Electrochemistry

The branch of chemistry that deals with the interrelationship between electrical energy and chemical transformation is called electrochemistry.

Electrochemical cell

The assembly of two electrodes that are either set in the same container or different containers containing different electrolytes so that the arrangement would generate electricity due to chemical reaction or consume electricity to cause the chemical reaction. This arrangement is known as an electrochemical cell.
There are two types of it.

i. Electrolytic cell
ii. Galvanic cell

Differences between Electrolytic cell and Galvanic cell
Electrolytic cellGalvanic cell
It is a device where electricity is consumed in order to cause a chemical reaction. eg. electrolysis
It is a device where electricity is generated due to a redox reaction. eg. dry cell
Both electrodes are suspended in the same container containing the same electrolyte.Two electrodes are suspended in different containers containing different electrolytes.
Cell reaction is non-spontaneous in nature.Cell reaction is spontaneous in nature.
No salt bridge is used.Salt bridge is used.
Anode is positive and cathode is negative.The anode is negative and the cathode is positive.
Electrode potential

To understand the origin of electrode potential, let us suppose a metallic plate is placed in its solution as shown in the diagram.

electrochemistry- electrode potential

When a metal plate is placed in its solution, then the metal atom from the metal plate tends to oxidize into a metal ion and go to the solution leaving behind the electron on the metal plate. This entire process can be shown in the following oxidation half-reaction.

M \rightarrow M^{n+} + ne^{-}

When the electron density on the metal plate becomes high, it acts as a negative electrode. As a result, positive ions from the solution are attracted towards the metal plate in order to neutralize the electrode and thus free metal atom is formed due to reduction. This process can be shown in the following reduction half-reaction.

M^{n+} + ne^{-} \rightarrow M

When the oxidation and reduction processes are balanced at the same time, the system reaches a state of dynamic equilibrium. In this case, there is the formation of two different layers between metal and solution. This charge layer is called an electrical double layer.

Due to net charge separation, an electrical double layer develops a certain amount of potential known as electrode potential or since only one electrode is used, it is also known as single electrode potential. Electrode potential is defined as the potential developed at the solid-liquid interface when any metallic rod is dipped into an electrolyte under equilibrium. It is denoted by E. It depends upon the following factors:

  • Nature of metallic rod (electrode)
  • Concentration of solution
  • Temperature of the system

If the concentration of the solution is one molar at 25°C (298K) and 1 atm pressure (760 mm Hg), under this condition, the electrode potential is called standard electrode potential. It is denoted by E°. There are two types of standard electrode potential.

  1. Standard oxidation potential: It measures the tendency of losing electrons. It is represented by E°M/Mn+.
  2. Standard reduction potential: It measures the tendency of gaining electrons. It is represented by E°Mn+/M.

Both potentials are equal but opposite in sign.
i.e. Standard oxidation potential = – Standard reduction potential

In earlier days, electrode potential was measured in terms of oxidation potential but nowadays, according to IUPAC recommendation, electrode potential is measured in terms of reduction.

 \left.\begin{matrix}
  & E^{o}Zn/Zn^{++} = +0.77\\ 
 & E^{o}Cu/Cu^{++} = -0.34\\
\end{matrix}\right\} \underset{(Not\ Recommended)}{standard\ oxidation\ potential}\\

\left.\begin{matrix}
  & E^{o}Zn^{++}/Zn = -0.77\\ 
 & E^{o}Cu^{++}/Cu = +0.34\\
\end{matrix}\right\} \underset{(Recommended\ by\ IUPAC)}{standard\ reduction\ potential}\\
Types of Electrodes in Electrochemistry

We can’t measure the absolute potential of a single electrode cell or single half cell because it cannot form a complete close circuit connecting through a voltmeter due to a lack of connection terminal. So, in order to measure the potential of such a half cell, it should be coupled with any other standard electrode known as a reference electrode. They are the electrode with known potential and are used to determine the potential of unknown electrodes. There are two types of reference electrodes. They are:

  1. Primary reference electrode eg. standard hydrogen electrode (SHE)
  2. Secondary reference electrode eg. calomel electrode
Construction of standard hydrogen electrode (SHE)

