RI 2025 Electrochemistry 1 Notes

1 RAFFLES INSTITUTION YEAR 6 H2 CHEMISTRY 2025 Lecture Notes 21a – Electrochemistry 1 _______________________________________________________________________________ A Content  Redox processes: electron transfer and changes in oxidation number (oxidation state)  Electrode potentials (i) Standard electrode (redox) potential, E ; the redox series (ii) Standard cell potentials, Ecell, and their uses (iii) Batteries and fuel cells B 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 G = – nFE to electrochemical cells, including the calculation of E 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 C References 1. Chemistry for Advanced Level by Peter Cann and Peter Hughes 2. Chemistry the Central Science by T L Brown and H E Lemay 3. Chemistry by Zumdahl

2 PRE-REQUISITE KNOWLEDGE Candidates should be able to describe and explain redox processes in terms of electron transfer and/or of changes in oxidation number (oxidation state). Please refer to Lecture Notes 1b Redox Reactions pages 2 – 8 to refresh your understanding on redox reactions and how to balance redox equations. An oxidation number is a number which is assigned to an element in a substance to show its state of oxidation. Note: The oxidation number of an element in a substance is related to the number of electrons lost, gained, or shared as a result of chemical bonding.  A positive change in oxidation state indicates that oxidation (loss of electrons) has occurred e.g. +1, +2.  A negative change in oxidation state indicates that reduction (gain of electrons) has occurred e.g. –1, –2. Rules for Assigning Oxidation Number 1. The oxidation number of an atom in the elemental state = 0 . Example: O.N. of Ca, I 2, S8 or P4 = 0 2. The oxidation number of a monoatomic ion is simply the charge on the ion. Example: O.N. of C l = 1; O.N. of Ca2+ = +2 3. The sum of the oxidation numbers of all the atoms or ions in an electrically neutral compound = 0. Examples: CaBr2 +2 + 2( 1) = 0 Al2O3 2(+3) + 3( 2) = 0 CO 2 +4+2( 2)=0 4. In polyatomic ions, the sum of the oxidation numbers = the charge on the ion. Examples: SO42 : The sum of the oxidation numbers (O.N of S = +6, O.N. of O = 2) = +6 + 4(2) = 2 = charge on the ion CrCl63 : The sum of the oxidation numbers (O.N of Cr = +3, O.N. of Cl = 1) = +3 + 6(1) = 3 = charge on the ion 1. In a covalent bond, the more electronegative atom has the negative oxidation number while the less electronegative atom has the positive oxidation number.  Electronegativity is the ability of an atom in a molecule to attract shared electrons in a bond.  Electronegativity values: F > O > N > C > H Examples: O.N. of F in OF 2 = 1 O.N. of O in OF 2 = +2

3 6. The oxidation numbers of groups 1, 2, 13 and 17 elements, hydrogen and oxygen are listed in the table. Note the exceptions. Element O.N. Examples Exception Group 1 (e.g. Na) Always +1 NaB r -- Group 2 (e.g. Mg) Always +2 MgO -- Group 13 (e.g. Al) Always +3 A l2O3 -- Group 17 Usually 1 (F is always 1) NaF Except when forming compounds with more electronegative atom E.g. O.N. for Cl in ClO2 = +4 O.N. for C l in in ClF = +1 Hydrogen Usually +1 HC l 1 in metal hydrides such as NaH (since Na is more electropositive than H) Oxygen Usually 2 CaO 1 in peroxides* such as H2O2 (hydrogen peroxide) and BaO2 (barium peroxide) *Peroxide ion is O22 0.5 in KO2 (potassium superoxide*) *Superoxide ion is O2 +2 in OF2 (fluorine monoxide) since F is more electronegative than O Balancing Redox Equations Candidates should be able to construct redox equations using the relevant half–equations.

4 A. REDOX PROCESSES & ELECTROLYTES A.1 Redox: Reduction & Oxidation 1. Oxidation Reduction o Increase in oxidation state of the element involved o Loss of electron(s) o Occurs at anode (AN OX) o Decrease in oxidation state of the element involved o Gain of electron(s) o Occurs at cathode (RED CAT) Reducing Agent Oxidising Agent o A substance that reduces another substance o Is itself oxidised o Electron donor, i.e. loses e – o A substance that oxidises another substance o Is itself reduced o Electron acceptor, i.e. gains e – 2. A redox reaction is one in which both reduction and oxidation occur. In fact, both processes must occur at the same time i.e. one cannot occur without the other. Cu2+(aq) + Zn(s)  Cu(s) + Zn2+(aq) 3. The equation for a redox reaction may be separated into two halfequations: Cu2+(aq) + Zn(s)  Cu(s) + Zn2+(aq) reduction halfequation: Cu 2+(aq) + 2e  Cu(s) oxidation halfequation: Zn(s)  Zn2+(aq) + 2e The half equations show the transfer of electrons that occurs between the relevant species in a redox reaction. Note that the number of electrons lost must be equal to the number of electrons gained in the reaction. Thus, a balanced redox equation does not contain any electrons. 4. A disproportionation reaction is a redox reaction in which the same element in a substance is simultaneously oxidised and reduced. 2H +(aq) + 2NO2–(aq)  H2O(l) + NO(g) + NO2(g) In this reaction, some NO 2 ions are being reduced to NO while simultaneously, other NO 2 ions are being oxidised to NO2. Oxidation number of N increases from +3 in NO 2 to +4 in NO2. Oxidation number of N decreases from +3 in NO 2 to +2 in NO. 5. Some everyday examples of redox reacti ons include rusting, respiration and combustion, reactions in batteries and cells, extraction of ores, and electroplating.

5 A.2 Electrolytes 1. An electrolyte is a compound which will conduct an electric current when it is in aqueous solution or in the molten state. An electrolyte conducts electricity due to the flow of charge carried by its ions. 2. A strong electrolyte is a compound which is fully ionised in aqueous solution. All ionic compounds are strong electrolytes since they are made up of ions which become mobile in the molten or aqueous state. E.g. NaOH(aq), PbCl2(l). sodium hydroxide: NaOH(aq)  Na+(aq) + OH(aq) molten lead(II) chloride: PbC l2(l)  Pb2+(l) + 2Cl(l) Some covalent compounds ionise completely in water and are also strong electrolytes. E.g. hydrogen chloride or HNO3. HC l(g) + H2O(l)  H3O+(aq) + Cl(aq) HNO 3(l) + H2O(l)  H3O+(aq) + NO3(aq) 3. A weak electrolyte only partially ionises in aqueous solution. E.g. water,