Electrochemistry

Oxydation-Reduction Reactions

Balancing Oxydation-Reduction Reactions Voltaic Cells

Cell EMF

Spontaneity of Redox Reactions Spontaneity of Redox Reactions

Effect of Concentration of Cell EMF Batteries

Corrosion Electrolysis

Oxidation Oxidation Oxidation

Oxidation----reductionreductionreductionreduction reactionsreactionsreactionsreactions → the transfer of electrons from one species to another.

The oxidation state of one or more substances in the reaction changes.

substances in the reaction changes.

The transfer of electrons → can be used to produce energy in the form of electricity.

Electrochemistry Electrochemistry Electrochemistry

Electrochemistry is the study of the relationships between chemical reactions and electrical energy.

We identify a reaction as oxidation-reduction by comparing the oxidation numbers of atoms in the reactants and products.

If the oxidation numbers change, the If the oxidation numbers change, the reaction is an oxidation-reduction.

zinc metal loses electrons to become a cation, and thus zinc has been

oxidized

Hydrogen gains electrons to become hydrogen gas.

Hydrogen has been

reduced

reduced

Zinc is the reducingreducing agentreducingreducing agentagentagent, or reductant,reductant,reductant,reductant,

and hydrogen ion is the oxidizingoxidizing agentoxidizingoxidizing agentagentagent, or

oxidant oxidant oxidant oxidant.

The amount of each element must be the same on both sides of the equation.

Balancing the number of electrons

Balancing the mass automatically balances the total charge on each side.

In many reactions, though, balancing the mass does not result in balancing charge.

In the next equation the law of conservation of mass appears to have been obeyed.

However, note that the total charge on the left side is +3, while the total charge on the right side is only +2.

The two processes, oxidation of manganese metal and reduction of chromium(III) ion, do not correspond to the transfer of the same number of electrons.

If we multiply the manganese on each side by If we multiply the manganese on each side by 3 and the chromium on each side by 2, we get

It is not always convenient to balance oxidation-reduction reactions by inspection.

The equation

would be impossible to balance by inspection.

To balance such an equation, we use a technique known as the method of halfhalfhalfhalf----reactionsreactionsreactionsreactions.

A half-reaction is an equation that shows either oxidation or reduction alone.

The electrical energy produced by a spontaneous oxidation-reduction reaction can be "harnessed" using a voltaic cell (also called a galvanic cell).

In a voltaic cell the two half-reactions are made to occur in separate compartments (half-cells).

to occur in separate compartments (half-cells).

The electrons can be transferred only through the wire.

◦ oxidation and reduction half-reactions occur at separate electrodes

Anode - the electrode at which oxidation takes place.

◦ the negative (-) electrode

◦ produces electrons

Cathode - the electrode at which reduction Cathode - the electrode at which reduction takes place.

◦ the positive (+) electrode

Oxidation-reduction reactions part I Oxidation-reduction reactions part I

Galvanic cells I: the copper-zinc cell Galvanic cells I: the copper-zinc cell

Galvanic cells II: the zinc-hydrogen cell Galvanic cells II: the zinc-hydrogen cell

EXAMPLE:

EXAMPLE:

EXAMPLE:

EXAMPLE:

Describe how you would construct a galvanic cell based on the following reaction:

Pb2+ ₍

_aq

_s

_s

_aq

_aq

_s

_s

_aq

• Pb2+ (aq) + 2 e- → Pb (s) • Zn (s) → Zn2+ (aq) + 2 e

-• Looking at the two half-reactions, we find that

the Pb2+ _{is being reduced, and the Zn is being} oxidized.

Therefore, the anode compartment of our cell

• Therefore, the anode compartment of our cell

would consist of a strip of zinc metal immersed in a solution containing Zn2+ _{ions (such as zinc} nitrate).

