Chapter six Electrochemistry. Oxidation NumbersOxidation Numbers Oxidation-reduction reactionOxidation-reduction reaction Oxidizing agent and reducing

Chapter sixChapter six

ElectrochemistryElectrochemistry

• Oxidation NumbersOxidation Numbers

• Oxidation-reduction reactionOxidation-reduction reaction

• Oxidizing agent and reducing Oxidizing agent and reducing agentagent

• Redox coupleRedox couple

6-1 Oxidation-reduction Concepts

● Oxidation Numbers

─ The charge on an atom or a monatomic ion

─ The charge that an atom in a substance would

have if the shared pair of electrons belonged

to the more electronegative atom in the bond

Rules for Assigning an Rules for Assigning an Oxidation Number (Ox#)Oxidation Number (Ox#)

1. 1. For an atomFor an atom in its elemental form in its elemental form Ox # = 0:Ox # = 0:

(O(O22,Cl,Cl22, H, H22, ) (not combine with different element), ) (not combine with different element)

2. . For a monatomic ion: For a monatomic ion: Ox # = ion chargeOx # = ion charge CaCa2+2+, +2; Br +2; Br--, -1;, -1;

Ox. # = charge of molecule or ion.Ox. # = charge of molecule or ion. Sum of oxidation states = 0 in neutral compoundsSum of oxidation states = 0 in neutral compounds

General rules General rules

Sum of oxidation states = charge of the ionSum of oxidation states = charge of the ion

HH22SOSO44, Cr, Cr22OO772-2-

(1) (1) For oxygen: For oxygen: Ox# = -2Ox# = -2 in most compounds in most compounds

Ox# = -1 in peroxides (HOx# = -1 in peroxides (H22OO22) )

4. Rules for specific atoms or periodic table groups

(2) (2) ForFor hydrogen hydrogen: : Ox# = +1Ox# = +1 in combination with nonmetals in combination with nonmetals

Ox# = -1Ox# = -1 in combination with metals in combination with metals

with electropositive element (i.e., Na, K) H =-1with electropositive element (i.e., Na, K) H =-1

(3) (3) For fluorine: For fluorine: Ox# = -1Ox# = -1 in all compounds in all compounds

Except with F, K

OF2; KO2

(4) (4) For Group 1A: For Group 1A: Ox# = +1Ox# = +1 in all compounds in all compounds

(5) (5) For Group 2A: For Group 2A: Ox# = +2Ox# = +2 in all compounds in all compounds

(6) (6) For Group 7A: For Group 7A: Ox# = -1Ox# = -1 in most compounds in most compounds

x = +3 (for N)x = +3 (for N)x + 2(-2) = -1x + 2(-2) = -1NONO22--

x = +5 (for N)x = +5 (for N)X + 3(-2) = -1X + 3(-2) = -1NONO33--

x = +3 (for P)x = +3 (for P)x + 3(-2) = -3x + 3(-2) = -3POPO333-3-

x = +5 (for P)x = +5 (for P)x + 4(-2) = -3x + 4(-2) = -3POPO443-3-

x = +4 (for S)x = +4 (for S)x + 3(-2) = -2x + 3(-2) = -2SOSO332-2-

x = +6 (for S)x = +6 (for S)x + 4(-2) = -2x + 4(-2) = -2SOSO442-2-

x = +4 (for C)x = +4 (for C)+1 + +1 + x + 3(-2) = -1x + 3(-2) = -1HCOHCO33--

Calculation of Oxidation Numbers:Calculation of Oxidation Numbers:

Example 1

Al2S3

Al is a monatomic ion with a 3+ charge, so its oxidation state is +3

(Rule 2).When combined with metals in binary compounds, S is a monatomic ion with a 2- charge, so

its oxidation state is -2 (Rule 2).

2 Al at +3 each = +63 S at -2 each = -6

sum = 0

+3 -2

Reminder: Nonmetals (like sulfur) as well as metals not in group I or II can have many oxidation states, so they must be carefully analyzed.

Example 2

Na2CO3

Oxygen has a -2 oxidation state (Rule 1).

Na is a monatomic ion with a 1+ charge, so its oxidation state is +1 (Rule 4).

2 Na at +1 each = +23 O at -2 each = -6

sum = -4

The overall sum must be 0 for a compound (Rule

2), so carbon must have a +4 oxidation state.

+1 -2+4

●● Oxidation-reduction (redox) Oxidation-reduction (redox) reaction:reaction:

Redox reactionRedox reaction:: A reaction in which one or A reaction in which one or

more electrons are transferred from one more electrons are transferred from one

atom to another.atom to another.

