Electrochemistry (2 lectures) Mr. Zaheer E. Clarke 1:00 p. m. / 2:15 p Electrochemistry (2 lectures) Mr. Zaheer E. Clarke 1:00 p.m. / 2:15 p.m. ½ full question on C10K Paper 1

Oxidation/Reduction Reactions In Cells Most chemical reactions involve the transfer of electrons between atoms & molecules? Not always clearly seen! Eg. 2Mg (s) + O2(g) 2MgO(s) May be written: 2Mg 2Mg2+ + 4e- and 4e- + O2 2O2-

2Mg 2Mg2+ + 4e- 4e- + O2 2O2- Mg has been oxidized (lost electrons) OIL (Oxidation Is Lost) Mg is the reducing agent 4e- + O2 2O2- O2 has been reduced (gained electrons) RIG (Reduction Is Gain) O2 is the oxidizing agent

The steps involved in the electron transfer when metallic elements (CONDUCTING ELECTRODES) & solutions of their salts (CONDUCTING SOLUTIONS) are combined can be isolated & observed clearly oxidation/reduction When the reactions occur spontaneously, the separation of the oxidation/reduction sites gives rise to potential difference which can drive e- through external resistive circuit Eg. Zn-Cu couple

What happens when a Zn metal strip is inserted in a CuSO4 solution? Solution will lose its blue colour – Cu metal is being deposited on the Zn strip Zn is going into solution as Zn2+ ions while Cu2+ is coming out of solution as metallic Cu How can this be written in in terms of chemical equations? Overall Reaction Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

Reaction proceeds spontaneously (Gθ is –ve) & will continue until it reaches equilibrium Half Equations Zn(s) Zn2+(aq) + 2e- Cu2+(aq) + 2e- Cu(s) Oxidation Reduction

Galvanic Cell Presently both oxidation & reduction occurs at the same site! – Zn metal strip If we SEPARATE these sites where oxidation & reduction occurs we would have what is called a Galvanic Cell What is a Galvanic Cell? A Galvanic or Voltaic Cell is one in which a spontaneous chemical reaction drives electrons from Anode to Cathode in an external circuit

Galvanic Cell Example of a Galvanic Cell is a Daniell Cell (Early) Galvanic Cells are formed by separating the site of oxidation & reduction in a spontaneous redox reaction, using solutions of ionic substances separated by a porous pot

Galvanic Cell Vs Electrolytic Cell In electrolytic cells an external source of electricity is used to drive a chemical reaction e.g. electrolysis of a salt solution Electrolysis – Anode is +ve (external electricity), cathode is –ve Galvanic Cell – Anode is -ve, cathode is +ve Always holds true - Anode = Oxidation Cathode = Reduction

Cell Potential EMF of a Daniell cell is 1.10V but this is not seen I practice Factors that affect the measured potential in a cell (e.g. Daneill Cell)? Thickness & porosity of the porous pot Cleanness of the electrodes Electrical Resistance of the Measuring Device Internal Resistance Combat Porous pot must be as thin & porous as possible Electrodes clean High Resistance Voltmeter

Liquid Junction Potential The liquid junction, porous pot, is also a source of “lost” potential Why is this? Build-up of charge results from the different mobilities of the ions as they move across the wall of the porous pot to neutralize the charge Potential difference exists between the inner & outer surfaces of the wall of the porous pot This potential that results is called a LIQUID JUNCTION POTENTIAL

Liquid Junction Potential How do we overcome the LIQUID JUNCTION POTENTIAL? Use a salt bridge to reduce the effect of liquid junction potentials (i.e. potentials which arise because of the difference in mobilities of the ions) Once these precautions are taken the EMF of the cell depends solely on the concentrations of the solutions & the metals used as the electrodes

Daniell Cell with Salt Bridge instead of Porous Pot Salt bridge consists of 5% agar jelly mixed with a saturated solution of KCl or KNO3 (K+, Cl- and NO3- have similar mobilities) The salt bridge reduces the LJP because of the large difference in concn. of the ions in the bridge compared to in the electrolyte solutions Effects due to mobility & availability of ions at the interface between the bridge & the solutions becomes negligible EMF of a Daniell Cell is 1.10V when the solutions are 1M

Daniell Cell The spontaneous reaction that drives this cell is: Cu2+(aq) + Zn(s) Cu(s) + Zn2+(aq) Cu2+ has a greater tendency to pull electrons than Zn2+ and that difference in electron pulling potential is what appears as a difference in electrical potential Cu2+ is a better oxidizing agent than Zn2+

Electrode Potential & Half Cells If a copper/silver or a zinc/silver cell was constructed a different potential would be observed for each cell (i.e. not 1.10 V) In a copper/silver cell, the silver is the +ve electrode and the copper is the –ve electrode Ag+ has a greater tendency to pull electrons than Cu2+ Ag+ is a better oxidizing agent than Cu2+ The spontaneous reaction Cu + 2Ag+ Cu2+ + Ag EMF = 0.46V

