Electrochemistry

(

2

lectures) Mr. Zaheer E. Clarke 1:00 p.m. / 2:15 p.m.

½

full

question on C10K Paper 1

1

Oxidation/Reduction Reactions In Cells

• Most chemical reactions involve the transfer of electrons between atoms & molecules?

– Not always clearly seen!

– Eg.

• 2Mg (s) + O 2(g)

• May be written:

• 2Mg 2MgO (s) 2Mg 2+ + 4e and • 4e + O 2 2O 2 2

• 2Mg 2Mg

2+

+ 4e

Mg has been oxidized (lost electrons)

OIL

(

O

xidation

I

s

L

ost) Mg is the reducing agent

• 4e

-

+ O

2

2O

2 O 2 has been reduced (gained electrons)

RIG

(

R

eduction

I

s

G

ain) O 2 is the oxidizing agent 3

• 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 4

• What happens when a Zn metal strip is inserted in a CuSO 4 solution?

• Solution will lose its blue colour – Cu metal is being deposited on the Zn strip • Zn is going into solution as Zn 2+ ions while Cu coming out of solution as metallic Cu 2+ is • How can this be written in in terms of chemical equations?

• Overall Reaction Zn (s) + Cu 2+ (aq) Zn 2+ (aq) + Cu (s) 5

Zn metal strip CuSO 4 • Reaction proceeds spontaneously (  G θ continue until it reaches equilibrium is –ve) & will • Half Equations Zn (s) Zn 2+ (aq) + 2e Cu 2+ (aq) + 2e Cu (s) Oxidation Reduction 6

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 7

Galvanic Cell

• Example of a Galvanic Cell is a

Daniell Cell (Early)

8

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 – is –ve Anode is +ve (external electricity), cathode • Galvanic Cell – Anode is -ve, cathode is +ve • Always holds true Anode =

Oxidation

Cathode =

Reduction

9

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.

2.

3.

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 10

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

11

Liquid Junction Potential How do we overcome the

LIQUID JUNCTION POTENTIAL?

– Use a salt bridge potentials to reduce the effect of liquid junction (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

12

Daniell Cell with Salt Bridge instead of Porous Pot

• •

Salt bridge

consists of 5% agar jelly mixed with a saturated solution of KCl or KNO 3 (K + , Cl and NO 3 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 13

Daniell Cell

• The spontaneous reaction that drives this cell is: Cu 2+ (aq) + Zn (s) Cu (s) + Zn 2+ (aq) • Cu 2+ has a greater tendency to pull electrons than Zn 2+ and that difference in electron pulling potential is what appears as a difference in electrical potential • Cu 2+ is a better oxidizing agent than Zn 2+ 14

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 Cu 2+  Ag + is a better oxidizing agent than Cu 2+ • The spontaneous reaction Cu + 2Ag + Cu 2+ + Ag EMF = 0.46V

15

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 | Cu 2+ half cell has a characteristic potential Zn | Zn 2+ 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 16

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

E θ

in

Volts

• The SHE is represented as: Pt (s) | H 2(g) | H + (aq) 17

Standard Electrode Potentials

• • Standard potentials are measured with the

side test electrode on the right hand

• The measured potential is

+ve

if the electrode has a than the H 2

greater tendency to pull electrons

electrode (SHE) and

–ve

if it has a

lower tendency Reduction Potentials

– – – – – Cu 2+ + 2e Zn 2+ Ag + + 2e + e   Ag Zn Pb 2+ + 2e  Pb Pb 4+ + 2e  Cu  Pb 2+

E θ E θ E θ E θ E θ = + 0.34V

= - 0.76V

= + 0.80V

= - 0.13V

= + 1.67V

18

Standard Cell Notation

• A vertical line vertical lines represents a phase boundary represent the salt bridge while double (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) | ZnSO 4(aq) || CuSO 4(aq) | Cu (s) Zn (s) | Zn 2+ (aq) || Cu 2+ (aq) | Cu (s) or 19

Standard Cell Notation & E

θ • (R) • (L) Cu 2+ Zn 2+ + 2e + 2e   Cu Zn

E E θ θ = + 0.34V

= - 0.76V

• Overall (R) - (L) Cu 2+ (aq) + Zn (s) Cu (s) + Zn 2+ (aq) • Overall E θ = E θ R – E θ L = 0.34 – (-0.76) = 1.10V

• When E θ is +ve direction written the reaction is spontaneous in the • If the Zn electrode was written as the cathode , the E θ would be –ve & the reaction would be spontaneous in the opposite direction 20

E

θ

– Indicator of Spontaneity

•  G θ is the maximum non-expansion (useful) work available from the reaction •  G θ can be equated to the electrical work done ( assuming constant pressure & temp.

) as the cell runs down & reaches equilibrium •  G θ = - (electrical work that can be done by the system) = (charge transferred) x (potential against which the charge is transferred) 21

E

θ

– Indicator of Spontaneity

• Work done = -

ν

e N A E θ

ν

– number of electrons transferred for each single oxid./red.

e – charge on each electron N A – is the single reactions per mole of reaction/Avogadro’s constant e N A = Faraday constant = 9.6485 x 10 4 C mol -1 •  G θ = -

ν

F E θ – work is done reversibly – constant pressure & temperature 22



G

θ

& E

θ •  G θ = -

ν

F E θ – When

E θ is +ve

, 

G θ is -ve

= reaction is

spontaneous

– When

E θ is -ve

, 

G θ is +ve

= reaction is

not spontaneous

•  G θ for the Daniell Cell  G θ = -(2)(96485)(1.10) = 212267 J mol -1 = 212.3 kJ mol -1 23

Nernst Equation

• Recall  G at any stage of rxn =  G  + RT ln Q • -

ν

F E = -

ν

F E θ + RT ln Q •

E = E θ – (RT/ ν F) ln Q

Nernst Equation • At unit activity E = E θ of the components (a = 1), ln Q = 0 & • At equilibrium ln Q = ln K & E = 0 (  G = 0) 24

Nernst Equation

• If we have a Cu 2+ /Cu electrode in one half & the SHE in the other Pt (s) | H 2(g) | H + (aq) || Cu 2+ (aq) | Cu (s) E θ = + 0.34V

(R) (L) Cu 2+ + 2e  2H + + 2e  Cu H 2 E θ = + 0.34 V E θ = 0 V (R) - (L) Cu 2+ (aq) + H 2(g) Cu (s) + 2H + (aq) E θ = + 0.34 V • Q = [a H+ ] 2 [a Cu ]/[a H2 ][a Cu2+ ] 25

Nernst Equation

• Q = [a H+ ] 2 [a Cu ]/[a H2 ][a Cu2+ ] • 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 = E θ – (RT/

ν

F) ln Q E = 0.763 – 0 = 0.763 V 26

Nernst Equation - pH

• When H + concentration is remains the same

NOT

1.00 M but everything Q = [a H+ ] 2 E = E θ – (RT/ E = 0.34

ν

F) ln{[a H+ ] 2 } – (RT/

ν

F) ln{[a H+ ] 2 } • The measured potential is related to the activity/ concentration of H + and E θ 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 27

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 28