It consists of a platinum wire which is sealed in a glass tube. The lower end of the platinum wire is attached to the platinum foil which is plated by platinum black. It helps with the adsorption of hydrogen and acts as a site for the reaction. This whole adjustment is dipped into the beaker of 1 M HCl solution with hydrogen gas at 1 atm pressure at 25°C which is constantly bubbled into a solution as shown below:

Standard hydrogen electrode

If hydrogen gas is maintained at 1 atm pressure and 25°C temperature and platinum foil is immersed into 1M HCl solution, the electrode is called a standard hydrogen electrode. Depending upon the nature of the counter electrode, the standard hydrogen electrode acts as a cathode as well as an anode.
If it acts as an anode:

H_{2}(g) \rightleftharpoons  2H^{+} + 2e^{-} (Oxidation\ half)

Then it is represented as:

^\circleddash Pt, H_{2}(g)(1atm)/H^{+}(aq.1M)//........)

If it acts as a cathode:

2H^{+} + 2e^{-} \rightleftharpoons H_{2(g)}\ (Reduction\ half)

Then it is represented as:

(.....//H^{+}(aq.1M)/H_{2}(g)(1\ atm),Pt^\oplus 

The electrode potential of a standard hydrogen electrode is arbitrarily taken as zero. In fact, the cell potential of the hydrogen electrode is neither zero nor independent of temperature. All other electrodes are calibrated with respect to standard hydrogen electrodes.
It is somewhat difficult to handle and can’t be used in the presence of oxidizing or reducing agents. The platinum black used is easily poisoned by mercury and H2S. Therefore other reference electrodes such as calomel electrodes are used. Such electrodes are called secondary reference electrodes.

Advantages of standard hydrogen electrode
  • It can be used over the entire pH range.
  • It gives no salt error.
  • It is highly accurate.
Disadvantages of standard hydrogen electrode
  • It is difficult to maintain a 1 molar concentration of H+ ion.
  • 1 atm pressure of H2 gas cannot be maintained uniformly.
  • The hydrogen electrode gets poisoned even if traces of impurities are present.
  • Platinum used is expensive.
Calomel electrode

It is a secondary reference electrode and its potential is determined with respect to standard hydrogen electrode. It is preferred over hydrogen electrodes due to its easiness of handling.

Construction

Calomel electrode consists of mercury at the bottom over which mercury-mercurous chloride is kept. KCl solution is placed over the mercury-mercurous chloride. A platinum wire is inserted for the electrical connection. The KCl solution used in the calomel electrode may be saturated or normal or decinormal. The electrode potential of the calomel electrode depends upon the concentration of KCl used. If a saturated KCl solution is used, the electrode is known as a saturated calomel electrode.

calomel electrode
Electrode reaction

Depending upon the electrode potential value with respect to the counter electrode, the calomel electrode acts as an anode half or cathode half cell.
If the calomel electrode acts as an anode:

\begin{align*}2Hg \rightarrow &  Hg_{2}^{2+} + 2e^{-}\\
Hg_{2}^{2+} +2Cl^{-} &\rightarrow   Hg_{2}Cl_{2}\\
\hline
2Hg + 2Cl^{-} \rightarrow &  Hg_{2}Cl_{2} + 2e^{-}

\end{align*}

Then it is represented as:

\circleddash (Pt)Hg(l)/Hg_{2}Cl_{2}(s),KCl(aq)//......

If it acts as a cathode:

\begin{align*}Hg_{2}Cl_{2} \rightarrow &  Hg_{2}^{2+} + 2Cl^{-}\\
Hg_{2}^{2+} +2e^{-} &\rightarrow   2Hg\\
\hline
Hg_{2}Cl_{2} + 2e^{-} \rightarrow &  2Hg + 2Cl^{-}
\end{align*}

Then it is represented as:

...//KCl(aq.)/Hg_{2}Cl_{2}(s)/Hg(l)Pt\oplus  
Advantages of calomel electrode
  1. It is easy to set up.
  2. It is convenient and easy to transport.
  3. It is compact and smaller in size.
  4. No separate salt bridge is required as it has already a side tube containing KCl solution.
  5. Potential doesn’t change with time and a slight change in temperature.
Electrochemical series

The standard reduction potential of several electrodes has been measured by using a standard hydrogen electrode(reference electrode). These electrodes are arranged in order of increasing or decreasing their standard reduction potential. The vertical arrangement of elements according to increasing standard reduction potential as compared to the electrode potential of standard hydrogen electrode from top to bottom is called electrochemical series (ECS). It is also called electrochemical activity or electromotive series.

