Electrochemistry Course: B.Tech. Subject: Engineering Chemistry Unit: V(A)
Jul 19, 2015
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.
29
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