JC H2 Chemistry Electrochemistry: From Galvanic Cells to Electrolysis

Electrochemistry is a topic that combines chemical knowledge with electrical concepts, making it challenging for students who struggle with either chemistry or physics. At Sophia Education, our JC Chemistry tutors have developed a systematic approach that makes electrochemistry accessible and logical.

Galvanic Cells: Spontaneous Redox Reactions

A galvanic cell converts chemical energy to electrical energy through a spontaneous redox reaction. The anode is where oxidation occurs (negative terminal in a galvanic cell). The cathode is where reduction occurs (positive terminal). Electrons flow from anode to cathode through the external circuit. Ions flow through the salt bridge to maintain electrical neutrality.

Standard Electrode Potentials (E°)

E°cell = E°cathode - E°anode. A positive E°cell indicates a spontaneous reaction. The more positive the E° value, the stronger the oxidising agent. The more negative the E° value, the stronger the reducing agent. Use the electrochemical series to predict whether a reaction is spontaneous.

Electrolytic Cells: Non-Spontaneous Reactions

An electrolytic cell uses electrical energy to drive a non-spontaneous redox reaction. The anode is connected to the positive terminal of the power supply (oxidation occurs). The cathode is connected to the negative terminal (reduction occurs). For aqueous solutions, consider whether water or the dissolved ions are preferentially discharged.

Faraday's Laws of Electrolysis

First Law: The mass of substance deposited is proportional to the quantity of charge passed. Second Law: The masses of different substances deposited by the same quantity of charge are proportional to their molar masses divided by the number of electrons transferred. Formula: m = (Q × M) / (n × F), where F = 96500 C mol⁻¹.