Electrode Potential & Half Cells Each electrode, i.e. the ion & its neutral atom, [Ag+/Ag] Contributes a characteristic potential to the overall cell potential Independent of the other electrode in the pair Cu | Cu2+ half cell has a characteristic potential Zn | Zn2+ half cell has a characteristic potential Ag | Ag+ half cell has a characteristic potential To assign a potential to each half cell one must assign an electrode as a “standard electrode” & measure each electrode relative to this standard electrode

Standard Hydrogen Electrode (SHE) The standard to which all electrodes are compared is the Standard Hydrogen Electrode Its characteristic potential is ZERO at ALL temperatures Potentials measured against the SHE are called Reduction Potentials and are represented by Eθ in Volts The SHE is represented as: Pt(s) | H2(g) | H+(aq)

Standard Electrode Potentials Standard potentials are measured with the test electrode on the right hand side The measured potential is +ve if the electrode has a greater tendency to pull electrons than the H2 electrode (SHE) and –ve if it has a lower tendency Reduction Potentials Cu2+ + 2e-  Cu Eθ = + 0.34V Zn2+ + 2e-  Zn Eθ = - 0.76V Ag+ + e-  Ag Eθ = + 0.80V Pb2+ + 2e- Pb Eθ = - 0.13V Pb4+ + 2e- Pb2+ Eθ = + 1.67V

Standard Cell Notation A vertical line represents a phase boundary while double vertical lines represent the salt bridge (no liquid junction potential) Standard Notation for Cells is based on this assumption: Right hand electrode is the cathode (where reduction occurs) Daniell Cell can be written as Zn(s) | ZnSO4(aq) || CuSO4(aq) | Cu(s) or Zn(s) | Zn2+ (aq) || Cu2+ (aq) | Cu(s)

Standard Cell Notation & Eθ (R) Cu2+ + 2e-  Cu Eθ = + 0.34V (L) Zn2+ + 2e-  Zn Eθ = - 0.76V Overall (R) - (L) Cu2+(aq) + Zn(s) Cu(s) + Zn2+(aq) Overall Eθ = EθR – EθL = 0.34 – (-0.76) = 1.10V When Eθ is +ve the reaction is spontaneous in the direction written If the Zn electrode was written as the cathode, the Eθ would be –ve & the reaction would be spontaneous in the opposite direction

Eθ – Indicator of Spontaneity Gθ is the maximum non-expansion (useful) work available from the reaction Gθ can be equated to the electrical work done (assuming constant pressure & temp.) as the cell runs down & reaches equilibrium Gθ = - (electrical work that can be done by the system) = - (charge transferred) x (potential against which the charge is transferred)

Eθ – Indicator of Spontaneity Work done = -ν e- NA Eθ ν – number of electrons transferred for each single oxid./red. e- – charge on each electron NA – is the single reactions per mole of reaction/Avogadro’s constant e- NA = Faraday constant = 9.6485 x 104 C mol-1 Gθ = -ν F Eθ work is done reversibly constant pressure & temperature

Gθ & Eθ Gθ = -ν F Eθ Gθ for the Daniell Cell When Eθ is +ve, Gθ is -ve = reaction is spontaneous When Eθ is -ve, Gθ is +ve = reaction is not spontaneous Gθ for the Daniell Cell Gθ = -(2)(96485)(1.10) = 212267 J mol-1 = 212.3 kJ mol-1

Nernst Equation Recall Gat any stage of rxn = G + RT ln Q -ν F E = -ν F Eθ + RT ln Q E = Eθ – (RT/ ν F) ln Q Nernst Equation At unit activity of the components (a = 1), ln Q = 0 & E = Eθ At equilibrium ln Q = ln K & E = 0 (G = 0)

Nernst Equation If we have a Cu2+/Cu electrode in one half & the SHE in the other Pt(s) | H2(g) | H+(aq) || Cu2+(aq) | Cu(s) Eθ = + 0.34V (R) Cu2+ + 2e-  Cu Eθ = + 0.34 V (L) 2H+ + 2e-  H2 Eθ = 0 V (R) - (L) Cu2+(aq) + H2(g) Cu(s) + 2H+(aq) Eθ = + 0.34 V Q = [aH+]2 [aCu]/[aH2][aCu2+]

Nernst Equation Q = [aH+]2 [aCu]/[aH2][aCu2+] Perfect gas: a = p / p Pure liquids and solids , a = 1 For solutions at low concentration: a = [conc.]/ [1 mol dm-3] Q = [1.00/1.00]2 [1]/[1.0/1.0][1.00/1.00] = 1 E = Eθ – (RT/ ν F) ln Q E = 0.763 – 0 = 0.763 V

Nernst Equation - pH When H+ concentration is NOT 1.00 M but everything remains the same Q = [aH+]2 E = Eθ – (RT/ ν F) ln{[aH+]2} E = 0.34 – (RT/ ν F) ln{[aH+]2} The measured potential is related to the activity/ concentration of H+ and Eθ of the cell pH 4.00, E = 0.577V or pH 7.00, E = 0.754V pH can be measured electrically E.g. pH meter

Applications Galvanic cells are used in flashlights, clocks, watches, remote controllers as Dry Cells Rechargeable batteries are used in cars to start engines, cell phones, video cameras, computers Fuel Cells