• The cathode compartment would consist of a

strip of lead immersed in a solution containing Pb2+ _{ions (such as lead (II) nitrate).}

• Pb2+ (

aq

s

s

aq

-• The two half-cells would be connected to

each other with a salt bridge and an external wire.

wire.

• Electrons flow through the wire from the zinc

anode to the lead cathode.

• Anions move from the cathode compartment

towards the anode while cations migrate from the anode compartment toward the cathode.

Redox chemistry of iron and copper Redox chemistry of iron and copper

Single vertical line, , represents a phase boundary.

Double vertical line, ǁ, represents a salt bridge.

• Shorthand for the anode half-cell is always

written on the left of the salt-bridge symbol, followed on the right of the symbol by the shorthand for the cathode half-cell.

– Reactants in each half cell are written first, followed

– Reactants in each half cell are written first, followed by products.

– Electrons move through the external circuit from left to right.

– For Zn (s) + Cu2+ (aq) → Zn2+ (aq) + Cu (s): Zn (s)Zn2+ (aq_{) ǁ Cu2+ (}aq)Cu (s).

Cell involving a gas.

◦ Additional vertical line due to presence of additional phase.

◦ List the gas immediately adjacent to the appropriate electrode.

Detailed notation includes ion concentrations Detailed notation includes ion concentrations and gas pressures.

EXAMPLE:

EXAMPLE:

EXAMPLE:

EXAMPLE:

Give the shorthand notation for a galvanic cell that employs the overall reaction

Pb(NO₃)₂(

aq

s

s

aq

aq

s

s

aq

SOLUTION: The two half-reactions for this overall reaction are:

Pb2+ ₍

_aq

_s

s

_aq

e-From these half-reactions, we know that lead is being reduced and nickel is being oxidized.

• Therefore, Ni is the anode and Pb is the

cathode. The cell notation is:

• Ni(

s

aq

aq

s

• This cell would consist of a strip of nickel as

the anode dipping into an aqueous solution the anode dipping into an aqueous solution of Ni(NO₃)₂ and a strip of Pb as the cathode dipping into an aqueous solution of Pb(NO₃)₂.

• The two half-cells would be connected by a

• Electrons spontaneously flow from one species to another—and through the wire from the anode to the cathode of a voltaic cell—because of a difference in potential energy, or a

potential difference

energy, or a

potential difference

• The potential difference between two

electrodes is measured in volts.

The potential difference that drives electrons

through the wire in a voltaic cell is called the

electromotive force or

electromotive force or

emf.

For a voltaic cell the emf is denoted E_cell and referred to as the cell potential.

Cell potential - measured with a voltmeter.

◦ Gives a positive reading when the + and - terminals of the voltmeter are connected to cathode (+) and anode (-), respectively.

can use voltmeter-cell connections to determine which electrode is the anode and which is the cathode

The value of a cell potential depends on what half-reactions are taking place in the two compartments of the cell.

The cell potential measured under standard conditions,

E

M

E

M

potential potential

For the zinc and copper voltaic cell in Figure 20.5,

E

That is for the reaction

at 25°C, where the concentrations of copper and zinc ions are both 1

M

Such potentials can be measured experimentally, but many of them can be calculated from tabulated standardstandardstandardstandard reductionreductionreductionreduction potentials

potentials potentials

The more positive the standard cell potential, the greater the driving force for electrons to flow from the anode to the cathode.

Because the cathode of a voltaic cell is always the half-reaction with the more positive (or the half-reaction with the more positive (or less negative) standard reduction potential, the standard cell potential of a voltaic cell is always positive.

The standard reduction potentials for the various half-reactions are measured against a standardstandardstandardstandard hydrogen

hydrogen hydrogen

hydrogen electrodeelectrodeelectrodeelectrode (SHE).

The half-reaction of interest and the SHE, both under standard conditions, are made into a under standard conditions, are made into a voltaic cell, as shown in Figure 20.11, and the cell potential is measured experimentally.

