ASRJC 2025 H2 Chem Electrochemistry Notes

2025 JC2 H2 Electrochemistry ©2025/ASRJC/CHEM 0 ANDERSON SERANGOON JUNIOR COLLEGE JC2 H2 CHEMISTRY ELECTROCHEMISTRY Content I. Overview II. Electrode Potentials & Electrochemical Cells III. Electrolysis & Electrolytic Cells IV. Summary 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 (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 H 2/O2 fuel cell and improved batteries (as in electric vehicles) in terms of smaller size, lower mass and higher voltage (l) state the relationship, F = Le, between the Faraday constant, the Avogadro constant and the charge on the electron (m) predict the identity of the substance liberated during electrolysis from the state of electrolyte (molten or aqueous), position in the redox series (electrode potential) and concentration (n) calculate: (i) the quantity of charge passed during electrolysis (ii) the mass and/or volume of substance liberated during electrolysis (o) explain, in terms of the electrode reactions, the industrial processes of: (i) the anodising of aluminium (ii) the electrolytic purification of copper [technical details are not required] References 1. Chemistry for Advanced Level, Cann and Hughes, Murray 2. Understanding Advanced Physical Inorganic Chemistry, Jeanne Tan and Kim Seng Chan 3. Chemistry, The Molecular Nature of Matter and Change (Fourth Edition), Silberberg, McGraw Hill 4. Chemistry The Central Science (Ninth Edition), Brown, LeMay, Bursten, Prentice Hall

2025 JC2 H2 Electrochemistry ©2025/ASRJC/CHEM 1 I. Overview Electrochemistry is the study of interactions between chemical changes and flow of electrons. There are two types of cells: Note that in electrochemistry, we are dealing with redox reactions. Electrochemistry Electrochemical cell Electrolytic cell Part 1 Electrochemical cells (or Galvanic / Voltaic cells) Example: Alkaline battery chemical reaction flow of electrons  converts chemical energy from a spontaneous reaction to generate electrical energy Part 2 Electrolytic cells Example: Down Cell (Electrolysis of molten NaCl) Dow flow of electrons chemical reaction  converts electrical energy to cause a non–spontaneous reaction to occur  What are the factors affecting the discharge of a substance during electrolysis?  How can we calculate the amount of substance liberated during electrolysis?  What are some industrial applications of electrolysis?  What happens in a redox reaction?  How can the tendency of a substance to be reduced or oxidised be measured?  What are the factors affecting electrode potential?  How can the direction and feasibility of a redox reaction be predicted, and what are the limitations?

2025 JC2 H2 Electrochemistry ©2025/ASRJC/CHEM 2 metal rod, M(s) Mn+(aq) II Electrode Potentials & Electrochemical Cells The tendency of a substance to be oxidised or reduced can be measured by its electrode potential. (A) Electrode Potentials When a metal rod (called the electrode) is dipped into a solution containing ions , Mn+, of the same metal (M), two possible electron transfer reactions can occur: oxidation and reduction (i) Oxidation: The M atoms have a tendency to lose electrons. Some atoms will go into the solution as Mn+. The electrons will be left behind on the metal. As time passes, there will be a build–up of electrons on the M surface. M(s) Mn+(aq) + ne– (ii) Reduction: The Mn+ ions on the metal surface may gain electrons, resulting in the regeneration of M. Mn+(aq) + ne– M(s) When the rates of oxidation and reduction are equal, dynamic equilibrium is established. At this point, there will be a constant negative charge on the metal, and a constant number of ions present in the solution around the metal. Mn+(aq) + ne– M(s) For metals that are more reactive (i.e. more readily oxidised) e.g. Zn Position of equilibrium lies more to the left , meaning that Zn tends to dissolve and go into the solution as positively charged Zn2+ ions. Zn2+(aq) + 2e Zn(s) This leaves a surplus of electrons on the metal rod and thus a negative charge on the Zn electrode. The negatively charged electrode attracts a layer of positive ions on the metal surface. This results in a separation of charges across the metal –solution interface, creating a potential difference between the metal and the solution . This potential difference is known as the absolute electrode potential, E of the metal. For metals that are less reactive (i.e. more readily reduced) e.g Cu Position of equilibrium lies more to the right , meaning that the M n+ ions in the solution tend to gain electrons from the metal rod and are thus deposited as metal atoms on the electrode. This results in a net deficit of electrons on the copper rod and thus a positive charge on the electrode. Cu2+(aq) + 2e Cu(s) oxidation reduction copper electrode Cu2+(aq) from CuSO4(aq) + + + + represents Cu Zn rod (electrode) Zn2+(aq) from ZnSO4(aq) + represents Zn2+ – – – – – –

2025 JC2 H2 Electrochemistry ©2025/ASRJC/CHEM 3 (B) Half–Cells An electrochemical cell consists of two half-cells. Half–cells are systems in which the oxidised state is in equilibrium with the reduced state. oxidised form + ne– reduced form There are three main types of half–cells: 1 Metal and its ion Half–cell e.g. Zn2+(aq)Zn(s) half–cell Zn2+(aq) + 2e– Zn(s) The solid Zn metal acts as the electrode when dipped in a solution containing its ions. 2 Non–metal and its ion Half–cell e.g. Cl2(g)Cl–(aq) half–cell ½Cl2(g) + e– Cl–(aq) An inert electrode ( e.g. platinum electrode) dipped in a solution containing C l– ions will develop an electrode potential when gaseous Cl2 is bubbled into the solution. 3 Ion–Ion Half–cell e.g. Fe3+(aq), Fe2+ (aq) half–cell Fe3+(aq) + e– Fe2+(aq) It consists of an inert electrode (e.g. Pt) dipped into a solution containing ions of the