Redox reactionRedox reaction: : AA reaction in which reaction in which

oxidation numbers of some elements are oxidation numbers of some elements are

changed in a reaction process.changed in a reaction process.

22FeClFeCl33 + SnCl + SnCl22 = 2FeCl = 2FeCl22 + SnCl + SnCl44

22FeFe3+3+ + Sn + Sn2+2+ = 2Fe = 2Fe2+2+ + Sn + Sn4+4+

Redox reaction

22Mg + OMg + O22 = 2MgO = 2MgO

Electron transferCu(s) + 2 Ag+(aq) Cu2+(aq) + 2 Ag(s)

Loss of Electrons =

OXIDATION (LEO)Gain of Electrons =

REDUCTION (GER)

- 2 e-

2 x +1 e-

Redox

reaction ReductionReduction reactionreaction

Oxidation reactionOxidation reaction

── lossloss of electrons of electrons

── oxidation number oxidation number increasesincreasesA g →A gA g →A g+ + + + ee--

── gaingain in electrons in electrons

── decreasedecrease in oxidation number in oxidation numberFeFe2+2+ + 2 + 2ee-- →Fe→Fe

Redox process always occurs together. In redox process, one can’t occur without the other.

Half- reaction

Half- reaction

Redox Reaction Example

• A redox reaction:

Cu(s) + 2Ag+(aq) Cu2+(aq) + 2Ag(s)

Ox. #: 0 +1 +2 0

oxidation reduction

Redox Reaction Example (cont.)

• We can envision breaking up the full redoxreaction into two 1/2 reactions:

Cu(s) + 2Ag+(aq) Cu2+(aq) + 2Ag(s)

oxidation reduction

Cu(s) Cu2+(aq) + 2e-

Ag+(aq) + e- Ag(s)

“half-reactions”

Oxidation

Reduction

Oxidizing agentOxidizing agent

●● Oxidizing agent and reducing Oxidizing agent and reducing agentagent

─ electron acceptor; species is reduced.

Reducing agentReducing agent

─ electron donor, species is oxidized.

)()(2

)(2

)( saqaqs CuFeCuFe

oxidizing agent

reducing agentLoss of 2 e-1 oxidation

Gain of 2 e-1 reduction

– An oxidizing agent is a species that oxidizes

another species; it is itself reduced.

– A reducing agent is a species that reduces

another species; it is itself oxidized.

Oxidizing agent:Oxidizing agent:

Reducing agent:Reducing agent: Losses one or more electronsLosses one or more electrons

Causes reductionCauses reduction

Undergoes oxidationUndergoes oxidationBecomes more positive or less Becomes more positive or less

negativenegative

Gains one or more electronsGains one or more electrons

Causes oxidationCauses oxidation

Undergoes reductionUndergoes reductionBecomes more negativeBecomes more negative

or less positiveor less positive

ExampleExample

CHCH4 4 ((gg) + 2 O) + 2 O22 ( (gg) ) CO CO22 ( (gg) + 2H) + 2H22O (O (ll))

CC HH OO-4-4 +1+1 00+4+4 +1+1 -2-2

Which species is oxidized ? (lost electrons/Ox state became more positive)

Which species is reduced ? (gained electrons/Ox state became more negative)

●● Redox coupleRedox couple ((Pair of Electrons Pair of Electrons Oxidation and Reduction)Oxidation and Reduction)

88HH++ + MnO + MnO44 + 5Fe + 5Fe2+2+ Mn Mn2+2+ + 5Fe + 5Fe3+3+ + 4H + 4H22OO

ReductionReduction:: 8H8H++ + MnO + MnO44 + 5e + 5e Mn Mn2+2+ + 4H + 4H22OO

OxidationOxidation:: 5Fe5Fe2+2+ 5Fe 5Fe3+3+ + 5e + 5e

The overall reaction is split into two The overall reaction is split into two half-half-

reactionsreactions, one involving , one involving oxidationoxidation and one and one

reductionreduction..

22FeFe3+3+ + Sn + Sn2+2+ = 2Fe = 2Fe2+2+ + Sn + Sn4+4+

ReductionReduction: : FeFe3+3+ + + ee-- →Fe →Fe2+2+

OxidationOxidation: Sn: Sn2+2+ - 2 - 2ee-- →Sn →Sn4+4+

Redox couplesRedox couples

Oxidized species Oxidized species // reduced species reduced species

SnSn4+4+ / Sn / Sn2+2+ ; Fe ; Fe3+ 3+ / Fe/ Fe2+2+

The oxidized and reduced states of each substance The oxidized and reduced states of each substance

taking part in a half-reaction form a taking part in a half-reaction form a redox coupleredox couple..

