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Chemistry: The Central Science

Chapter 20: Electrochemistry

Redox reaction power batteries

Electrochemistry is the study of the relationships between electricity and chemical

reactions

o It includes the study of both spontaneous and nonspontaneous processes

20.1: Oxidation States and Oxidation-Reduction Reactions

Oxidation numbers of all the elements involved in the reaction can be tracked to

determine whether the reaction is a redox reaction

In some reactions, the oxidation numbers change, but we cannot say that any

substance literally gains or loses electron

o E.g. Combustion of hydrogen to form water

In this reaction, hydrogen is oxidized from 0 to +1 oxidation state and

oxygen is reduced from the 0 to the -2 oxidation state

Water is not an ionic substance, however, and so there is not a

complete transfer of electrons from hydrogen to oxygen as water is

formed

o Using oxidation states is a convenient form of “bookkeeping,” but you should

not generally equate the oxidation state of an atom with its actual charge in a

chemical compound

The substance the oxidizes the other substance (thus becoming reduced) is called

the oxidizing agent or oxidant

The substance that reduces the other substance (thus becoming oxidized) is called

the reducing agent or reductant

20.2: Balancing Oxidation-Reduction Equations

When balancing a redox reaction, the gains and the losses of electrons must be

balanced

Half-Reactions

o Although oxidation and reduction must take place simultaneously, it is often

convenient to consider them as separate processes

o E.g.

Oxidation:

Reduction:

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Equations that show either oxidation or reduction alone are called half-

reactions

In the overall reaction, the number of electrons lost in the

oxidation half-reaction must equal to the number of electrons

gained in the reduction half-reaction

Balancing Equations by the Method of Half-Reactions

o The use of half-reactions to balance oxidation-reduction equations usually

begin with a “skeleton” ionic equation that shows only the substances

undergoing oxidation and reduction

o For balancing a redox reaction that occurs in acidic aqueous solution, the

procedure is as follows:

Divide the equation into two half-reactions, one for oxidation and the

other for reduction

Balance each half-reaction

First, balance the elements other than H and O

Next, balance the O atoms by adding H2O as needed

Then, Balance the H atoms by adding H+ as needed

Finally, balance the charge by adding e- as needed

At this point, you can check whether the number of electrons in

each half-reaction equals corresponds to the changes in

oxidation state

Multiply the half-reactions by integers, if necessary, so that the number

of electrons lost in one half-reaction equals the number of electrons

gained in the other

Add the two half-reactions and, if possible, simplify by canceling

species appearing on both sides of the combined equation

Check to make sure that atoms and charges are balanced

Balancing Equations for Reactions Occurring in Basic Solution

o One way to balance these reactions is to balance the half-reactions initially as

if they occurred in acidic solution

Then, count the H+ in each half-reaction, and add the same number of

OH- to each side of the half-reaction

The OH- will neutralize the protons on the side containing H+ and the

other side ends up with OH-

20.3: Voltaic Cells

The energy released in a spontaneous redox reaction can be used to perform

electrical work

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o This task is accomplished through a voltaic (or galvanic) cell, a device in which

the transfer of electrons takes place through an external pathway

E.g. spontaneous reaction occurs when a strip of zinc is placed in

contact with a solution containing Cu2+

Zn metal is in contact with Zn2+(aq) in one compartment of the

cell, and Cu metal is in contact with Cu2+(aq) in another

compartment

o Consequently, the reduction of the Cu2+ can only occur

by a flow of electrons through an external circuit

Two solid metals that are connected by the external circuit are

called electrodes

o Electrode at which oxidation occurs is called the anode

o Electrode at which reduction occurs is called the cathode

Each compartments of a voltaic cell is called a half-cell

o Anode:

o Cathode:

For a voltaic cell to work, the solutions in the two half-cells must

remain electrically neutral

As Zn is oxidized in the anode compartment, Zn2+ enter the

solution

As Cu2+ at the cathode reduces, the positive charge from the

solution is removed

A salt bridge serves this purpose

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o A salt bridge consists of a U-shaped tube that contains

an electrolyte, such as NaNO3(aq), whose ions will not

react with other ions in the cell or with the electrode

materials

Anions always migrate toward the anode and the cations

toward the cathode

In any voltaic cell the electron flow from the anode through the

external circuit to the cathode

A Molecular View of Electrode Processes

o Redox reaction between Zn(s) and Cu2+(aq) lead to an increase in Zn2+(aq) and

Cu, and a decrease in Zn(s) and Cu2+(aq)

o In the case of the voltaic cell, the Zn atom “loses” two electrons and becomes

a Zn2+(aq) in its compartment

The electron travels through the wire and attached to Cu2+(aq),

forming Cu(s) in its compartment

o The redox reaction between Zn and Cu2+ is spontaneous regardless of

whether they react directly or in the separate compartments of a voltaic cell

20.4: Cell EMF under Standard Conditions

The electrons flow spontaneously toward the electrode with the more positive

electrical potential

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The difference in potential energy per electrical charge (the potential difference)

