Electrochemistry

Galvanic Cells or Voltaic Cells

A review of the redox reaction can be found here.

• Based on spontaneous reaction
• Consists of two half-cells
• Oxidation occurs at the anode
• Reduction occurs at the cathode
• Converts chemical potential energy to electrical energy
• The strongest oxidizing agent always undergoes reduction at the cathode.

A galvanic cell or voltaic cell consists of the following components, as shown in the figure below:

• Two electrodes - Solid conductors (anode and cathode); usually 2 metals. Oxidation and reduction reaction happen at electrode. 
• Electrolytes – These are the ions in a neutral solution which help conduct electricity via ionic conduction.
• External circuit – Usually a wire that allows for the movement of electrons from anode to cathode.
• Salt bridge - Contains an electrolytic solution that doesn’t interfere with the reaction.
• Internal circuit – Salt bridge or porous barrier that allows movement of positive and negative ions in solution.
• Half-cells- The redox reaction happens at each half cell: oxidation at one half-cell, and reduction at the other. 

Figure 1. Galvanic cells and electrodes. From LibreTexts (2021, March 3). Chemistry. https://chem.libretexts.org/@go/page/260

Cell Notation

Cell notations are short hand notation of galvanic (or voltaic) cells that describe the following:

• Reaction conditions (temperature, pressure and concentration),
• The anode and the cathode,
• Electrode components.

Cell Notation Rules

• The anode half-cell reaction (where the oxidation happens) is given first, followed by the cathode half-cell reaction.
• Reactants are written first, followed by the products for a given half-cell.
• The spectator ions (the ions that neither lose nor gain electrons) are not included in the cell notation.
• A single vertical line is used between two chemical species in different phases that are in physical contact, such as a solid electrode and a liquid electrolyte.  
• A double vertical line is used to represent a salt bridge or porous membrane that separates the half-cells.
• The phase of chemicals involved are shown in brackets.
• If the reaction conditions are different than the standard temperature and pressure, then they are also shown in brackets. If nothing is shown, then the reaction is assumed to happen at standard conditions.

Example:

Figure 2. Voltaic Cells. From LibreTexts (2020, August 15). Chemistry. https://chem.libretexts.org/@go/page/285

In the above example:

The reduction half-reaction for the above cell notation is given below. Since Ag⁺ is gaining electrons, it is being reduced, and this reduction happens at cathode. Since Ag⁺ is gaining electrons to become Ag, the reactant is Ag⁺ and it is listed first in the cell notation. The product is Ag and it is listed second for the same half-reaction. 

Ag⁺_(aq)+ e⁻ ⇌ Ag_(s)

The oxidation half-reaction for the above cell notation is given below. Since Cu is losing electrons, it is being oxidized, which happens at anode.  Cu is losing electrons to become Cu²⁺. Therefore, the reactant is Cu and it is listed first in the cell notation. The product is Cu²⁺ and it is listed second for the same half-reaction. 

Cu_(s)⇌Cu²⁺_(aq)+2e⁻

Inert and Active Electrodes

An inert electrode is a non-reactive electrode whose primary purpose is to transport electrons. It does not exchange electrons with the solution.

A reactive electrode (such as copper in the above example) is something that participates in the reaction. In the example above, copper is being oxidized to produce Cu²⁺.

Examples of commonly used inert electrodes: platinum, rhodium, carbon, and gold.

Examples of commonly used active electrodes: copper, sliver, zinc and lead.

Cell Potentials

Cell potential is the potential difference that exists between two half-cells. It is calculated using the following formula:

E^o_Cell=E^o_Cathode+ E^o_Anode

where E^o is the standard reduction potential found in your data booklet, as shown below:

Calculating the cell potential for the reaction in the previous section would go as follows:

Cu_(s) → Cu²⁺_(aq) + 2e^-     E^o_cathode= -0.340 V

Since copper is being oxidized, we use the oxidation half-cell of the redox reaction. Therefore, we reverse the sign of E^o

(Ag⁺ + e^- → Ag_(s) ) x2       E^o_Anode = +0.800 V

This reduction half-reaction is multiplied by two, balancing the number of electrons between both half-reactions.

E^o_Cell = E^o_Cathode +  E^o_Anode

E^o_Cell = 0.800 V + (-0.340 V)

E^o_Cell = +0.460V

Note: Since the cell potential is positive, it is a spontaneous redox reaction!

Electrolytic Cells

Voltaic cells are driven by a spontaneous redox reaction which converts chemical energy into electrical energy. 

Electrolytic cells, on the other hand, are driven by a non-spontaneous reaction that converts electrical energy into chemical energy. The redox reaction is driven by electrical energy supplied by an external energy source, as shown in the picture below:

Figure 3: The electrolysis of molten sodium chloride. From Openstax (2021) "Electrolysis." Chemistry 2e. https://openstax.org/books/chemistry-2e/pages/17-7-electrolysis

The half-reactions for the figure above are as follows:

Oxidation half reaction : 2Cl^-_(l)  →   Cl_2(g) + 2e^-

Reduction half reaction : Na⁺_(l) + e-  →  Na_(l)

The table below shows the comparison of voltaic cells and electrolytic cells: