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ERT 108 : PHYSICAL CHEMISTRY Equilibrium Electrochemistry. By; Mrs Hafiza Binti Shukor. ERT 108/3 PHYSICAL CHEMISTRY SEM 2 (2010/2011). By; Mrs Hafiza Binti Shukor. Subtopics. Half-Reactions and Electrodes Varieties of Cells The Electromotive Force Standard Potentials

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ERT 108 : PHYSICAL CHEMISTRY Equilibrium Electrochemistry

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Ert 108 physical chemistry equilibrium electrochemistry

ERT 108 :

PHYSICAL CHEMISTRY

Equilibrium Electrochemistry

By; Mrs Hafiza Binti Shukor

ERT 108/3PHYSICAL CHEMISTRY

SEM 2 (2010/2011)

By; Mrs Hafiza Binti Shukor


Subtopics

Subtopics

  • Half-Reactions and Electrodes

  • Varieties of Cells

  • The Electromotive Force

  • Standard Potentials

  • Applications of Standard Potentials

  • Impact on Biochemistry:

    Energy Conversion in Biological Cells

ERT 108/3PHYSICAL CHEMISTRY

SEM 2 (2010/2011)


Equilibrium electrochemistry

Equilibrium Electrochemistry

Electrochemical system :

heterogeneous system in which there is a

difference of electrical potential between 2 or

more phases.

ERT 108/3PHYSICAL CHEMISTRY

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Equilibrium electrochemistry1

Equilibrium Electrochemistry

  • An electrochemical cell consists:

  • two electrodes (or metallic conductors)

  • an electrolyte (an ionic conductor – may be a solution, a liquid or a solid).

  • An electrode & its electrolyte comprise an electrode compartment.

  • two electrodes may share the same compartment

  • if the electrolytes are different, the two compartments may be joined by a salt bridge [a tube containing a concentrated electrolyte solution (potassium chloride in agar jelly)] that completes the electrical circuit & enables the cell to function.

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Equilibrium electrochemistry2

Equilibrium Electrochemistry

  • A galvanic cell is an electrochemical cell that produces electricity as a result of the spontaneous reaction occurring inside it. [ Eg: Daneil Cell]

  • A electrolytic cell is an electrochemical cells in which a non-spontaneous reaction is driven by an external source of current.

Terminals (electrods) made from same metal

Ionic conductor (at least 1 phase) to allow electronic charge to be transferred

ERT 108/3PHYSICAL CHEMISTRY

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Varieties of cells

Varieties of Cells

Anode

(Negative)

Cathode (positive)

  • Daniel cell:

  • Porous ceramics barrier separates a compartment containing (prevent extensive mixing of the solutions by convection currents but allow ions to pass from solution to the other):

    the redox couple at one electrode is Cu2+/Cu and at the other is Zn2+/Zn.

Cu2+(aq) + Zn Cu + Zn2+(aq)

ERT 108/3PHYSICAL CHEMISTRY

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Half reactions and electrodes

Half-Reactions and electrodes

  • Oxidation– the removal/loss of electronsfrom a species.

  • Reduction– the addition/ gain of electronsto a species.

  • Redox reaction – transfer of electrons from one species to another.

  • Reducing agent (reductant) – the electron donor.

  • Oxidizing agent (oxidant) – the electron acceptor.

  • Any redox reaction (or even not redox reaction) may be expressed as the difference of two reduction half-reactions.

  • Half-reactions – conceptual reactions showing the gain of electrons @ the reduced & oxidized species in half-reaction form a redox couple.

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Example of half reactions and electrodes

Example of Half-Reactions and electrodes

During the operation of cell, the electrochemical reactions of:

Zn - 2e- Zn2+(aq)

oxidation

Occur…call it as Half Reactions

Cu2+(aq) + 2e- (Cu) Cu

reduction

Electron flow process

2e -(Zn) 2e- (Cu)

Overall galvanic cell reaction

Zn + Cu 2+(aq) Zn2+(aq) + Cu

a)Zn rod in ZnSO4 solution

b)Cu rod in CuSO4 solution

ERT 108/3PHYSICAL CHEMISTRY

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Half reactions and electrodes1

Half-Reactions and electrodes

  • Anode:

  • the electrode at which the oxidation occurs.

    (-): removal of e-.

  • Cathode:

  • the electrode at which the reduction occurs.

    (+): addition of e-.

ERT 108/3PHYSICAL CHEMISTRY

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Half reactions and electrodes2

Half-Reactions and electrodes

Zn - 2e- Zn2+(aq)

Cu2+(aq) + 2e- (Cu) Cu

Zn + Cu 2+(aq) Zn2+(aq) + Cu

ERT 108/3PHYSICAL CHEMISTRY

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Example 1

Example 1

  • Express the following reactions in terms of reduction half-reactions.

  • The dissolution of silver chloride in water:

    (Note: it is not a redox reaction.)

  • The formation of H2O from H2 and O2 in acidic solution.

ERT 108/3PHYSICAL CHEMISTRY

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Answer example 1

Answer (Example 1)

  • The two reduction half-reactions:

    The redox couples are AgCl/Ag,Cl-& Ag+/Ag.

  • Formation of H2O from H2 and O2 in acidic solution.The two reduction half-reactions:

    The redox couples are H+/H2 & O2,H+/H2O


Liquid junction potentials

Liquid junction potentials

  • Liquid junction potentials (Elj):

  • an additional source of potential difference across the interface of the two electrolytes.

  • E.g. In the Daniel cell

    (i) two different electrolyte solutions are in contact,

    (ii) different concentration of hydrochloric acid.

  • The contribution of the liquid junction to the potential can be reduced (to about 1 to 2 mV) by joining the electrolyte compartments through a salt bridge.

