.Connected to negative end of battery in

electrolytic cell.Gains weight in electrolytic cell.

Electrode where plating takes place in electroplating.

36

Voltaic Cell


37

Current and Voltage

• The number of electrons that flow through the system per second is the current.Electrode surface area dictates the number of

electrons that can flow.

• The amount of force pushing the electrons through the wire is the voltage.The farther the metals are separated on the

activity series, the larger the voltage will be.


38

Current

The amount of water that passesa point each secondis called the currentof the river.

The number of electrons that passa point each secondis called the currentof the electricity.


39

Voltage

Gravity is the force pullingthe water downthe river.

Voltage is the force pushingthe electrons down the wire.


40

Dead Battery

As the reactionproceeds, thereactants getconsumed andthe voltaic cell“dies.” The current decreasesuntil electronscan no longerflow throughthe wire.


41

LeClanché’s Acidic Dry Cell• Electrolyte in paste form.

ZnCl2 + NH4Cl.Or MgBr2.

• Anode = Zn (or Mg).Zn(s) Zn2+(aq) + 2 e-

• Cathode = graphite rod.• MnO2 is reduced.

2 MnO2(s) + 2 NH4+(aq) + 2 H2O(l) + 2 e- 2 NH4OH(aq) + 2 Mn(O)OH(s)

• Cell voltage = 1.5 v.• Expensive, nonrechargeable, heavy, easily

corroded.


42

Alkaline Dry Cell• Same basic cell as acidic dry cell, except

electrolyte is alkaline KOH paste.

• Anode = Zn (or Mg).Zn(s) Zn2+(aq) + 2 e-

• Cathode = brass rod.

• MnO2 is reduced.2 MnO2(s) + 2 NH4

+(aq) + 2 H2O(l) + 2 e- 2 NH4OH(aq) + 2 Mn(O)OH(s)

• Cell voltage = 1.54 v.

• Longer shelf life than acidic dry cells and rechargeable; little corrosion of zinc.

Tro's Introductory Chemistry, Chapter 16 43

Lead Storage Battery• Six cells in series.

• Electrolyte = 6 M H2SO4.

• Anode = Pb.Pb(s) + SO4

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

• Cathode = Pb coated with PbO2.

• PbO2 is reduced.PbO2(s) + 4 H+(aq) + SO4

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

• Cell voltage = 2.09 v.

• Rechargeable, heavy.

44

Fuel Cells• Like batteries in which reactants are constantly being

added.So it never runs down!

• Anode and cathode both Pt-coated metal.• Electrolyte is OH– solution.• Anode reaction: 2 H2 + 4 OH– → 4 H2O(l) + 4 e-.• Cathode reaction: O2 + 4 H2O + 4 e- → 4 OH–.


45

Nonspontaneous Redox Reaction

• The reverse of a spontaneous reaction is nonspontaneous.

• To get it to run, an outside energy source must be supplied.

• Nonspontaneous redox reactions can be made to work by using a battery to force the electrons to flow in the nonspontaneous direction.


46

Electrolysis • Electrolysis is the process of using

electricity to break a compound apart.

• Electrolysis is done in an electrolytic cell.

• Electrolytic cells can be used to separate elements from their compounds.Generate H2 from water for fuel cells.Recover metals from their ores.


47

Electrolytic Cell• The + terminal of the battery = anode.• The - terminal of the battery = cathode.• Cations attracted to the cathode; anions attracted to the

anode.• Cations pick up electrons from the cathode and are

reduced; anions release electrons to the anode and are oxidized.

• In electroplating, the work piece is the cathode.Cations are reduced at the cathode and plate onto the surface.The anode is made of the plate metal, the anode oxidizes and

replaces the metal cations lost from the solution.


48

Electrolytic Cell—Electroplating


49

Corrosion• Corrosion is the spontaneous oxidation of a

metal by chemicals in the environment.• Since many materials we use are active metals,

corrosion can be a very big problem.


50

Preventing Corrosion• One way to reduce or slow corrosion is to coat

the metal surface to keep it from contacting corrosive chemicals in the environment.Paint.Some metals, like Al, form an oxide that strongly

attaches to the metal surface, preventing the rest from corroding.

• Another method to protect one metal is to attach it to a more reactive metal that is cheap.Sacrificial electrode.

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