2.8 Electrolytic Cells and Electrolysis | Electrochemistry | Chemistry 12

This section describes the principles of electrolytic cells, the process of electrolysis, and the factors that determine the products of an electrolytic reaction.

Electrochemical Cells

Electrochemical cells are devices that interconvert electrical and chemical energy. They are classified into two types:

Electrolytic Cells: These cells convert electrical energy into chemical energy by driving non-spontaneous redox reactions.
Galvanic (Voltaic) Cells: These cells convert chemical energy into electrical energy from spontaneous redox reactions.

For a detailed explanation of oxidation-reduction concepts, refer to Oxidation Reduction Concepts→.

Electrochemical cells diagram showing electrolytic and galvanic cell types — Figure 2.8.1: Classification of electrochemical cells.

Electrolytic Cells

An electrolytic cell is a device that uses an external electrical current to force a non-spontaneous redox reaction to occur. This forced chemical change is known as electrolysis.

Energy Conversion: Electrical Energy → Chemical Energy
Reaction Type: Non-spontaneous (requires external energy input)
Electrodes:
- Anode: The positive electrode where oxidation occurs. It is connected to the positive terminal of the power source.
- Cathode: The negative electrode where reduction occurs. It is connected to the negative terminal of the power source.

The electrochemical series determines the relative reactivity of species as oxidizing or reducing agents. See Electrode Electrode Potential And Electrochemical Series→ for more details.

Electrolytic cell setup showing electrode polarity and ion movement — Figure 2.8.2: Electrolytic cell with positive anode and negative cathode.

The Process of Electrolysis

Electrolysis is a chemical process initiated by passing a direct electric current through an electrolyte. An electrolyte is a substance in a molten state or aqueous solution that contains free ions.

Apparatus: An electrolytic cell typically consists of:
- An electrolyte (molten ionic compound or aqueous solution)
- Two inert electrodes (such as graphite or platinum)
- An external power source (such as a battery or DC power supply)
- Connecting wires to complete the circuit
Conduction Mechanisms:
- Metallic Conduction: Electrons carry the current through the external wires and electrodes.
- Electrolytic Conduction: Ions carry the current through the electrolyte. Cations move towards the cathode, and anions move towards the anode.
Mechanism: The power source pushes electrons into the cathode (making it negative) and pulls electrons from the anode (making it positive). This electron flow forces the non-spontaneous redox reaction to occur.

Units in Electrochemistry

Coulomb (C): The SI unit of electric charge. It represents the charge on approximately $6.25 \times 1 0^{18}$ electrons.

One Faraday (1 F) is equivalent to 96,487 Coulombs per mole. See Experimental Determination Of Avogadro Constant By Electrolytic Method→ for the relationship between these constants.

Faraday (F): A convenient unit for chemists, representing the charge of one mole of electrons.

$1 F = 96, 487 C m o l^{- 1}$

Ampere (A): The SI unit of electric current, defined as the flow of one coulomb per second.

$1 A = 1 C s^{- 1}$

Factors Influencing the Products of Electrolysis

The specific substances liberated at the electrodes during electrolysis depend on three critical factors:

1. The State of the Electrolyte:

Molten State: When an ionic compound is molten, only the cations and anions from the electrolyte itself are present and available to be discharged at the electrodes.
Aqueous State: When an ionic compound is dissolved in water, the solution contains ions from the electrolyte and ions ( $H^{+}$ and $O H^{-}$ ) formed from the autoionization of water, as well as water molecules ( $H_{2} O$ ). Therefore, water can also participate in the redox reactions.

2. Position in the Electrochemical (Redox) Series:

At the Cathode (Reduction): If multiple cations are present, the cation that is more easily reduced (that is, less reactive, lower in the electrochemical series) will be preferentially discharged. For example, $C u^{2 +}$ will be reduced to $Cu$ before $N a^{+}$ ions or water. If the metal cation is more reactive than hydrogen (such as $N a^{+}$ or $M g^{2 +}$ ), water will be reduced instead to produce $H_{2}$ gas.
At the Anode (Oxidation): If multiple anions or species capable of oxidation are present, the species that is more easily oxidized will be preferentially discharged.

The electrochemical series ranks species by their standard electrode potentials. See Electrochemical Series→ for detailed information.

