RI 2025 Electrochemistry (II)

1 RAFFLES INSTITUTION YEAR 6 H2 CHEMISTRY 2025 Lecture Notes 21b – Electrochemistry 2 (Electrolysis) _____________________________________________________________________________ A Content Electrolysis (i) Factors affecting the amount of substance liberated during electrolysis (ii) Industrial uses of electrolysis (iii) The Faraday constant; the Avogadro constant; their relationship B Learning outcomes Candidates should be able to: (a) state the relationship, F = Le, between the Faraday constant, the Avogadro constant and the charge on the electron (b) predict the identity of the substance liberated during electrolysis from the state of electrolyte (molten or aqueous), position in the redox series (electro de potential) and concentration (c) calculate: (i) the quantity of charge passed during electrolysis (ii) the mass and/or volume of substance liberated during electrolysis (d) 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] 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 D Online Resources 1. http://www.chemguide.co.uk/physical/redoxeqiamenu.html#top

2 1 INTRODUCTION 1.1 Basic set up • The process of passing electricity (from an external source) to force a non–spontaneous redox reaction to occur is called electrolysis. • The cell in which electrolysis occurs is called the electrolytic cell. • A typical electrolytic cell has the following components: • During electrolysis, there are two possibilities for “chemical change”: a) electrolyte – its components are oxidised or reduced. b) electrode – a ___________anode is oxidised. • Due to the redox processes occurring, some possible physical observations are: a) the electrolyte changes colour or colour intensity, b) bubbling at the cathode or anode (i.e. a gas is evolved), c) a solid metal deposits on the cathode, d) the anode decreases in size and mass (i.e. it dissolves). Note: • Substances are described as being liberated (for substances produced) or discharged (for species that underwent redox reaction) during electrolysis. • Any oxidation always takes place at the anode. • Any reduction always takes place at the cathode. A + – e– e–  External source of electrical energy i.e. battery  Electrolyte (molten or aqueous)  Electrodes (cathode/anode) e– flow to this electrode • chemical species accept e– from this electrode • reduction occurs here This electrode is the cathode. e– flow from this electrode • chemical species lose e– here • oxidation occurs here This electrode is the anode.

3 1.2 Comparison between Electrochemical (or Voltaic) Cell & Electrolytic Cell Type of Cell Electrochemical (or Voltaic) Cell Electrolytic Cell Diagram Spontaneous redox reaction in cell produces electricity. Electricity from external source forces a ________________ redox reaction to occur. Energy conversion chemical → electrical energy electrical → chemical energy External Source No Yes Salt Bridge Usually Yes No Electrode Anode Cathode Anode Cathode Half-equation Oxidation Reduction Oxidation Reduction Sign − + + − Electron flow in external circuit From anode to cathode From anode to cathode Electrochemical (or Voltaic) cell • Spontaneous redox reaction which provides source of electrical power Electrolytic cell • Supply electrical energy to make non-spontaneous redox reactions occur Note: Unlike in electrochemical cells, we DO NOT consider Ecell > 0 when setting up electrolytic cells. Non-spontaneous reactions are taking place instead, driven by an external power source. Key considerations for choosing which reactions take place are discussed in section 2. Useful tip: the polarity of electrodes in electrolytic cells are opposite that of electrochemical cell. An easy way to remember this is the cathode forces reduction and provides electrons from the power source. Hence it is negative. e– cathode anode V (–) (+) + e– cathode anode (+) A − (−) Chemical Energy Electrical Energy

4 2 FACTORS AFFECTING SELECTIVE DISCHARGE OF IONS Candidates should be able to 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. When an electrical potential is applied across the two electrodes in an electrolytic cell, not all species present undergo simultaneous discharge. In fact, the cation or anion that is discharged is the one that requires the least energy. This is called selective/preferential discharge. The order in which ions are discharged at the electrodes are influenced by: (A) the physical state of electrolyte (molten or aqueous); note that E values do not actually apply in molten states. (B) electrode potential, E, (C) the relative concentration of the species, (D) the nature of the electrodes (inert or reactive) (A) Physical State of Electrolyte (Molten or Aqueous) Compare the electrolysis of molten NaCl(l) and dilute NaCl(aq) molten NaCl(l) Species present in electrolyte Na+(l) and Cl–(l) Electrodes Inert electrodes e.g. graphite Cathode (–) Na+(l) migrates to the cathode (negative electrode) and is discharged i.e. Na+(l) gains electrons is reduced to Na(l). Na+(l) + e− → Na(l) Anode (+) Cl– migrates to the anode (positive electrode) and is discharged, i.e. Cl– loses electrons and is oxidised to Cl2. 2Cl–(l) → Cl2(g) + 2e− Overall equation 2NaCl(l) → 2Na(l) + Cl2(g) Setup Going into the next example, we will observe in aqueous dilute NaCl, the options for discharge are different, and selection criteria need to be applied. + e– cathode anode (+) − (−) carbon electrode carbon electrode Na+ liquid sodium, Na(l) chlorine gas, Cl2(g) Cl−

5 (B) Electrode Potential Recall: A more positive (or less negative) E implies a higher tendency for reduction to occur and conversely, a less positive (or more negative) E implies a higher tendency for oxidation to occur. Dilute aqueous NaCl Species present in electrolyte Na+(aq), Cl–(aq) and H2O(l) Electrodes Inert electrodes e.g. graphite Cathode (–) Na+(aq) + e− ⇌ Na(s) E = –2.71 V 2H2O(l) + 2e– ⇌ H2(g) + 2OH–(aq) E = –0.83 V Since the E (H2O/H2) is _____________ than E (Na+/Na), H2O has a greater tendency to be reduced than Na+ and is thus preferentially reduced. Reaction at cathode: _____________________________ Anode (+) O2(g) + 4H+(aq) + 4e– ⇌ 2H2O(l) E = +1.23 V Cl2(g) + 2e− ⇌ 2Cl−(aq) E = +1.36 V Since the E (O2/H2O) is _______________ than E (Cl2/Cl−), H2O has a greater tendency to be oxidised than Cl− and is thus preferentially oxidised. Reaction at anode: ______________________________ Overall equation 2H2O(l) → 2H2(g) + O2(g) The electrolysis of NaC l(aq) results in the electrolysis of water! Note: • For the half-equations and overall redox equation, use ‘→’ instead of ‘⇌’. • [H+] and [OH−] from the auto-ionisation of water are extremely low at 10 –7 mol dm–3 so their discharge is not significant. In summary, • cations are __________ at the cathode. If more than one cation is present, and the concentration of the cations are almost the same , the cation with a _____________ E will generally be reduced. [The reduction of H2O, if present, must also be considered.] • anions are __________ at the anode. If more than one anion is present, and the concentration of the anions are almost the same , the anion with a ____________ E will generally be oxidised. [The oxidation of H2O, if present, must also be considered.]

6 (C) Relative