Thermodynamics has many applications other than expanding gases inside steam

engines. One of the most important areas of applied thermodynamics is

ELECTROCHEMISTRY.

Electrochemistry is an area of chemistry that deals with the inter-conversion of

electrical energy and chemical energy. It is typically investigated through the use of

electrochemical cells, systems that incorporate a redox reaction to produce or utilize electrical

energy.

There are two types of electrochemical cells:

• One type of cell does work by releasing free energy from a spontaneous reaction (∆G

< 0) to produce electricity. A battery houses this type of cell (Voltaic or galvanic cell)

• The other type of cell does work by absorbing free energy from a source of electricity

to drive a nonspontaneous (∆G > 0). (Electrolytic cells).

Redox Reactions and Electrochemical Cells

Electrochemical processes are redox (oxidation-reduction) reactions in which the energy

released by a spontaneous reaction is (converted to electricity in which electricity is used to

drive a nonspontaneous chemical reaction (electrolysis) to redox reactions, electrons are

transferred from one substance to another.

The reaction between zinc metal and sulfur is an example of a redox reaction:

0 0 +2 -2

Zn + S → ZnS

RA OA

GEROA (Gain Electron, Reduction, Oxidizing Agent)

LEORA (Loss Electron, Oxidation, Reducing Agent)

Recall that the numbers above the elements are the oxidation numbers of the elements.

By assigning an oxidation number to each atom, we can see which species was oxidized and

which reduced and, from that, which is the oxidizing agent and which is the reducing agent.

The loss of electrons by an element during oxidation is marked by an increase in the element's
oxidation number. In reduction, there is decrease in oxidation number resulting from a gain of

electrons by an element. In the preceding reaction Zn metal is oxidized and S is reduced. Keep

in mind three key points:

• Oxidation (electron loss) always accompanies reduction (electron gain). The oxidizing

agent is reduced, and the reducing agent is oxidized.

• The total number of electrons gained by the atoms/ ions of the oxidizing agent always

equals the total number lost by the atoms/ions of the reducing agent.

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Electrochemistry

SCI 401 – Chemistry for Engineers

• Redox reaction can be defined as one in which the oxidation numbers of the species

change.

Rules for Assigning an Oxidation Number (O.N.)

1. The oxidation number of an element in the free or uncombined state is always zero

(including

2. The algebraic sum of the oxidation numbers of all the atoms in the formula of a

compound is zero.

3. The oxidation number of an ion is the same as the charge on the ion.

4. The sum of the oxidation numbers of the atoms in a polyatomic ion must be equal to

the charge on the ion.

Rules for Specific Atoms or Periodic Table Groups

1. Group IA elements are always +1

2. Group Il A are always + 2.

3. Hydrogen is usually +1 except in hydrides.

4. Oxygen is usually -2, except in peroxides

5. Group VII A elements are -1.

LEORA (Loss Electron, Oxidation, Reducing Agent)

Example:

Use oxidation numbers to decide which of the following are redox reactions. For redox

reaction, identify the species that undergo the reduction and oxidation reaction; and the

oxidizing and reducing agent.

1. CaO(s) + CO2 (g) → CaCO3(s) not redox reaction CO3

2. 4KNO3(s) → 2K2O(s) + 2N2(g) + 5O2 (g) redox reaction NO3

Balancing Redox Reactions

We balance a redox reaction by making sure that the number of electrons lost by the

reducing agent equals the number of electrons gained by the oxidizing agent. There are two

methods used to balance redox equations: the oxidation number method and the half-reaction

+2 -2

x + 2(-2) = 0;

x =+4

+4 -2 +2 -2

x + 3(-2) = -2;

x = +4

+4 -2

+1 -1 +1 -2 0 0

x + 3(-2) = -1;
+1 + x = 5

5 -2

+5 -2

OA RA

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Electrochemistry

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method — but only the second method will be discussed in this section. The following steps

are used in balancing a redox reaction by the half-reaction method.

1. Write the unbalanced equation for the reaction in ionic form.

2. Separate the equation into two half-reactions.

3. Balance each half-reaction for number and type of atoms and charges. Balance the

atoms other than O and H in each half-reaction separately.

4. For reactions in an acidic medium, add H2O to balance the O atoms and H

to balance

the H atom.

5. Add electrons (e

) on one side of each half-reaction to balance the charges. If

necessary, equalize the number of electrons in the two half-reactions by multiplying

one or both half-reactions by appropriate coefficients,

6. Add the two half-reactions and balance the final equation by inspection. The electrons

on both sides must cancel.

charges on both sides of the equation.

Example:

1. Balance the equation showing the oxidation of Fe+2 ions to Fe+3 ions by dichromate ions in

an acidic medium. The dichromate ions are reduced to Cr3+ ions.

Step 1) Fe+2 + Cr2O7

-2→Fe+3 + Cr+3

Step 2 - 3) Fe+2 →Fe+3

Cr2O7

-2→ 2Cr+3

Step 4) Cr2O7

-2 + 14H

+→ 2Cr+3 + 7H2O

Step 5) Fe+2 →Fe+3 + e

6Fe+2 →6Fe+3 + 6e

6e

- + Cr2O7

-2 + 14H

+→ 2Cr+3 + 7H2O Cr2O7

-2

(-2 + 14(+1) = 12) (2(3) + 7(2)(1) + 7(-2) = 6)

Step 6) 6Fe+2 →6Fe+3 + 6e

- Oxidation

6e

- + Cr2O7

-2 + 14H

+→ 2Cr+3 + 7H2O Reduction

6Fe+2+ Cr2O7

-2 + 14H

+→ 6Fe+3 + 2Cr+3 + 7H2O Overall Reaction

2x + 7(-2) = -2;

x=6

+6 -2 +1 -2 +6 -2

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Electrochemistry

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Electrochemical Cells

Electrochemical cell is the experimental apparatus for generating electricity through

the use of a spontaneous redox reaction (∆G < 0). An electrochemical cell is sometimes

referred to as a galvanic cell or voltaic cell, after tile scientists Luigi Galvani and Alessandro

Volta, who constructed early version of this device.