3. Concentration of the Electrolyte:

The relative concentration of ions can significantly influence which species is discharged. A very high concentration of a particular ion can sometimes override the electrochemical series preference. This is particularly relevant when comparing the oxidation of halide ions ( $C l^{-}$ , $B r^{-}$ , $I^{-}$ ) with the oxidation of hydroxide ions ( $O H^{-}$ ) or water. A high concentration of halide ions often favors their discharge even if $O H^{-}$ or $H_{2} O$ are thermodynamically easier to oxidize under standard conditions.

Key Reactions and Formulae

The following examples illustrate how the factors above determine the products of electrolysis in different scenarios.

Electrolysis of Fused (Molten) Sodium Chloride (NaCl)

In molten NaCl, only $N a^{+}$ and $C l^{-}$ ions are present.

At Anode (Oxidation): Chloride ions lose electrons and are oxidized to chlorine gas. $2 C l^{-} (l) \to C l_{2} (g) + 2 e^{-}$
At Cathode (Reduction): Sodium ions gain electrons and are reduced to molten sodium metal. $2 N a^{+} (l) + 2 e^{-} \to 2 Na (l)$
Net Reaction: $2 N a^{+} (l) + 2 C l^{-} (l) \to 2 Na (l) + C l_{2} (g)$

Electrolysis of Concentrated Aqueous Sodium Chloride (Brine)

In concentrated aqueous NaCl, ions present are $N a^{+}$ , $C l^{-}$ , $H^{+}$ (from water dissociation), $O H^{-}$ (from water dissociation), and $H_{2} O$ molecules.

At Anode (Oxidation): Despite $O H^{-}$ and $H_{2} O$ being easier to oxidize under standard conditions, the very high concentration of $C l^{-}$ ions makes them preferentially oxidized to chlorine gas. $2 C l^{-} (a q) \to C l_{2} (g) + 2 e^{-}$
At Cathode (Reduction): $N a^{+}$ is a very reactive metal and is more difficult to reduce than water. Therefore, water molecules are reduced to hydrogen gas and hydroxide ions. $2 H_{2} O (l) + 2 e^{-} \to H_{2} (g) + 2 O H^{-} (a q)$
Net Reaction: The overall reaction produces hydrogen gas, chlorine gas, and sodium hydroxide in solution. $2 H_{2} O (l) + 2 C l^{-} (a q) \to H_{2} (g) + C l_{2} (g) + 2 O H^{-} (a q)$

Electrolysis of Dilute Aqueous Sodium Chloride

In dilute aqueous NaCl, the ions present are $N a^{+}$ , $C l^{-}$ , $H^{+}$ , $O H^{-}$ , and $H_{2} O$ molecules, but the concentration of $C l^{-}$ is low.

At Anode (Oxidation): With a low concentration of $C l^{-}$ , the oxidation of $C l^{-}$ is less favorable. Instead, hydroxide ions (derived from water) are oxidized, producing oxygen gas and water. $4 O H^{-} (a q) \to O_{2} (g) + 2 H_{2} O (l) + 4 e^{-}$
At Cathode (Reduction): Similar to concentrated NaCl, $N a^{+}$ is still harder to reduce than water. Water is reduced to hydrogen gas and hydroxide ions. $4 H_{2} O (l) + 4 e^{-} \to 2 H_{2} (g) + 4 O H^{-} (a q)$
Net Reaction: The overall reaction is effectively the electrolysis of water. $2 H_{2} O (l) \to 2 H_{2} (g) + O_{2} (g)$

Galvanic (Voltaic) Cells

Unlike electrolytic cells, galvanic cells utilize spontaneous chemical reactions to produce electricity.

Summary

Electrolytic cells are devices that utilize external electrical energy to drive non-spontaneous chemical reactions.
During electrolysis, oxidation occurs at the positive anode, and reduction occurs at the negative cathode.
The identity of the products formed at the electrodes is determined by:
1. The physical state of the electrolyte.
2. The relative positions of the ions in the electrochemical series.
3. The concentration of the ions.

Electrolyte	Product at Anode (+)	Product at Cathode (-)
Molten NaCl	Chlorine gas ( $C l_{2}$ )	Molten Sodium ( $Na$ )
Conc. Aqueous NaCl	Chlorine gas ( $C l_{2}$ )	Hydrogen gas ( $H_{2}$ )
Dilute Aqueous NaCl	Oxygen gas ( $O_{2}$ )	Hydrogen gas ( $H_{2}$ )

Chemical Significance: Electrolysis is used for producing sodium, magnesium, and aluminum, and for manufacturing chlorine gas and sodium hydroxide.

