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Electrochemistry(2 lectures)

Mr. Zaheer E. Clarke

1:00 p.m. / 2:15 p.m.

½ full question on C10K Paper 1

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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-

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• 2Mg 2Mg2+ + 4e-

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

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• 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

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• 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)

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• Reaction proceeds spontaneously (Gθ is –ve) & will continue until it reaches equilibrium

• Half Equations

Zn(s) Zn2+(aq) + 2e-

Cu2+(aq) + 2e-

Cu(s)

CuSO4

Zn metal strip

Oxidation

Reduction

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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

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Galvanic Cell

• Example of a Galvanic Cell is a Daniell Cell (Early)

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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

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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)?

1. Thickness & porosity of the porous pot

2. Cleanness of the electrodes

3. Electrical Resistance of the Measuring Device

• Internal Resistance

• Combat– Porous pot must be as thin & porous as possible– Electrodes clean– High Resistance Voltmeter

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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

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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

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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

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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+

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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 reactionCu + 2Ag+ Cu2+ + Ag EMF = 0.46V

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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

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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)

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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

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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)

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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 – Eθ

L = 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

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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)

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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

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Gθ & Eθ

Gθ = -ν F Eθ

– 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

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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)

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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+]

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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

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Nernst Equation - pH

• When H+ concentration is NOT 1.00 M but everything remains the sameQ = [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

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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