NJC 2026 Electrochemistry Part 1 (Student)

National Junior College SH2 H2 Chemistry 1 Electrochemistry Electrochemistry is the study of the inter-conversion of chemical and electrical energies. This inter-conversion of energies occurs in cells in which redox reactions (i.e. reactions involving the transfer of electrons from one chemical species to another) produce or utilise electrical energy. Comparison of Electrochemical (Part 1) and Electrolytic cell (Part 2) ELECTROCHEMICAL CELL ELECTROLYTIC CELL Energy change chemical energy → electrical energy electrical energy → chemical energy Cell Reaction Spontaneous Chemical reaction takes place to generate electric current within celll (battery). Non-spontaneous External electric current is applied to drive the reaction to occur. Cathode (reduction) Positive (+) electrode Negative (–) electrode Anode (oxidation) Negative (–) electrode Positive (+) electrode Note: 1. In both electrochemical and electrolytic cells, oxidation always takes place at the anode while reduction always takes place at the cathode. 2. Polarity of electrode always follow the polarity of the battery. [ElectroCHEMICAL cell] The two half-cells form the battery. Anode is the negative terminal of the battery that releases electrons while cathode is the positive terminal of the battery that accepts electrons. [ElectroLYTIC cell] The polarities of the electrode are reversed, as the electrodes follows the polarity of the battery terminal it is connected to. • anode (+ve) is connected to the +ve terminal of battery and attracts anions (−ve) • cathode (−ve) is connected to the −ve terminal of battery and attracts cations (+ve)

National Junior College SH2 H2 Chemistry 2 Electrochemical Cell (Part 1) Content • Redox processes: electron transfer and changes in oxidation number (oxidation state) • Electrode potentials (i) standard electrode (redox) potentials, Eo; the redox series (ii) standard cell potentials, Eo , and their uses (iii) batteries and fuel cells Learning Outcomes Candidates should be able to: (a) describe and explain redox processes in terms of electron transfer and/or of changes in oxidation number (oxidation state) (b) define the terms: (i) standard electrode (redox) potential (ii) standard cell potential (c) describe the standard hydrogen electrode (d) describe methods used to measure the standard electrode potentials of: (i) metals or non-metals in contact with their ions in aqueous solution (ii) ions of the same element in different oxidation states (e) calculate a standard cell potential by combining two standard electrode potentials (f) use standard cell potentials to: (i) explain/deduce the direction of electron flow from a simple cell (ii) predict the spontaneity of a reaction (g) understand the limitations in the use of standard cell potentials to predict the spontaneity of a reaction (h) construct redox equations using the relevant half-equations (see also Section 13) (i) state and apply the relationship ∆Go = −nFEo to electrochemical cells, including the calculation of Eo for combined half reactions (j) predict qualitatively how the value of an electrode potential varies with the concentration of the aqueous ion (k) state the possible advantages of developing other types of cell, e.g. the H2/O2 fuel cell and improved batteries (as in electric vehicles) in terms of smaller size, lower mass and higher voltage References ● Chemistry; The Molecular Nature of Matter and Change, by Martin Silberberg, 2010, Publisher: McGraw Hill ● Chemistry for Advanced Level, by Peter Cann and Peter Hughes, 2002, Publisher: John Murray ● Chemguide: https://www.chemguide.co.uk/physical/redoxeqiamenu.html#top Copyright © National Junior College All Rights Reserved. No part of this publication may be reproduced or transmitted in any form or by any means, electronic or mechanical, including photocopy, recording or any other information storage and retrieval system, without prior permission in writing from the copyright owner.

National Junior College SH2 H2 Chemistry 3 Prerequisite to learn and understand this topic well An understanding of Redox concepts, Moles and Stoichiometry and Le Chatelier’s principle are imperative to do well for this chapter. 1 Introduction Electrochemistry is the study of the inter-conversion of chemical and electrical energies. Success Criteria • I can differentiate between an electrochemical cell and electrolytic cell. This inter-conversion of energies occurs in cells in which redox reactions (i.e. reactions involving the transfer of electrons from one chemical species to another) produce or utilise electrical energy. There are two types of cells. Electrochemical cell (also called galvanic cell or battery) Electrolytic cell Generation of Electricity Use of Electricity for Work Redox reaction is spontaneous (∆G < 0) Redox reaction is non-spontaneous (∆G > 0) Chemical → Electrical energy Electrical → Chemical energy Examples: dry cells, batteries, fuel cells Examples: extraction of aluminium metal from molten aluminium oxide, purification of copper, electroplating Electrochemical Cell (Battery) (Battery) Electrolytic Cell

National Junior College SH2 H2 Chemistry 4 1.1 About Redox Reactions Success Criteria ● I know the difference between reduction and oxidation in terms of electron transfer and oxidation numbers Oxidation: A process where a chemical species loses electrons (increase in oxidation number) Reduction: A process where a chemical species gains electrons (decrease in oxidation number). A redox reaction is a process involving electron transfer: simultaneous loss of electrons from the species undergoing oxidation and gain of electrons by the species undergoing reduction. Video here: Worked Example 1 (a) Describe and explain what is observed when a piece of copper metal is placed in a solution of silver nitrate. Description: ● As the reaction proceeds, the colourless solution turns blue and the pink copper pipe is covered with silver metal, which falls off and collects under the pipe. Explanation: ● Cu(s) lose electrons and get oxidised to Cu2+(aq) ions ● Ag+(aq) ions gain electrons and get reduced to Ag(s) ● A redox reaction has taken place. (b) Write appropriate redox equations to describe the reactions taking place in the beaker. [R]: Ag+(aq) + e− ⎯→ Ag(s) [O]: Cu(s) ⎯→ Cu2+(aq) + 2e− Overall: Cu(s) + 2Ag+(aq) ⎯→ Cu2+(aq) + 2Ag(s) (c) What would you expect to observe if a piece of Ag metal is placed in a solution of Cu2+? Observations: Since the reaction of Ag +(aq) and Cu(s) is spontaneous, the reverse reaction is expected to be non-spontaneous. Hence, the intensity of the blue Cu2+(aq) remains the same while Ag metal remains with nothing deposited on it. Cu(s) AgNO3 (aq)

National Junior College SH2 H2 Chemistry 5 2 The Electrochemical Cell Success Criteria: ● I can identify the polarity of the electrode based on the reaction occurring. ● I can predict the products formed at each electrode for the given example and understand how the observations occur for each half-cell. In Worked Example 1, Cu(s) and Ag+(aq) are in direct physical contact, thus the electrons are transferred directly from Cu to Ag+. The energy released from the reaction is given off as heat, and no usable energy is harnessed from this spontaneous redox reaction. To capture this energy in the form of electrical energy, an electrochemical cell can be set up below by separating the two half-reactions above. Note: A half-cell consists of both an electrode and an electrolyte(s). A half-cell contains 2 species of the sa