This section describes the principles of electrolytic cells, the process of electrolysis, and the factors that determine the products of an electrolytic reaction.

Electrochemical Cells

Electrochemical cells are devices that interconvert electrical and chemical energy. They are classified into two types:

Electrolytic Cells: These cells convert electrical energy into chemical energy by driving non-spontaneous redox reactions.
Galvanic (Voltaic) Cells: These cells convert chemical energy into electrical energy from spontaneous redox reactions.

For a detailed explanation of oxidation-reduction concepts, refer to Oxidation Reduction Concepts→.

Electrolytic Cells

An electrolytic cell is a device that uses an external electrical current to force a non-spontaneous redox reaction to occur. This forced chemical change is known as electrolysis.

Energy Conversion: Electrical Energy → Chemical Energy
Reaction Type: Non-spontaneous (requires external energy input)
Electrodes:
- Anode: The positive electrode where oxidation occurs. It is connected to the positive terminal of the power source.
- Cathode: The negative electrode where reduction occurs. It is connected to the negative terminal of the power source.

The electrochemical series determines the relative reactivity of species as oxidizing or reducing agents. See Electrode Electrode Potential And Electrochemical Series→ for more details.

The Process of Electrolysis

Electrolysis is a chemical process initiated by passing a direct electric current through an electrolyte. An electrolyte is a substance in a molten state or aqueous solution that contains free ions.

Apparatus: An electrolytic cell typically consists of:
- An electrolyte (molten ionic compound or aqueous solution)
- Two inert electrodes (such as graphite or platinum)
- An external power source (such as a battery or DC power supply)
- Connecting wires to complete the circuit
Conduction Mechanisms:
- Metallic Conduction: Electrons carry the current through the external wires and electrodes.
- Electrolytic Conduction: Ions carry the current through the electrolyte. Cations move towards the cathode, and anions move towards the anode.
Mechanism: The power source pushes electrons into the cathode (making it negative) and pulls electrons from the anode (making it positive). This electron flow forces the non-spontaneous redox reaction to occur.

Units in Electrochemistry

Coulomb (C): The SI unit of electric charge. It represents the charge on approximately $6.25 \times 1 0^{18}$ electrons.

One Faraday (1 F) is equivalent to 96,487 Coulombs per mole. See Experimental Determination Of Avogadro Constant By Electrolytic Method→ for the relationship between these constants.

Faraday (F): A convenient unit for chemists, representing the charge of one mole of electrons.

$1 F = 96, 487 C m o l^{- 1}$

Ampere (A): The SI unit of electric current, defined as the flow of one coulomb per second.

$1 A = 1 C s^{- 1}$

Factors Influencing the Products of Electrolysis

The specific substances liberated at the electrodes during electrolysis depend on three critical factors:

1. The State of the Electrolyte:

Molten State: When an ionic compound is molten, only the cations and anions from the electrolyte itself are present and available to be discharged at the electrodes.
Aqueous State: When an ionic compound is dissolved in water, the solution contains ions from the electrolyte and ions ( $H^{+}$ and $O H^{-}$ ) formed from the autoionization of water, as well as water molecules ( $H_{2} O$ ). Therefore, water can also participate in the redox reactions.

2. Position in the Electrochemical (Redox) Series:

At the Cathode (Reduction): If multiple cations are present, the cation that is more easily reduced (that is, less reactive, lower in the electrochemical series) will be preferentially discharged. For example, $C u^{2 +}$ will be reduced to $Cu$ before $N a^{+}$ ions or water. If the metal cation is more reactive than hydrogen (such as $N a^{+}$ or $M g^{2 +}$ ), water will be reduced instead to produce $H_{2}$ gas.
At the Anode (Oxidation): If multiple anions or species capable of oxidation are present, the species that is more easily oxidized will be preferentially discharged.

The electrochemical series ranks species by their standard electrode potentials. See Electrochemical Series→ for detailed information.

3. Concentration of the Electrolyte:

The relative concentration of ions can significantly influence which species is discharged. A very high concentration of a particular ion can sometimes override the electrochemical series preference. This is particularly relevant when comparing the oxidation of halide ions ( $C l^{-}$ , $B r^{-}$ , $I^{-}$ ) with the oxidation of hydroxide ions ( $O H^{-}$ ) or water. A high concentration of halide ions often favors their discharge even if $O H^{-}$ or $H_{2} O$ are thermodynamically easier to oxidize under standard conditions.