Notation of redox coupleNotation of redox couple

Ox + Ox + nn e e -- RedRed

Oxidized state + Oxidized state + nn e e -- Reduced stateReduced state

Redox couple half-reactions:Redox couple half-reactions:

6-2 Voltaic Cells (Primary Cells)

Electron transferZn(s) + CuSO4 (aq) ZnSO4 (aq) + Cu(s)

- 2 e-

+2 e-

Loss of Electrons =

OXIDATION (LEO)Gain of Electrons =

REDUCTION (GER)

What could happen when we

put a piece of zinc metal into

the solution of copper sulfate?

Zn metal

Cu2+ ions

• Copper is deposited on the zinc

The copper plates out onto the The copper plates out onto the

surface of the zinc metalsurface of the zinc metal

• The blue copper (Ⅱ)ions are gradually

replaced by colorless zinc ions

• Chemical energy (reaction enthalpy)

is released as heat

CHEMICAL CHANGE CHEMICAL CHANGE ELECTRIC CURRENT ELECTRIC CURRENT

Zn metal

Cu2+ ions

With time, Cu plates out onto Zn metal strip, and Zn strip “disappears.”

Zn is oxidizedZn is oxidized and is the reducing agent and is the reducing agent Zn(s) Zn(s) Zn Zn2+2+(aq) + 2e(aq) + 2e--

CuCu2+2+ is reduced is reduced and is the oxidizing agentand is the oxidizing agentCuCu2+2+(aq) + 2e(aq) + 2e-- Cu(s) Cu(s)

Oxidation: Zn(s) Oxidation: Zn(s) Zn Zn2+2+((a qa q) + 2e) + 2e--

Reduction: CuReduction: Cu2+2+((a qa q) + 2e) + 2e-- Cu(s) Cu(s)--------------------------------------------------------------------------------------------------------Zn(s) + CuZn(s) + Cu2+2+((aqaq) ) Zn Zn2+2+((aqaq) + Cu(s)) + Cu(s)

Zn metal

Cu2+ ions

Electrons are transferred from Zn to Cu2+, but there is no useful electric current.

CHEMICAL CHANGE CHEMICAL CHANGE ELECTRIC CURRENT (2) ELECTRIC CURRENT (2)

2 e-

ToTo obtain a useful obtain a useful

current, we separate current, we separate

the oxidizing and the oxidizing and

reducing agents so that reducing agents so that

electron transfer occurs electron transfer occurs

thru an external wire. thru an external wire.

Zn

Zn2+ ions

Cu

Cu2+ ions

wire

saltbridge

electrons

CHEMICAL CHANGE CHEMICAL CHANGE ELECTRIC CURRENT (2) ELECTRIC CURRENT (2)

• This is accomplished in a VOLTAIC cell.

(also called GALVANIC cell)

• A group of such cells is called a battery.

Voltaic CellVoltaic Cell

A device in which A device in which

chemical energy ischemical energy is

changed to electricalchanged to electrical

energy.energy.

6-2.1 6-2.1 Cu- Zn Primary CellCu- Zn Primary Cell

(-) (+)

Zn

Zn2+ ions

Cu

Cu2+ ions

wire

saltbridge

electrons

• • Electrons travel thru external wire.Electrons travel thru external wire. Salt bridge Salt bridge allows anions and cations to allows anions and cations to move between electrode compartments.move between electrode compartments. This maintains electrical neutrality.This maintains electrical neutrality.

ANODE (-) CATHODE (+)

Negative electrode generates electron

Oxidation Occur

Positive electrode accepts electron

Reduction Occur

Zn half-cell Cu half-cellZn2+ / Zn Cu2+ / Cu

Voltaic CellsVoltaic Cells

A A voltaic cellvoltaic cell consists of two half-cells. consists of two half-cells.

• Each half-cell is a portion of the

electrochemical cell in which a half-reaction

takes place.• A simple half-cell can be made from a metal

strip dipped into a solution of its metal ion.

• For example, the zinc-zinc ion half cell consists of a zinc strip dipped into a solution of a zinc salt.

• Another simple half-cell consists of a

copper strip dipped into a solution of a

copper salt.

• In a voltaic cell, two half-cells are connected

in such a way that electrons flow from one

metal electrode to the other through an

external circuit.

As long as there is an external circuit, As long as there is an external circuit, electrons can flowelectrons can flow through it from one through it from one electrode to the other.electrode to the other.