The standard potential of the standard hydrogen electrode's half-reaction is arbitrarily assigned a

value of zero, so the measured potential

The standard reduction potentials in Table 20.1 can be used to compare the oxidizing power or reducing power of a substance.

The more positive the value of

E°

species, the more readily it undergoes

species, the more readily it undergoes

reduction and the better

oxidizing agent

As

E°

on the right side of the arrow becomes a stronger

reducing agent

Standard reduction potentials Standard reduction potentials

It is possible to use standard reduction potentials to predict whether or not a given

oxidation-reduction reaction will be

spontaneous.

For instance, will copper solid be oxidized by For instance, will copper solid be oxidized by a solution containing iron(II) ions?

To determine the answer, we calculate the standard cell potential for the cell as we have described it.

In our description, copper is being oxidized so the copper electrode is the anode.

so the copper electrode is the anode.

Iron is being reduced, making the iron electrode the cathode.

E_°red = 0.34 V

E_°red = - 0.44 V

The negative E_°cell tells us that this reaction, as written, will not occur spontaneously. (Its reverse reaction will be spontaneous.)

Our ability to predict the spontaneity of a chemical reaction by calculating a standard cell potential points to a relationship between

the sign of

E°

G°

standard change in Gibbs free energy. standard change in Gibbs free energy.

Quantitatively, this relationship is expressed as

n

F

Both

n

F

Therefore, a positive value for

E°

correspond to a negative value for

G°

denoting a spontaneous reaction.

The relationship between cell potential and The relationship between cell potential and free energy change holds at conditions other than standard, as well.

E = Eo − 0.0592

Nernst equation: (in volts at 25o_C).

◦ Enables us to calculate cell potentials under nonstandard-state conditions.

E = Eo − 0.0592

n log Q

Calculate

E

s

_aq

_aq

_s

2 Cr (

s

_E

( ) E_cello = +0 74. + −0 13. = +0 61. V E = Eo − 0.0592 n log Q E = E − n log Q E_cell = E_cello − 0.0592 6 log Cr3+

[ ]

[ ]

Three different ways to determine the value of an equilibrium constant

K

Equilibrium constants for redox reactions tend to be either very large or very small in comparison with equilibrium constants for acid-base reactions.

◦ Positive value of Eo _{corresponds to K > 1.}

◦ Positive value of E corresponds to K > 1.

Important application of Nernst equation -electrochemical determination of pH using a pH meter.

Consider a cell with a hydrogen electrode as the anode and a second reference electrode the anode and a second reference electrode as the cathode.

◦ Pt (s)H₂ (1 atm)_{H+ (? M)ǁreference cathode.}

can measure the pH of a solution by

E_cell = 0.0592pH+ E_ref

pH= Ecell − Eref 0.0592

can measure the pH of a solution by measuring

E

Actual pH measurements use a glass electrode with a calomel electrode as the reference.

EXAMPLE

EXAMPLE

EXAMPLE

EXAMPLE::::

The following cell has a potential of 0.49 V. Calculate the pH of the solution in the anode compartment. Pt(s) H₂(g) (1 atm)H+_{(pH = ?)Cl}-_{(aq) (1M)Hg} 2Cl2 Pt(s) H₂(g) (1 atm)H+_{(pH = ?)Cl}-_{(aq) (1M)Hg} 2Cl2 (s) Hg (l)

The cell reaction is

Hg₂Cl₂ (s) + H₂ (g) → 2 Hg (l) + 2 Cl- _{(aq) + 2}

The cell reaction is Hg₂Cl₂ (s) + H₂ (g) → 2 Hg (l) + 2 Cl- _{(aq) + 2} H+ _(aq) V 28 . 0 V 28 . 0 V 00 . 0 o o o = + = + = E E E o 0.00 V 0.28 V 0.28 V Cl Hg, Cl Hg o H H o -2 2 + 2 + = + = = E _→ E _→ E pH= Ecell − Eref 0.0592 55 . 3 0592 . 0 V 28 . 0 V 49 . 0 pH = − =

EXAMPLE

EXAMPLE

EXAMPLE

EXAMPLE::::

Calculate the equilibrium constant for the following reaction at 25o_C.