Some electrochemical series list
electrochemical series
Applications of electrochemical series

1. To predict the cathode and anode of a galvanic cell: The electrode having lower standard reduction potential acts as an anode while the electrode having higher standard reduction potential acts as a cathode. For example,

\begin{align*}E^{o}\ Zn^{++}/Zn = -0.77V\\
E^{o}\ Cu^{++}/Cu = +0.34V
\end{align*}

Here, Eo Zn++/Zn < Eo Cu++/Cu. So, Zinc acts as an anode and copper acts as a cathode.

2. To compare the relative reactivity of elements: The element having a lower value of standard reduction potential is relatively more reactive than the elements having a higher value of standard reduction potential. For example,

\begin{align*}
E^{o} Zn^{++}/Zn = -0.77V\ \ \ E^{o}Ag^{+}/Ag = +0.80V\\
E^{o} Cu^{++}/Cu = +0.34V\ \ \ E^{o}Fe^{++}/Fe = -0.44V
\end{align*}

Here, Eo Zn++/Zn value is less than other elements. So, Zn is more reactive while Eo Ag+/Ag value is higher than other elements. So, Ag is the least reactive among them. Hence the reactivity order is Zn>Fe>Cu>Ag.

3. To predict the relative strength of oxidizing and reducing agents: The substance having a high value of Eo shows a high tendency to gain the electron. So, it is a relatively strong oxidizing agent than the substance having a lower value of Eo.
On the other hand, the substance having a lower value of Eo has a high tendency to lose the electron. So, it is a relatively strong reducing agent than the substance having a higher value of Eo. For example,

\begin{align*}
E^{o} I_{2}/I^{-} = +0.54V\\
E^{o} Cl_{2}/Cl- = +1.36V
\end{align*}

Here, chlorine is a stronger oxidizing agent than iodine but on the other hand, iodine is a stronger reducing agent than chlorine.
In the electrochemical series, Fluorine is the strongest oxidizing agent and lithium is the strongest reducing agent.

4. To calculate the standard emf of the cell: By knowing the Eo of cathode and anode, we can calculate the standard emf of the cell by applying the following relation: For example,

\begin{align*}
E^{o}\ Cell = E^{o}& Cathode - E^{o} Anode\\
E^{o}Zn^{++}/Zn &= -0.77V\\
E^{o} Cu^{++}/Cu &= +0.34V
\end{align*}

Since, Eo Zn++/Zn < Eo Cu++/Cu, Zinc acts as an anode and copper acts as a cathode.
So, Eo Cell = 0.34 – (-0.77) = 1.11V

5. To predict the metal which displaces hydrogen from acid or not: The metal having low value of Eo than hydrogen can displace hydrogen from acid. For example,

\begin{align*}
E^{o} Zn^{++}/Zn = -0.77V\\
E^{o} Cu^{++}/Cu = +0.34V\\
E^{o} H^{+}/H2 = 0.00V
\end{align*}

Here, Zn is more reactive than H2. So, it can displace hydrogen from acid whereas Cu cannot.

6. To predict whether the cell reaction is spontaneous or not: If the emf of the cell is positive, the reaction is feasible or spontaneous. If the emf of the cell is negative, the reaction is not feasible or non-spontaneous.