between two electrodes is measured in units of volts

o

C is coulomb

V is volt

One electron has a charge of 1.60 × 10-19 C

Electromotive force (emf) – The potential difference between the two electrodes of

a voltaic cell providing the driving force that pushes the electron through the circuit

o The emf of a cell, denoted Ecell, is also called the cell potential

Ecell is measured in volts so it’s often referred to as cell voltage

o For any cell reaction that proceeds spontaneously such as that in a voltaic cell,

the cell potential will be positive

Under standard conditions (25°C and 1 M for aqueous or 1 atm for gases), the emf is

called the standard emf, or the standard cell potential, and is denoted E°cell

Standard Reduction (Half-Cell) Potentials

o Standard reduction potentials (E°red) – the standard electrode potentials

tabulated for reduction reactions

o

For all spontaneous reactions at standard conditions, E°cell > 0

o The reference half-reaction is the reduction of H+(aq) to H2(g) under standard

conditions, which is assigned a standard reduction potential of exactly 0 V

An electrode designed to produced this half-reaction is called a

standard hydrogen electron (SHE), or the normal hydrogen electrode

(NHE)

An SHE consists of a platinum wire connected to a piece of

platinum foil covered with finely divided platinum that serves as

an inert surface for the reaction

The electron is encased in a glass tube so that the hydrogen gas

under standard conditions (1 atm) can bubble over the platinum

The solution contains H+(aq) under standard (1 M) conditions

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o Whenever we assign an electrical potential to a half-reaction, we write the

reaction as a reduction

o Changing the stoichiometric coefficient in a half-reaction does not affect the

value of the standard reduction potential

o The more positive the value of E°red, the greater the driving force for reduction

under standard conditions

Strength of Oxidizing and Reducing Agents

o The more positive the E°red value for a half-reaction, the greater the tendency

for the reactant of the half-reaction to be reduced and oxidize another species

o The half-reaction with the smallest reduction potential is most easily reversed

as an oxidation

o Solutions of reducing agents are difficult to store for extended periods

because of the ubiquitous presence of O2, a good oxidizing agent

20.5: Free Energy and Redox Reactions

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E: A positive value of E indicates a spontaneous process; a negative value of E

indicates a nonspontaneous one

The activity series consists of the oxidation reactions of the metals, ordered from the

strongest reducing agent at the top to the weakest reducing agent at the bottom

o E.g.

Ni is oxidized and Ag+ is reduced

Positive value of E° indicates that the displacement of silver by

nickel is a spontaneous process

EMF and ΔG

o

ΔG is the change in Gibbs free energy

n is a positive number without units that represents the number of

electrons transferred in the reaction

F is called Faraday’s constant, which is the quantity of electrical charge

on one mole of electrons

o A positive value of E and a negative value of ΔG both indicate that a reaction is

spontaneous

o When the reactants and products are all in their standard states

20.6 Cell EMF under Nonstandard Conditions

As a voltaic cell is discharged, the reactants of the reaction are consumed, and the

products are generated, so the concentrations of these substances change

o The emf progressively drops until E = 0, at which point we say the cell is

“dead”

The Nernst Equation

o

This is the Nernst equation

At T = 298 K the quantity 2.303 RT/F equals 0.0592, with the

units of volts (V)

o

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o At E = 0, ΔG = 0

The system is at equilibrium

o In general, increasing the concentration of reactants of decreasing the

concentration of products increases the driving force for the reaction,

resulting in a higher emf and vice versa

Concentration Cells

o Cell emf depends on the concentration so a voltaic cell can be constructed

using the same species in both the anode and cathode compartments as long

as the concentration are different

A cell based solely on the emf generated because of a difference in a

concentration is called a concentration cell

o E.g. Nickel

Oxidation of Ni(s) occurs in the half-cell containing the more dilute

solution, thereby increasing the concentration of Ni2+(aq)

n (the number of electron being transferred) is equal to 2

20.7: Batteries and Fuel Cells

A battery is a portable, self-contained electrochemical power source that consists of

one or more voltaic cells

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o When cells are connected in series, the battery produces a voltage that is the

sum of the emfs of the individual cells

Higher emfs can also be achieved by using multiple batteries in series

o Some batteries are primary cells, meaning that they cannot be recharged

A secondary cell can be recharged from an external power source after

its emf has dropped

Lead-Acid Battery

o A 12-V lead-acid automotive battery consists of six voltaic cells in series, each

producing 2 V

The electrode reactions that occur during discharge are

Because the reactants are solids, there is no need to separate the cell

into anode and cathode compartments

Solids are excluded from the reaction quotient Q, the relative

amounts of Pb(s), PbO2(s), and the PbSO4(s) have no effect on

the emf, helping the battery maintain a relatively constant emf

o Lead-acid batter can be recharged

During recharging, an external source of energy is used to reverse the

direction of the overall cell reaction

Alkaline Battery

o Alkaline batteries are nonrechargeable (primary battery)