5 S O 2- _{(aq) + I (s) + 6 H O (l)} → _{10 SO} 2- _{(aq) + 2 IO}3- _{(aq) +}

5 S₂O₈2- _{(aq) + I}

2 (s) + 6 H2O (l) → 10 SO42- (aq) + 2 IO3- (aq) +

S₂O₈2- _{(aq) + 2 e}- → _{2 SO} 42- (aq) Eo = +2.01 V I₂(s) + 6H₂O (l) → 2IO3- _{(aq) + 12 H}+_{(aq) +10e}- _Eo = -1.20 V Ecello = 2.01+ −( 1.20)= +0.81 V

The value of

n

log K = (10)(0.81) (0.0592) =137 K = 10137 Eo = 0.0592 n log K

• Most important practical application of

galvanic cells is their use as batteries.

• Features required in a battery depend on the

application.

• C. General features. • C. General features.

– Compact and lightweight.

– Physically rugged and inexpensive.

– Provide a stable source of power for relatively long

• A battery is a self-contained source of electrochemical energy made from one or more voltaic cells.

• Ordinary flashlight batteries consist of a

single voltaic cell, while car batteries are six single voltaic cell, while car batteries are six identical voltaic cells connected in series.

• It is worth noting that electrochemistry is one of only a very few commercially viable methods of generating electricity.

Car Battery

Cathode:

Car Battery

Cathode:

Corrosion - the oxidative deterioration of a metal.

Well-known example of corrosion -conversion of iron to rust.

◦ Requires both oxygen and water.

◦ Requires both oxygen and water.

◦ Involves pitting of the metal surface.

rust is deposited at a location physically separated from the pits

• Corrosion is the undesirable oxidation of a metal.

• A familiar example of corrosion is the rusting of

iron.

• Iron metal is oxidized to Fe2+ by oxygen.

• The Fe2+ is then further oxidized to Fe3+ in a

Corrosion of iron can be prevented by coating it with paint or with another metal such as tin or zinc.

Coating with paint or with a less readily oxidized metal such as tin protects the iron oxidized metal such as tin protects the iron simply by preventing oxygen and water from reaching the iron surface.

If such a coating is damaged and the iron is exposed, corrosion of iron will occur in the exposed area.

• Some metals create their own sealant by oxidation.

• Aluminum, for example, oxidizes fairly

readily.

• Coating iron with a

more

• But unlike tin, zinc protects the iron even if the zinc coating is damaged.

the zinc coating is damaged.

• It does so by making the iron serve as the

• When the zinc coating is damaged and iron is exposed, the zinc itself is oxidized rather than the iron.

E_°red = - 0.44 V

• Iron, with the more positive (less negative)

E°

less easily oxidized.

• Zinc will be oxidized, serving as the

sacrificial

anode

anode

• This type of corrosion prevention is called

cathodic cathodic cathodic

cathodic protectionprotectionprotectionprotection.

E_°red = - 0.44 V

• One of the many uses of cathodic protection is to

prevent corrosion of underground pipes and

storage tanks—many of which are iron.

• A more easily oxidized metal is placed in electrical

contact with the object to be protected by an contact with the object to be protected by an insulated copper wire.

• The iron object becomes the cathode (and the more

easily oxidized metal the sacrificial anode) in a voltaic cell.

• Oxidation eventually consumes the sacrificial

• A voltaic cell is one in which a spontaneous chemical reaction is used to generate a voltage.

• Electrolysis is the use of a voltage to drive a nonspontaneous reaction.

nonspontaneous reaction.

