Electrochemistry

Course: B.Tech.

Subject: Engineering Chemistry

Unit: V(A)

Arrhenius Theory of Electrolytic

dissociation

Svante August Arrhenius (19 February 1859 – 2

October 1927) was a Swedish Scientist, and one

of the founders of the science of physical

chemistry.

The Arrhenius equation, lunar crater Arrhenius

and the Arrhenius Labs at Stockholm University

are named after him.

In order to explain the properties of electrolyticsolutions, Arrhenius put forth, in 1884, acomprehensive theory which is known as theory ofelectrolytic dissociation or ionic theory.

The main points of the theory are:

An electrolyte, when dissolved in water, breaks up intotwo types of charged particles,

one carrying a positive charge

the other a negative charge

These charged particles are called ions.

Positively charged ions are termed cations

Negatively charged as anions

Theory

In its modern form, the theory assumes that solid electrolytes are composed of ions which are held together by electrostatic forces of attraction.

When an electrolyte is dissolved in a solvent, these forces are weakened and the electrolyte undergoes dissociation into ions. The ions are solvated.

A+B- --> A+ + B-

A+B-+ aq --> A+(aq)+B- (aq)

NaCl Na+ + Cl-

K2SO4 2K+ + SO42

The process of splitting of the molecules

into ions of an electrolyte is called

ionization.

The fraction of the total number of

molecules present in solution as ions is

known as degree of ionization or degree

of dissociation.

It is denoted by α= (Number of molecules

dissociated into ions)/(Total number of

molecules)

Ions present in solution constantly re-unite to form neutral molecules and, thus, there is a state of dynamic equilibrium between the ionized and non-ionized molecules.

[A+ ][B- ] /[AB] =K

K is known as ionization constant.

The electrolytes having high value of K are termed strong electrolytes

those having low value of K as weak electrolytes

Applying the law of mass action to

above equilibrium When an electric current is passed through

the electrolytic solution, the positive ions

(cations) move towards cathode and the

negative ions (anions) move towards

anode and get discharged, i.e., electrolysis

occurs.

The ions are discharged always in

equivalent amounts, no matter what their

relative speeds are.

Transport Number

Transport number or transference number is the ratio of the current carried by a given ionic species through a cross section of an electrolytic solution to the total current passing through the cross section.

The transport number is equal to the ratio of the velocity, or mobility, of a given ion to the sum of the velocities, or mobilities, of the cation and anion.

It is a characteristic dependent on the

mobilities of all the ions in the

electrolytic solution,

on the concentrations of the ions

on the temperature of the solution

The transport number is usually

determined by the Hittorf method—that

is, by the change in the concentrations of

the ions near the electrodes.

Kohlrausch’s law

where,

According to Kohlrausch’slaw. “conductivity

of ions is constant at infinite dilution and it

does not depend on nature of co-ions.”

2

For AxBy type electrolyte,

Here Z+and Z- are the charges present on cation and anion.

Ksp, the solubility-product constant

An equilibrium can exist between a partially soluble substance and its

solution:

For example:

BaSO4 (s) Ba2+ (aq) + SO42- (aq)

When writing the equilibrium constant expression for the dissolution of BaSO4, we remember that the concentration of a solid is constant.

The equilibrium expression is therefore:

K = [Ba2+][SO42-]

K = Ksp, the solubility-product constant.

Ksp = [Ba2+][SO42-]

The Solubility Expression

AaBb(s) aAb+ (aq) + bBa- (aq)

Ksp = [Ab+]a [Ba-]b

Example: PbI2 (s) Pb2+ + 2 I-

Ksp = [Pb2+] [I-]2

The greater the ksp the more soluble the solid is in H2O.

Solubility and Ksp

Three important definitions:

1) solubility: quantity of a substance that

dissolves to form a saturated solution

2) molar solubility: the number of moles of

the solute that dissolves to form a liter of

saturated solution

3) Ksp (solubility product): the equilibrium

constant for the equilibrium between an

ionic solid and its saturated solution

An oxidation-reduction (redox) reaction

involves the transfer of electrons (e - ).

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The oxidation numbers of the atoms will change….

one goes up (oxidation) and one goes down (reduction)

Sodium transfers its electrons to chlorine

Redox Reaction:Oxidation-Reduction

Find the oxidation numbers of each element in

a reaction and see which ones have changed.

Rules for oxidation number

◦ An element that is not in a compound has an oxidation number of zero (0)

◦ Group 1 Metals are always 1+

◦ Group 2 Metals are always 2+

◦ Fluorine is always 1-

◦ Oxygen is always 2- except when combined with F (OF2) or the peroxide ion (O2

2-)

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Reduction is the gain of electrons.

21

Nonmetals gain electrons to form – ions

•The oxidation number goes down

(reduces)

A half-reaction can be written to represent

reduction.

22

Cu2+ + 2e- Cu0

In reduction half reactions,

electrons are written on the left

because electrons are gained

Oxidation is the loss of electrons.