• Because zinc has a greater tendency to lose electrons than copper, zinc atoms in the zinc electrode lose electrons to form zinc ions.

• The electrons flow through the external circuit

to the copper electrode where copper ions gain

the electrons to become copper metal.

The two half-cells The two half-cells must also be connected must also be connected

internallyinternally to allow ions to flow between them. to allow ions to flow between them.

• Without this internal connection, too much positive charge builds up in the zinc half-cell (and too much negative charge in the copper half-cell) causing the reaction to stop.

• Figure A and B show the two half-cells of a voltaic cell connected by salt bridge.

• A A salt bridgesalt bridge is a U shape tube of an is a U shape tube of an electrolyte in a gel that is connected to electrolyte in a gel that is connected to the two half-cells of a voltaic cell.the two half-cells of a voltaic cell.

• The salt bridge allows the flow of ions but

prevents the mixing of the different

solutions that would allow direct reaction

of the cell reactants.

6-2.2 6-2.2 Cell ReactionCell Reaction

The two half-cell reactions, as noted earlier, are:The two half-cell reactions, as noted earlier, are:

e2)aq(Zn)s(Zn 2

)s(Cue2)aq(Cu2

oxidation half-reaction

reduction half-reaction

Note that the sum of the two half-reactionsNote that the sum of the two half-reactions

• Note that electrons are given up at the

anode and thus flow from it to the

cathode where reduction occurs.

)s(Cu)aq(Zn)aq(Cu)s(Zn 22

is the net reaction that occurs in the voltaic

cell; it is called the cell reaction.

6-2.3 6-2.3 Notation for Voltaic CellsNotation for Voltaic Cells

It is convenient to have a It is convenient to have a shorthandshorthand way of way of

designating particular voltaic cells.designating particular voltaic cells.

• The anode (oxidation half-cell) is written on the left. The cathode (reduction half-cell) is written on the right.

• The cell consisting of the zinc-zinc ion half-cell and the copper-copper ion half-cell, is written

)s(Cu|)aq(Cu||)aq(Zn|)s(Zn 22

anode cathode

（－）（＋）

• A boundary between different phases

(e.g., an electrode and a solution) is

represented by a single vertical line (│)

• The boundary between half-cell

compartments, usually a salt, is

represented by a double vertical line

( )

• The cathode in a voltaic cell has a positive sign

• The anode in a voltaic cell has a negative sign because electrons flow from it.

Notation for Voltaic CellsNotation for Voltaic Cells

• The two electrodes are connected by a

salt bridge, denoted by two vertical bars.



anode cathodesalt bridge


• The cell terminals are at the extreme

ends in the cell notation.





• A single vertical bar indicates a phase

boundary, such as between a solid

terminal and the electrode solution.




2Fe3+(c1) + Sn2+(c2) = 2Fe2+(c3) + Sn4+(c4)

Half reactionFe3+ + e Fe2+

Sn2+ - 2e Sn4+

Sn4+ / Sn2+

Fe3+ / Fe2+

Pt│Sn4+ (c4), Sn2+(c2)

Pt│Fe3+ (c1), Fe2+(c3)

Cell notation

Pt│Sn4+ (c4), Sn2+(c2) Fe3+ (c1), Fe2+(c3) │Pt

When the When the half-reaction involves a gashalf-reaction involves a gas, , an inert material such asan inert material such as platinum platinum serves serves as a terminal and an electrode surface on as a terminal and an electrode surface on which the reaction occurs.which the reaction occurs.

• hydrogen electrode

• The cathode half-reaction is

)g(He2)aq(H2 2

• The notation for the hydrogen electrode, written as a cathode, is

Pt|)g(H|)aq(H 2

• To write such an electrode as an anode, you simply reverse the notation.

)aq(H|)g(H|Pt 2

To fully specify a voltaic cell, it is To fully specify a voltaic cell, it is necessary to give the necessary to give the concentrations concentrations of solutionsof solutions and the and the pressure of gasespressure of gases..

• In the cell notation, these are written in parentheses ( ). For example,

Pt|)atm 0.1(H|)aq(H||)M 0.1(Zn|)s(Zn 22

Line NotationLine Notation solidsolidAqueousAqueousAqueousAqueoussolidsolid

Anode on the leftAnode on the leftCathode on the rightCathode on the right

Single line different phases.Single line different phases.

Double line salt bridge.Double line salt bridge.

If all the substances on one side are aqueous, If all the substances on one side are aqueous, a platinum electrode is indicated. a platinum electrode is indicated.