• Reactions that are driven by an externally

supplied voltage are called electrolysiselectrolysiselectrolysiselectrolysis reactions

reactions reactions

reactions, and electrochemical cells designed for the purpose of carrying out electrolysis reactions are called electrolyticelectrolyticelectrolyticelectrolytic cellscellscellscells

Sodium metal is produced commercially by electrolysis of molten sodium chloride.

Electrodes are immersed in molten sodium

chloride, and a voltage source drives

electrons from the anode to the cathode. electrons from the anode to the cathode.

Sodium is reduced at the cathode to molten sodium metal.

Chloride ions are oxidized to chlorine gas at the anode.

Cathode: Anode: E_°red = - 2.71 V E_°red = 1.36 V E° cell = - 4.07 V

A negative cell potential means that the reaction is nonspontaneous.

The value of the cell potential tells us that a minimum of 4.07 volts must be applied to the cell to drive the reaction in the desired direction.

E_°red = - 2.71 V Cathode:

Anode: E°red = 1.36 V

E°

cell = - 4.07 V

Electrolysis of molten salts is used commercially to produce a number of active metals.

E_°red = - 2.71 V Cathode:

Anode: E°red = 1.36 V

E°

cell = - 4.07 V

Electrolysis of molten salts is used commercially to produce a number of active metals.

• Electrolysis can also be done with aqueous solutions.

• Although chlorine gas can be produced from

a solution of sodium chloride, sodium metal cannot be produced.

cannot be produced.

• This has to do with the relative ease of

reduction of water versus sodium ions.

E_°red = - 2.71 V

• With a far less negative reduction potential, water—rather than sodium— will be reduced

at the cathode of an electrolytic cell

containing aqueous sodium chloride.

E_°red = - 2.71 V

E_°red = - 2.71 V

• In a comparison of reduction potentials, we might also predict that water, and not chloride ions, would be oxidized at the anode of such an electrolytic cell.

E_°red = 1.36 V

• The first half-reaction, with the more positive reduction potential, is more apt to occur as a reduction rather than an oxidation.

• Experimentally, however, chlorine gas is

produced at the anode during electrolysis of produced at the anode during electrolysis of aqueous sodium chloride.

E_°red = 1.36 V

• This is explained by kinetics.

• Although the oxidation of water is

thermodynamically favored, the oxidation of chloride ions is kinetically favored because of a lower activation energy.

a lower activation energy.

E_°red = 1.36 V

• Electroplating is an electrolytic process used to deposit a thin layer of one metal on another.

• In such a process the metal to be deposited is used as the anode, and the metal on which used as the anode, and the metal on which the deposit is to be made is the cathode.

• Electroplating is the process by which silver-plated flatware is finished.

• It is also the process used to coat iron car

• When the electrode in an electrolytic cell is involved in the reaction, it is called an

active

electrode.

• When the same metal is to be oxidized from

the anode and reduced (deposited) at the the anode and reduced (deposited) at the cathode, the standard cell potential is zero.

• It therefore requires only a very small voltage to drive such a process.

• Using stoichiometry, we can relate the

amount of metal deposited in an

electroplating process to the current and the length of time for which the current is applied.

• The charge passing through an electrolytic cell is measured in

coulombs

• How many grams of copper would be reduced by the application of 5.00 amps to a solution of copper sulfate for 35.0 minutes?

EXAMPLE

EXAMPLE

EXAMPLE

EXAMPLE::::

How many grams of Cl₂ would be produced in the electrolysis of molten NaCl by a current of 4.25 A for 35.0 min?

• Remember that a coulomb is an A.s or that an ampere is C/s. • 2 Cl- → Cl₂ + 2 e -• moles of electrons = 2 4.25 C s ×35.0 min × 60 s 1 min × 1 mol e -96,500 C × 1 mol Cl₂ 2 mol e- × 70.9 g Cl₂ 1 mol Cl- = 3.28 g Cl2