Introduction to Electrochemistry
“The branch of chemistry which deals with interconversion of electrical and chemical energy is called electrochemistry”.
Importance of Electrochemistry
Electrochemistry is fundamental to many real-world applications. Photosynthesis is a redox reaction which provides food for the entire planet, and respiration that keeps you alive, both are redox reactions.
Electrochemical cells power our devices, electroplating protects metals from corrosion, and electrolysis produces important industrial chemicals.
Oxidation, Reduction, and Redox Reactions
Oxidation Process
Oxidation is a process involving loss of electron or electrons.
Example: Fe²⁺ → Fe³⁺ + e⁻
Here Fe²⁺ has lost an electron, therefore it is oxidized.
Reduction Process
Reduction is a process involving gain of electron or electrons.
Example: Cl⁰ + e⁻ → Cl⁻
Here Cl⁰ has gained an electron, therefore it is reduced.
Redox Reactions
“Oxidation and reduction always take place together. The reactions in which this happens is called redox reactions.”
There are two ways of finding out whether or not a substance has been oxidized or reduced during a chemical reaction: electron transfer and changes in oxidation number.
Oxidation Number and Its Significance
Definition
“An oxidation number is a number given to each atom or ion in a compound that shows its degree of oxidation”.
Or “It is the apparent charge on an atom (per atom) of an element in a molecule or an ion”.
Oxidation Number Rules
1. The oxidation number of any uncombined element is zero.
2. In either a compound or an ion, the more electronegative element is given the negative oxidation number.
3. In compounds, many atoms or ions have fixed oxidation numbers:
Group 1 (alkali metals) have always +1 oxidation number.
Group 2 (alkaline earth metals) have always +2 oxidation number.
Group 17 (halogens) in binary compounds have always -1 oxidation number.
Hydrogen is +1 (except in metal hydrides, such as NaH, where it is -1).
Oxygen is -2 (except in peroxides, where it is -1, and in OF₂, where it is +2).
4. The oxidation number of an element in a mono-atomic ion is always the same as the charge on the ion.
5. The sum of the oxidation numbers in a compound is zero.
6. In ions, the algebraic sum of oxidation number is equal to charge on the ion.
Disproportionation Reaction
Definition
“A disproportionation reaction is a chemical reaction where a single substance acts as both the oxidizing and reducing agent, resulting in two different products with different oxidation states”.
This type of reaction involves the simultaneous oxidation and reduction of the same element.
Example
Decomposition of hydrogen peroxide (H₂O₂) to form water (H₂O) and oxygen (O₂) is a disproportionation reaction:
2H₂O₂ → 2H₂O + O₂
Oxidation number of oxygens in H₂O₂ = -1
Oxidation number of oxygens in H₂O = -2 (Oxygen is reduced from -1 to -2)
Oxidation number of oxygens in O₂ = 0 (Oxygen is oxidized from -1 to 0)
Oxidizing and Reducing Agents
| Oxidizing Agent | Reducing Agent |
|---|---|
| An oxidizing agent (atom, ion or molecule) is that substance which oxidizes some other substance, and is itself reduced to a lower oxidation state by gaining one or more electrons. | A reducing agent (atom, ion or molecule) is that substance which reduces some other substance, and is itself oxidized to a higher oxidation state by losing one or more electrons. |
| Its oxidation state decreases due to gain of electrons. | Its oxidation state increases due to loss of electrons. |
| Non-metals are good oxidizing agents. | Metals are good reducing agents. |
Electrolytic Cell
Definition
“An electrolytic cell is a device that converts electrical energy into chemical energy through a process called electrolysis”.
Construction and Working
It consists of two electrodes, typically made of metal or another conductive material.
These electrodes are immersed in an electrolyte a substance containing free ions that carry electric current.
When a direct electric current is applied, one electrode becomes negatively charged (the cathode) and the other positively charged (the anode).
Positive ions in the electrolyte migrate towards the cathode, where they gain electrons (a reduction process), while negative ions move towards the anode, where they lose electrons (an oxidation process).
Uses of Electrolytic Cell
Electrolytic cells are used in:
The electrolysis of sodium chloride to produce sodium metal and chlorine gas
The refining and plating of metals
The production of chemicals like caustic soda.
Electrode Potentials
Definition
“When a metal is put into a solution of its ions, an electric potential (voltage) is established between the metal and the metal ions in solution”.
This is called electrode potential and it indicates the ease of oxidation or reduction of a substance.
Standard Hydrogen Electrode (SHE)
The standard hydrogen electrode is one of several types of half-cell that can be used as reference electrode.
This electrode consists of: Hydrogen gas at 101 kPa (1atm) pressure, in equilibrium with H⁺ ions of concentration 1.00 mol dm⁻³ (1M HCl) and a platinum electrode covered with platinum black in contact with the hydrogen gas and the H⁺ ions.
Standard Electrode Potential
“The standard electrode potential for a half-cell is the voltage measured under standard conditions with a standard hydrogen electrode (SHE) as the other half-cell”.
Standard conditions: Concentration of ions (1.00 mol dm⁻³), Temperature (25 °C/298 K), Pressure of gas (1atm/101 kPa).
Nernst Equation
Equation
For a given electrode, e.g., a Cu(s) /Cu²⁺ electrode, the relationship is:
E = E° + (RT/zF) ln([oxidized form]/[reduced form])
Where E is the electrode potential under non-standard conditions, E° is the standard electrode potential, R is the gas constant (8.314 J K⁻¹ mol⁻¹), T is the Kelvin temperature, z is the number of electrons transferred, F is the Faraday constant, ln is the natural logarithm.
Simplified Form
At standard temperature (25°C): E = E° + (0.059/z) log([oxidized form]/[reduced form])