Key Reactions and Formulae

The following examples illustrate how the factors above determine the products of electrolysis in different scenarios.

Electrolysis of Fused (Molten) Sodium Chloride (NaCl)

In molten NaCl, only $N a^{+}$ and $C l^{-}$ ions are present.

At Anode (Oxidation): Chloride ions lose electrons and are oxidized to chlorine gas. $2 C l^{-} (l) \to C l_{2} (g) + 2 e^{-}$
At Cathode (Reduction): Sodium ions gain electrons and are reduced to molten sodium metal. $2 N a^{+} (l) + 2 e^{-} \to 2 Na (l)$
Net Reaction: $2 N a^{+} (l) + 2 C l^{-} (l) \to 2 Na (l) + C l_{2} (g)$

Electrolysis of Concentrated Aqueous Sodium Chloride (Brine)

In concentrated aqueous NaCl, ions present are $N a^{+}$ , $C l^{-}$ , $H^{+}$ (from water dissociation), $O H^{-}$ (from water dissociation), and $H_{2} O$ molecules.

At Anode (Oxidation): Despite $O H^{-}$ and $H_{2} O$ being easier to oxidize under standard conditions, the very high concentration of $C l^{-}$ ions makes them preferentially oxidized to chlorine gas. $2 C l^{-} (a q) \to C l_{2} (g) + 2 e^{-}$
At Cathode (Reduction): $N a^{+}$ is a very reactive metal and is more difficult to reduce than water. Therefore, water molecules are reduced to hydrogen gas and hydroxide ions. $2 H_{2} O (l) + 2 e^{-} \to H_{2} (g) + 2 O H^{-} (a q)$
Net Reaction: The overall reaction produces hydrogen gas, chlorine gas, and sodium hydroxide in solution. $2 H_{2} O (l) + 2 C l^{-} (a q) \to H_{2} (g) + C l_{2} (g) + 2 O H^{-} (a q)$

Electrolysis of Dilute Aqueous Sodium Chloride

In dilute aqueous NaCl, the ions present are $N a^{+}$ , $C l^{-}$ , $H^{+}$ , $O H^{-}$ , and $H_{2} O$ molecules, but the concentration of $C l^{-}$ is low.

At Anode (Oxidation): With a low concentration of $C l^{-}$ , the oxidation of $C l^{-}$ is less favorable. Instead, hydroxide ions (derived from water) are oxidized, producing oxygen gas and water. $4 O H^{-} (a q) \to O_{2} (g) + 2 H_{2} O (l) + 4 e^{-}$
At Cathode (Reduction): Similar to concentrated NaCl, $N a^{+}$ is still harder to reduce than water. Water is reduced to hydrogen gas and hydroxide ions. $4 H_{2} O (l) + 4 e^{-} \to 2 H_{2} (g) + 4 O H^{-} (a q)$
Net Reaction: The overall reaction is effectively the electrolysis of water. $2 H_{2} O (l) \to 2 H_{2} (g) + O_{2} (g)$

Galvanic (Voltaic) Cells

Unlike electrolytic cells, galvanic cells utilize spontaneous chemical reactions to produce electricity.

Summary

Electrolytic cells are devices that utilize external electrical energy to drive non-spontaneous chemical reactions.
During electrolysis, oxidation occurs at the positive anode, and reduction occurs at the negative cathode.
The identity of the products formed at the electrodes is determined by:
1. The physical state of the electrolyte.
2. The relative positions of the ions in the electrochemical series.
3. The concentration of the ions.

Electrolyte	Product at Anode (+)	Product at Cathode (-)
Molten NaCl	Chlorine gas ( $C l_{2}$ )	Molten Sodium ( $Na$ )
Conc. Aqueous NaCl	Chlorine gas ( $C l_{2}$ )	Hydrogen gas ( $H_{2}$ )
Dilute Aqueous NaCl	Oxygen gas ( $O_{2}$ )	Hydrogen gas ( $H_{2}$ )

Chemical Significance: Electrolysis is used for producing sodium, magnesium, and aluminum, and for manufacturing chlorine gas and sodium hydroxide.