Voltaic cell (Typical Galvanic cell)

A voltaic or galvanic cell is constructed by combining the anode half and cathode half cells. Out of two half cells, the electrode having lower reduction potential acts as an anode and the electrode having a higher reduction potential acts as a cathode. Hence, the anode is the negative electrode and it is connected to the negative terminal of the circuit. Cathode is the positive electrode and it is connected to the positive terminal of the circuit.
The anode half and cathode half cells are connected externally by an electric circuit while internally by a salt bridge. The potential developed due to net cell reaction is read out from a potentiometer or galvanometer.

galvanic cell

After the completion of the electric circuit, the following observations are seen:

  1. The weight of the anode is decreased.
  2. The concentration of ions in an anodic salt solution is increased.
  3. The concentration of ions in the cathodic salt solution is decreased.
  4. The weight of the cathode is increased.
  5. There is a flow of electrons from the anode to the cathode and hence the flow of electric current is from the cathode to the anode.
Salt bridge

It is a U-shaped glass tube containing inert electrolyte solution such as KCl, KNO3, K2SO4 or NH4NO3 in agar-agar or jelly medium which is used to connect internally two half cells.
Usually 3-4% agar-agar solution is heated in a saturated solution of inert electrolyte. Then, it is subjected to a U-shaped glass tube and allowed to set in the form of gel after cooling.

In a galvanic cell, the anode undergoes oxidation where cations move into the electrolyte. So, the electrolytes of the anode half-cell become more positively charged. Similarly, due to the reduction of cation at the cathode, the electrolyte of the cathode half cell becomes more negatively charged. It shows that when galvanic cell functions, net charge separation takes place between anode and cathode half cells. For its functioning, the electrical neutrality of the electrolyte must be maintained by allowing cations present in the salt bridge to move into the cathode half-cell while anions move into the anode half-cell.

The essential requirement of electrolytes used in salt the bridge are:

  • The mobility of the cation and anion of the electrolyte should be the same.
  • The electrolyte of the salt bridge should not involve electrochemical change.
  • The ions of the salt bridge should not chemically react with the species of the cell.
Functions of salt bridge
  1. It completes the electrical circuit by allowing the migration of ions from one half-cell to another.
  2. It maintains the electrical neutrality of the electrolyte of the cathode and anode half-cell without mixing their solution.
Cell notation for Galvanic cell
  1. The anodic half-cell is always written on the left-hand side and the cathodic half-cell is on the right-hand side.
  2. A single vertical line is used to separate the phase equation.
\underset{Anodic\ half\ cell}{Zn(s)/Zn^{++}}\ \ \ \underset{Cathodic\ half\ cell}{Cu^{++}/Cu(s)}
  1. A double vertical line is used to donate the salt bridge.
Zn(s)/Zn^{++}//Cu^{++}/Cu(s)
  1. The concentration of the solution is shown at a side by enclosing the small bracket.
Zn(s)/Zn^{++}(aq.1M)\ // \ Cu^{++}(aq.\ 1M)/Cu(s)
  1. If an inert electrode like platinum, graphite, etc. are used, they are shown at the corresponding side of half cell by separating comma(,)
    If H2 is used as an anode, then
-Pt,H_{2}(g)(...atm)/H^{+}(aq.\ 1M)//.....

If H2 is used as a cathode, then

......//H^{+}(aq.\ 1M)/H_{2}(g)(...atm),Pt^{+}
Zinc-Copper (Zn-Cu) voltaic cell (Daniel cell)

It is constructed by dipping a zinc rod into a zinc sulphate solution and a copper rod into a copper sulphate solution followed by connecting these two half cells externally by a metallic conductor and internally with a salt bridge.
Eo Zn++/Zn = -0.76V and Eo Cu++/Cu = +0.34V.
Due to lower reduction potential, zinc act as an anode and copper act as a cathode.

zn-cu voltaic cell
Cell notation for Zn-Cu voltaic cell
-Zn(s)/Zn^{++}(aq.\ 1M)//Cu^{++}(aq.\ 1M)/Cu(s)+
Electrode reaction
\begin{align*}
At\ anode:\ Zn \rightarrow Zn^{++}&+ 2e^{-}\\
At\ cathode:\ Cu^{++}+2e^{-}&\rightarrow Cu\\
\hline
Cell\ reaction:\ Zn + Cu^{++}&\rightarrow Zn^{++}+Cu
\end{align*}
\begin{align*}
E^{o}\ Cell &= E^{o}\ Cathode - E^{o}\ Anode\\
&= 0.34 + 0.76 \\
&= 1.1 V
\end{align*}
Copper silver (Cu-Ag) voltaic cell

It is constructed by dipping a copper rod into copper nitrate solution and a silver rod into silver nitrate solution followed by connecting these two half cells externally by metallic conductor and internally with a salt bridge.
Eo Cu++/Cu = +0.34V and Eo Ag+/Ag = +0.80.
Due to lower reduction potential, copper act as an anode and silver acts as a cathode.