o The anode of this battery consists of powdered zinc metal immobilized in a gel

in contact with a concentrated solution of KOH

The cathode is a mixture of MnO2(s) and graphite

Nickel-Cadmium, Nickel-Metal-Hydride, and Lithium-Ion Batteries

o Nickel-cadmium (nicad) battery

During discharge, cadmium metal is oxidized at the anode of the

battery while nickel oxyhydroxide is reduced at the cathode

Cadmium is a toxic heavy metal

Its use increases the weight of batteries and provides an

environmental hazard

o Nickel-metal-hydride (NiMH) batteries

Cathode reaction of NiMH is the same as that for the nickel-cadmium

batteries

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The anode consists of a metal alloy that has the ability to absorb

hydrogen ions

During the oxidation at the anode, the hydrogen atoms lose

electrons, and the resulting H+ ions react with OH- ions to form

H2O

Due to the robustness of the batteries toward discharge and recharge,

the batteries can last up to 8 years

o Lithium-ion (Li-ion) battery

Lithium is a very light element and therefore achieve a greater energy

density—the amount of energy stored per unit mass—than nickel-

based batteries

It is based on the ability of Li+ ions to be inserted into and removed

from certain layered solids

Hydrogen Fuel Cells

o The direct production of electricity from fuels by a voltaic cell could, in

principle, yield a higher rate of conversion of the chemical energy of the

reaction

Voltaic cells that perform this conversion using conventional fuels,

such as H2 and CH4 are called fuel cells

Strictly speaking, fuel cells are not batteries

o In the fuel cell for the reaction of hydrogen and oxygen, the anode and

cathode are separated by a thin polymer

Protons are able to pass through these polymers but electrons cannot

20.8: Corrosion

Corrosion reactions are spontaneous redox reactions in which a metal is attacked by

some substance in its environment and converted to an unwanted compound

For nearly all metals, oxidation is a thermodynamically favorable process in air at

room temperature

o When oxidation process is not inhibited in some way, it can be very

destructive to whatever object is made from the metal

o Oxidation can form an insulating protective oxide layer that prevents further

reaction of the underlying metal

Corrosion of Iron

o Rusting of iron requires both oxygen and water

Other factors—such as the pH of the solution, the presence of salts,

contact with metal more difficult to oxidize than iron, and stress on the

iron—can accelerate rusting

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o Corrosion of iron is electrochemical in nature

Electrons can move through the metal from a region where oxidation

occurs to another region where reduction occurs

Preventing the Corrosion of Iron

o Iron is often covered with a coat of paint or another metal such as tin or zinc

to protect its surface against corrosion

If the coating is broken and the iron is exposed to oxygen and water,

corrosion will begin

o Galvanized iron, which is iron coated with a thin layer of zinc, uses the

principles of electrochemistry to protect the iron from corrosion even after

the surface coat is broken

The Zn(s) is easier to oxidize than Fe(s)

Thus, even if the zinc coat is broken, the zinc will serves as the anode

and is corroded instead of iron

o Protecting a metal from corrosion by making it the cathode in an

electrochemical cell is known as cathodic protection

The metal that oxidized while protecting the cathode is called the

sacrificial anode

20.9: Electrolysis

Electrical energy can be used to cause nonspontaneous redox reactions to occur

o Such processes, which are driven by an outside source of electrical energy, are

called electrolysis reactions and take place in electrolytic cells

An electrolytic cell consists of two electrodes in a molten salt or a

solution

A battery or some other source of direct electrical current acts as an

electron pump, pushing electrons into one electrode and pulling them

from the other

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The electrode of the electrolytic cell that is connected to the negative terminal of the

voltage source is the cathode of the cell

Several practical applications of electrochemistry are based on active electrodes—

electrodes that participate in the electrolysis process

o Electroplating, for example, uses electrolysis to deposit a thin layer of one

metal on another metal to improve beauty or resistance to corrosion

Quantitative Aspects of Electrolysis

o For any half-reaction, the amount of a substance that is reduced or oxidized in

an electrolytic cell is directly proportional to the number of electrons passed

into the cell

o A coulomb is the quantity of charge passing a point in a circuit in 1 s when the

current is 1 ampere (A)

Coulombs = amperes × seconds

o Electrons can be thought of as reagents in electrolysis reactions

Electrical Work

o

-wmax means that a voltaic cell does work on its surrounding

o Eext > Ecell is needed to bring about a nonspontaneous electrochemical process

o When an external potential Eext is applied to a cell, the surroundings are doing

work on the system

n is the number of moles of electrons forced into the system by

the external potential

o n × F is the total electrical charge supplied to the system

by the external source of electricity

o watt (W) is a unit of electrical power

Watt-second is a joule

Kilowatt-hour (kWh) is equal to 3.6 × 106 J