The separation of half-reactions is the essential idea behind a voltaic ceil. The

components of each half-reaction are placed in a separate container, or half-cell, which consist
of one electrode dipping into an electrolyte solution. The two half-cells are joined by the

circuit, which consists of a wire and a salt bridge (the inverted U tube in the figure). In order

to measure the voltage generated by the cell, a voltmeter is inserted in the path of the wire

connecting the electrodes. A switch closes or opens the circuit. The oxidation half-cell (anode

compartment) is shown on the left and the reduction half-cell (cathode compartment) on the

right.

Here are the key points about the Zn/Cu2+

voltaic cell:

1. The oxidation half-cell. The anode compartment consists of a zinc bar (the anode)

immersed in a Zn2+ electrolyte (solution of ZnS04). The zinc bar conducts the released

electrons out of its half-cell.

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2. The reduction half-cell. The cathode compartment consists of a copper bar (the

cathode) immersed in a Cu2+ electrolyte (solution of CuS04). Copper metal conducts

electrons into its half-cell.

3. Relative charges on the electrodes. The electrode charges are determined by the source

of electrons and the direction of electron flow through the circuit. The electrons flow

left to right through the wire to the cathode. In any voltaic cell, the anode is negative

and the cathode is positive.

4. The purpose of the salt bridge. The oxidation half-cell and the reduction half-cell

originally contain neutral solutions. If the half-cells do not remain neutral, the resulting

charge imbalance would stop cell operation. To avoid this situation and enable the cell
to operate, the two half- cells are joined by a salt ridge, which act as a "liquid wire,"

allowing ions to flow through both compartments and complete the circuit. Therefore,

salt bridge maintains the neutrality of the solutions.

The conventional notation for representing electrochemical cells is the cell diagram

(cell notation). If we assume that the concentrations of Zn2+ and Cu2+ ions are

Zn(s) | Zn2+ (1M) || Cu2+ (1M) | Cu(s)

• The anode is written first, to the left of the double lines, while cathode is on

right.

• The single vertical line represents a phase boundary.

• The double vertical lines denote the salt bridge.

We can write the half- cell reactions as follows:

Anode (oxidation): Zn(s) Zn2+

(1M) + 2e- Eoox = 0.76V

Cathode (reduction): Cu2+

(1M) + 2e- Cu(s) Eored = 0.34V

Overall: Zn(s) + Cu

2+ (1M) Zn

2+

(1M) + Cu(s) Eocell= 1.10V

Standard Reduction Potentials

For a reduction reaction at an electrode when all solutes are 1M and all gases are at 1
atm, the voltage is called the standard reduction potential (EO). The larger (more positive)

the EO value, the greater the tendency for the substance to be reduced, and therefore the stronger

its tendency to act as an oxidizing agent (gains electrons more readily). The smaller (more

negative) EO value, the greater the tendency for the substance to be oxidized and act as reducing

agent (loses electrons more readily).

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Let us consider Cu2+, H+, and Zn2+, three oxidizing agent present in the voltaic cell.

We can rank their relative oxidizing strengths by writing each half-reaction as gain of electrons

(reduction), with its corresponding standard electrode potential:

Cu2+

gains two electrons more readily than I-1+, which gains them more readily than

Zn2+

. In terms of strength as an oxidizing agent, therefore, Cu+2> H+ > Zn2+. Therefore Cu2+

is

the strongest oxidizing agents, and Zn is the strongest reducing agent.

Example:

Arrange the following species in order of increasing strength as oxidizing agents under

standard-state conditions: Ce4+, O2, H2O2

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Electrochemistry
SCI 401 – Chemistry for Engineers

*For all half-reactions the concentration is 1M for dissolved species and the pressure is atm

for gases. These are standard state values.

Calculating the standard emf (Eo

) of an Electrochemical Cell

In an electrochemical cell, electric current flows from one electrode to the other (from

anode to cathode) because there is a difference in electrical potential energy between the

Please refer to the links for the Table of Standard Reduction Potential

https://chem.libretexts.org/Bookshelves/Ancillary_Materials/Reference/Reference_Tables

/Electrochemistry_Tables/P2%3A_Standard_Reduction_Potentials_by_Value

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electrodes. The difference in electrical potential between the anode and cathode is measured

by a voltmeter, and the reading (in volts) is called cell voltage. However, two other terms,

electromotive force or emf (E) and cell potential are also used to denote cell voltage. If all

solutes have a concentration of 1 M and all gases have a pressure of 1 atm (standard

conditions), the voltage difference between the two electrodes of the cell is called the

standard emf (EOcell). The standard emf is the sum of the standard oxidation potential and

the standard reduction potential for each half-reaction.

E0

cell = E0

ox + E0
red

We can use the sign of the emf of the cell to predict the spontaneity of redox reaction. Under

standard-state conditions for reactants and products, the redox reaction is spontaneous in the

forward direction if the standard emf of the cell is positive. If it is negative, the reaction is

spontaneous in opposite direction.