Cu-ag voltaic cell
Cell notation
-Cu(s)/Cu^{++}(aq.\ 1M)//Ag^{+}(aq.\ 1M)/Ag(s)+
Electrode reaction
\begin{align*}
At\ anode:\ Cu \rightarrow\ &Cu^{++}+ 2e^{-}\\
At\ cathode:\ Ag^{+}+e^{-}&\rightarrow Ag\ ] \times 2\\
\hline
Cell\ reaction:\ Cu + 2Ag^{+}&\rightarrow Cu^{++}+2Ag
\end{align*}
\begin{align*}
E^{o}\ Cell &= E^{o}\ Cathode - E^{o}\ Anode\\
&= 0.80 - 0.34 \\
&= 0.46 V
\end{align*}
Emf of a cell or cell potential

It is defined as the difference in potential between two half-cells or two electrodes of a galvanic cell.

E_{Cell} = E_{Cathode} - E_{Anode}

At standard temperature and pressure, the emf of cell or cell potential is known as standard emf or standard cell potential.

E^{o}_{Cell} = E^{o}_{Cathode} - E^{o}_{Anode}
Relation between cell potential and free energy

The work done by a galvanic cell equals the product of voltage (cell potential) and charge carried by an electron.

i.e. Work done by galvanic cell = Voltage x charge
= E x F x n
= nEF …….(I)

This work done of a galvanic cell is equal to the decrease in free energy of a system.
i.e. Work done by galvanic cell = decrease in free energy = – ΔG …….(ii)

From (i) and (ii):

ΔG = -nEF

where ΔG = free energy change
n = number of moles of electron
F = Faraday’s constant
E = emf of cell or cell potential

At STP, ΔGo = -nEoF

Commercial cell

Commercial cells are such electrochemical device which acts as an energy power source and provide considerably constant voltage over a period of time.

Battery: When two or more commercial cells are connected in series, the combination is called a battery. A battery is an arrangement of electrochemical cells used as a source of electrical energy.
A useful battery should have the following requirement:

  • It should be light and compact so that it can easily be transported.
  • It should have a long life when it is used and when not used.
  • The voltage of the battery should not vary appreciably during its use.
Types of commercial cell

There are mainly three types of commercial cells:

1. Primary cell: They are the commercial cell where electrical energy is obtained from chemical reactions as long as the active materials are present. They are not chargeable and cannot be reused. Thus the electrode reaction cannot be reversed. eg. dry cell, button cell

Dry cell (Leclanche cell)

It is the most familiar type of battery. It was invented by French engineer Leclanche in 1866. It is used in torches, toys, tape recorders, calculators, etc.

Construction of a dry cell

In a dry cell, the anode consists of a zinc container and the cathode is a graphite rod surrounded by powdered MnO2 and carbon. The space between the electrode is filled with a moist paste of NH4Cl and ZnCl2 acts as an electrolyte.
When the cell is working, the following reaction takes place:

dry cell reaction

NH3 liberated combines with ZnCl2 to form a diamine zinc(II) chloride complex which is poisonous.

2NH_{3} + ZnCl_{2} \rightarrow   (Zn(NH_{3})]Cl_{2}

If NH3 gas does not combine with ZnCl2, it would exert pressure and the seal of the cell would be broken. The dry cell does not have a long life since NH4Cl (acidic) corrodes the zinc container. The voltage of a dry cell varies from 1.25 – 1.50V.

dry cell electrochemistry

2. Secondary cells: They are also called storage cells or rechargeable cells or accumulator cells. The cell which can be recharged by passing direct current and their electrode reaction can be reversed are called secondary cells. They can be recharged, discharged and recharged many times. eg. Lead storage battery, Nickel-Cadmium cell, Lithium-ion cell, etc.