Cu(s)Cu(s)CuCu+2+2(aq)(aq)FeFe+2+2(aq),Fe(aq),Fe+3+3(aq)(aq)Pt(s)Pt(s)

6-2.46-2.4 Electromotive ForceElectromotive Force• The maximumThe maximum potential potential difference difference

between the electrodes of a voltaic cell is between the electrodes of a voltaic cell is

referred to as the referred to as the eelectrolectrommotive otive fforceorce

((emfemf) of the cell, denoted ) of the cell, denoted EE

• EE = = φφcathodecathode – – φφanodeanode = = φφ++ - -φφ--

• EE is a is a positive positive number.number.

The The standard emf, standard emf, EE oo, is the emf of a , is the emf of a

cell operating under standard conditionscell operating under standard conditions

of concentration (1 M), pressure (1atm),of concentration (1 M), pressure (1atm),

and temperature (25 and temperature (25 ooC).C).

Standard Notation for Electrochemical CellsStandard Notation for Electrochemical Cells

ANODE Zn / Zn2+ // Cu2+ / Cu CATHODE

OXIDATION

Anode electrode

Active electrolyte in oxidation half-reaction

Cathode electrode

Active electrolyte in reduction half-reaction

Salt bridge Phase boundaryPhase boundary

REDUCTION

Anode and CathodeAnode and Cathode

OOXIDATIONXIDATION occurs at the occurs at the AANODENODE..

RREDUCTIONEDUCTION occurs at the occurs at the CCATHODE.ATHODE.

Mnemonic: Mnemonic: OO and and AA are vowels; are vowels; RR and and CC are consonants are consonants

6-2.5 6-2.5 Types of ElectrodesTypes of Electrodes

((a) metal/metal ion a) metal/metal ion electrodeelectrode

(b) metal/ insoluble (b) metal/ insoluble salt electrodesalt electrode

(c) gas electrode(c) gas electrode

(d) redox electrode(d) redox electrode

Types of Electrode (continued)Types of Electrode (continued)EElleeccttrrooddee TTyyppee DDeessiiggnnaattiioonn

MMeettaall//mmeettaall iioonn MM((ss))||MM++((aaqq))

GGaass PPtt((ss))||XX22((gg))||XX++((aaqq)) oorrPPtt((ss))||XX22((gg))||XX--((aaqq))

MMeettaall//iinnssoolluubbllee ssaalltt MM((ss))||MMXX((ss))||XX--((aaqq))

RReeddooxx PPtt((ss))||MM++((aaqq)),,MM22++((aaqq))

Types of Electrode (continued)Types of Electrode (continued)RReeddooxx CCoouuppllee HHaallff rreeaaccttiioonn

MM++//MM MM++((aaqq)) ++ ee-- MM((ss))

XX++//XX22

XX22//XX--XX++((aaqq)) ++ ee-- ½½ XX22((gg))½½ XX22((gg)) ++ ee-- XX--((aaqq))

MMXX//MM//XX-- MMXX((ss)) ++ ee-- MM((ss)) ++ XX--((aaqq))

MM22++//MM++ MM22++((aaqq)) ++ ee-- MM++ ((ss))

l

l

l

l

l

• The types of electrodeThe types of electrodeMetal-metal ion electrodeMetal-metal ion electrode

Zn(s) Zn∣Zn(s) Zn∣ 2+2+( aq );( aq ); Cu(s ) Cu(s ) ∣∣CuCu2+2+ ( aq ) ( aq )

MM∣∣MMn+ n+ :: MMn+n+ + ne == M + ne == M

Gas electrodeGas electrode

Pt H∣Pt H∣ 22(p) H∣(p) H∣ ++((cc)) 2H2H++ + 2e == H + 2e == H22

Metal-insoluble salt electrodesMetal-insoluble salt electrodes

Pt Hg(∣Pt Hg(∣ ll) Hg∣) Hg∣ 22ClCl22(s) Cl∣(s) Cl∣ -- ((cc))

HgHg22ClCl2 2 (s)+ 2e == 2Hg + 2Cl(s)+ 2e == 2Hg + 2Cl--

Oxidation-reduction electrodesOxidation-reduction electrodes

FeFe3+3+ + e == Fe + e == Fe2+ 2+ Pt Fe∣Pt Fe∣ 3+3+((cc11),Fe),Fe2+2+((cc22) )

Documents

Chapter six Electrochemistry. Oxidation NumbersOxidation Numbers Oxidation-reduction reactionOxidation-reduction reaction Oxidizing agent and reducing