23

Metal atoms lose electrons to become + ions

The oxidation numbers go up (increases)

Cr2+ Cr4+ + 2e-

2N3- N20 + 6e-

A half-reaction can be written to represent oxidation.

24

Zn0 Zn2+ + 2e-

In oxidation half reactions,

electrons are written on the right

because electrons are lost

The sum of the oxidation numbers

of all the atoms in a compound is

zero.

Na2SO4

◦ Na is +1 because it is a group 1 metal

◦ O is -2

◦ The oxidation number of Sulfur must be calculated

2(+1) + X + 4(-2) = 0

(2 ) + X + (-8) =0

X = +6

CuO

Oxygen is -2

The oxidation number of

copper must be

calculated

X + -2 = 0

X = +2

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The sum of the oxidation numbers

of all the atoms in a polyatomic ion

is the charge of the ion.

PO43-

Oxygen is 2-

The oxidation number of

phosphorous must be

calculated

X + 4(-2) = -3

X + (-8) = -3

X = +5

NO3-

Oxygen is 2-

The oxidation number of

nitrogen must be

calculated

X + 3(-2) = -1

X = 5+

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Electrochemical and concentration

cells An electrochemical cell is a device capable

of either generating electrical energy from chemical reactions or facilitating chemical reactions through the introduction of electrical energy.

A common example of an electrochemical cell is a standard 1.5-volt "battery".

(Actually a single "Galvanic cell"; a battery properly consists of multiple cells, connected in either parallel or series pattern.)

A voltaic cell spontaneously converts chemical energy to electrical energy.

28

Batteries are voltaic cells

Electrons flow from the anode (- electrode) to

the cathode (+ electrode) through the wire in a

voltaic cell.

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An Ox -oxidation

takes place…electrons

are lost.

Red Cat -reduction

takes place…electrons

are gained.

Zn Zn2+ + 2e-Cu2+ + 2e - Cu0

- +

Electrons

released

here by

oxidation

Electrons

needed

here for

reduction

e-

e-

e-

e- e- e- e-

e-

e-

e-

e-

The salt bridge completes the circuit

allows ions to flow from one ½ cell to

the other ½ cell to maintain neutrality.

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Zn Zn2+ + 2e- Cu2+ + 2e - Cu0

- +

4

Electrolysis

An electrolytic cell requires electrical

energy to produce chemical change.

This process is known as electrolysis.

31

Uses of Electrolytic cells

Recharging a battery

Electroplating

◦ During copper plating, Cu2+ ions are reduced to

Cu0 metal at the cathode (Red Cat) which is the

negative electrode

Electrolysis

◦ The Hoffman apparatus uses electricity to break

water apart into hydrogen + oxygen

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Galvanic Cells

• Galvanic Cell: Electrochemical cell in which

chemical reactions are used to create spontaneous

current (electron) flow.

Salt bridge

Zn2+ Cu2+

Na+Zn Cu

SO42–

Voltmeter

(–) (+)

Oxidation half-reactionZn(s)

Salt bridge

Zn2+ Cu2+

Na+Zn Cu

SO42–

Zn2+(aq) + 2e–

Voltmetere–

Anode

(–) (+)

Zn2+Zn

Oxidation half-reactionZn(s)

Salt bridge

Zn2+ Cu2+

Na+Zn Cu

SO42–

Zn2+(aq) + 2e–

Voltmetere–

2e– lost

per Zn atom

oxidized

Anode

(–) (+)

e–

Zn2+Zn

Oxidation half-reaction

Reduction half-reaction Cu2+(aq) + 2e–

Zn(s)

Salt bridge

Zn2+ Cu2+

Na+Zn Cu

SO42–

Zn2+(aq) + 2e–

Cu(s)

Voltmetere– e–

2e– lost

per Zn atom

oxidized

Anode

(–)

Cathode

(+)

e–

Cu2+e–Cu

2e– gained

per Cu2+ ion

reduced

Zn2+Zn

Oxidation half-reaction

Reduction half-reaction Cu2+(aq) + 2e–

Zn(s)

Salt bridgeAnode

(–)

Cathode

(+)

Zn2+ Cu2+

Na+Zn Cu

SO42–

Zn2+(aq) + 2e–

Cu(s)

Voltmetere– e–

2e– lost

per Zn atom

oxidized

e–

Cu2+e–Cu

2e– gained

per Cu2+ ion

reduced

Zn2+Zn

Oxidation half-reaction

Reduction half-reaction

Overall (cell) reaction

Zn(s) + Cu2+(aq)

Cu2+(aq) + 2e–

Zn(s)

Salt bridge

Zn2+ Cu2+

Na+Zn Cu

SO42–

Zn2+(aq) + 2e–

Cu(s)

Zn2+(aq) + Cu(s)

Voltmetere– e–

Anode

(–)

Cathode

(+)

2e– lost

per Zn atom

oxidized

e–

3

References

1.Engineering Chemistry by Jain and Jain

2. https://www.askiitians.com/iit-jee-

chemistry/physical-chemistry/kohlrausch-

law.aspx

3. https://chemwiki.ucdavis.edu