Lead storage battery

Lead storage cell or lead acid accumulator generally consists of six cells which are connected in series to get a 12V battery. Each cell produces 2V. Here, spongy lead acts as an anode and a grid of lead packed with PbO2 acts as a cathode. An aqueous solution of H2SO4 is used as an electrolyte.

Electrode reaction during discharge:

lead storage battery

Electrode reaction during recharging:

2PbSO4 +2H2O + Energy Pb + PbO2 + 4H+ + 2SO4

It is the most commonly used battery in automobiles.

charging and discharging
Lithium-ion battery

A lithium-ion battery or Li-ion battery is a type of rechargeable battery. They are commonly used for portable electronics and electric vehicles and are growing in popularity for military and aerospace applications.
In the batteries, lithium-ion move from the negative electrode through an electrolyte to the positive electrode during discharge and back when charging. They use an intercalated lithium compound as the material at the positive electrode and typically graphite at the negative electrode.

Electrode reaction:

\begin{align*}
At\ cathode:CoO_{2} + Li^{+} + e^{-}\rightarrow LiCoCO_{2}\\
At\ anode: LiC_{6} + CoO_{2}\rightarrow C_{6}+LiCoO_{2} \\
\hline
Net\ reaction:LiC_{6}+CoO_{2} \rightarrow C_{6}+ LiCoO_{2}
\end{align*}
Fuel cells (Hydrogen-oxygen fuel cells)

A fuel cell is an electrochemical cell which converts the chemical energy of a fuel directly into electrical energy. Hydrogen is used as a fuel cell or internal combustion engine.

fuel cell

A fuel cell consists of two metallic rods named an anode and cathode made from graphite coated with Pt or Ag as a catalyst. NaOH solution is taken as an electrolyte. Hydrogen and oxygen are supplied to the cell from anode and cathode compartments respectively under 50 atm and 250oC. Hydrogen is supplied towards the anode and reacts with OH- and forms water that supplies electrons to the circuit and conducts electricity. Oxygen is supplied towards the cathode which reacts with water to form hydroxide ions.

\begin{align*}
At\ anode: H_{2} + 2OH^{-}\rightarrow 2H_{2}O+2e^{-}\ ] \times 2 \\
At\ cathode:O_{2} + 2H_{2}O + 4e^{-}\rightarrow 4OH^{-}\\
\hline
Net\ reaction:2H_{2}+O_{2} \rightarrow 2H_{2}O+ Energy
\end{align*}
Applications of fuel cell
  1. It is used as fuel for vehicles like cars, buses, etc.
  2. It is used for spacecraft and military purposes.
  3. Water is obtained which is used for drinking purposes.
Some Important Questions

1. An electrochemical cell is constructed by using aluminium and silver electrodes whose electrode potentials are, Eo Al3+/Al= -1.67V, Eo Ag+/Ag=0.80.
i. Represent an electrochemical cell using the above electrode.
ii. Write down the complete cell reactions.
iii. Calculate the emf of the cell.

2. The standard electrode potential for Mg++/Mg and Cu++/Cu are -2.37V and +0.34V respectively.
i. Draw a standard cell notation.
ii. Identify the anode and cathode.
iii. Write the cell reactions taking place at electrodes.
iv. Calculate the standard cell potential.

3. The standard reduction potential for Fe3+/Fe2+ and I2/I are +0.77V and +0.54V respectively.
i. Draw a standard cell notation.
ii. Identify the anode and cathode.
iii. Write the cell reactions taking place at electrodes.
iv. Calculate the standard cell potential.

4. Predict which one of the following reaction occur spontaneously?
i. 2Fe++ + Sn+4 2Fe+++ + Sn+2
ii. 2Fe+++ + Sn2+ 2Fe++ + Sn4+
Given standard reduction potentials of Fe+++/Fe++ and Sn4+/2+ are +0.77V and 0.15V respectively.

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