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Liquid junction potentials1

Liquid junction potentials

Galvanic cell without liquid junction.

Galvanic cell with liquid junction (salt bridge).

ERT 108/3PHYSICAL CHEMISTRY

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Concentration cells

Concentration cells

  • Electrolyte concentration cell:

  • the electrode compartment are identical except for the electrolytes concentrations.

  • Electrode concentration cell:

  • the electrodesthemselves have different concentrations either because they are gas electrodes operating at different pressures or because they are amalgams (solution in mercury) with different concentrations.

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Notation

Notation

1) Phase boundaries are denoted by a vertical bar.

Eg: Pt(s) ІH2(g)ІHCl(aq)ІAqCl(s)ІAg(s)

2) A liquid junction is denoted by

3) Interface is denoted by a double vertical line ||.

  • Fig 1:

  • Zn (s)|ZnSO4 (aq) CuSO4 (aq) |Cu (s)

  • Fig 2:

  • Zn (s)|ZnSO4 (aq)||CuSO4 (aq) |Cu (s)

Fig 1

Fig 2

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The electromotive force

The electromotive force

  • The cell reaction IS THE REACTION IN THE CELL WRITEN ON THE ASSUMPTION THAT:

    1st : write the right hand half-reaction as a reduction

    (Assumption: spontaneous reaction is 1 in which reduction is

    taking place in the right hand compartment).

    2nd : subtract from it the left-hand reduction half-reaction.

    (By implication, the electrode is the site of oxidation)

    In the cell: Zn(s)|ZnSO4(aq)||CuSO4(aq)|Cu(s)

    Right-hand electrode: Cu2+(aq)+2e- Cu(s)

    Left-hand electrode: Zn2+(aq)+2e- Zn(s)

    Overall cell reaction: Cu2+(aq)+ Zn(s) Cu(s) +Zn2+(aq)

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The nernst equation

The Nernst equation

  • A cell in which the overall cell reaction has not reached chemical equilibriumcan do electrical work as the reaction drives electrons through an external circuit.

  • the work that a given transfer of electrons can accomplish depends on the potential difference between the two electrodes.

  • This potential differences is called the cell potential and is measured in volts, V (1 V = 1 JC-1 s).

  • A cell in which the overall reaction is atequilibriumcan do no work, & then the cell potential is zero.

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The nernst equation1

The Nernst equation

  • When expressed in terms of a cell potential, the spontaneous direction of change can be expressed in terms of the cell emf.

  • the reaction is spontaneous when E>0.

  • the reverse reaction is spontaneous when E<0.

  • when the cell reaction is at equilibrium, the cell potential is zero.

Note: The potential difference is called the electromotive force (emf), E

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The nernst equation2

The Nernst equation

  • The Nernst equation relates the cell’s emf(E) to the activities ai of the substances in the cell’s chemical reaction & to the standard emfof the cell (E Ѳ)(the cell’s chemical reaction).

  • where F = Faraday constant, F=eNA

    v = the stoichiometric coefficient of the electron

    in the half-reactions.

    Q = the reaction (or activity) quotient ,

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The nernst equation3

The Nernst equation

  • A practical form of the Nernst equation is

  • because at 250c,

ERT 108/3PHYSICAL CHEMISTRY

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Cells at equilibrium

Cells at equilibrium

  • Suppose the reaction has reached equilibrium; then Q = K (K= the equilibrium constant of the cell reaction).

  • A chemical reaction at equilibrium cannot do work, & hence it generates zero potential difference between the electrodes of a galvanic cell.

  • Setting E=0, Q=K:

  • the Nernst equation:

ERT 108/3PHYSICAL CHEMISTRY

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Example 2

Example 2

Three different galvanic cells have standard electromotive force (EѲ) of 0.01, 0.1 and 1.0V, respectively, at 250C.

Calculate the equilibrium constants (K) of the reactions that occur in these cells assuming the charge number (v) for each reaction is unity.

ERT 108/3PHYSICAL CHEMISTRY

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Answer example 2

Answer (Example 2)

  • For EѲ = 0.01V,

    = 1.476

  • For EѲ = 0.1V, K = 49.0

  • For EѲ = 1.0V, K = 8.02 x 1016

ERT 108/3PHYSICAL CHEMISTRY

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Standard potentials

Standard Potentials

  • A standard potential of a couple is the cell potential in which it forms the right hand electrode and the left hand electrode is a standard hydrogen electrode

  • A galvanic cell is a combination of two electrodes, & each one can be considered as making a characteristics contributions to the overall cell potential.

  • although it is not possible to measure the contribution of a single electrode, we can define the potential of one of the electrodes as zero & then assign values to others on that basis.

  • the specially selected electrode is the standard hydrogen electrode (SHE): Pt(s)|H2(g)|H+(aq),

    EѲ=0 (at all temperatures).

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Standard potentials1

Standard Potentials

  • To achieve the standard conditions, the activity of the hydrogen ions must be 1 (pH=0) & the P of the hydrogen gas must be 1 bar.

  • The standard potential (EѲ) of another couple is then assigned by constructing a cell in which it is the right-hand electrode & the standard hydrogen electrode (SHE) is the left-hand electrode.

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Standard potentials2

Standard Potentials

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Example 3

Example 3

  • Consider the following galvanic a cell:

  • What is:

  • the cell reaction?

  • the standard electromotive force?

  • The equilibrium constant?

    Given Table 9.1 standard potential for:

Cu2+(aq)+2e- Cu(s) E0=0.34V

Zn2+(aq)+2e- Zn(s) E0= - 0.76V

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Answer example 3

Answer (Example 3)

(a) & (b)

Right-hand electrode: Cu2+(aq)+2e- Cu(s) E0=0.34V

Left-hand electrode: Zn2+(aq)+2e- Zn(s) E0= - 0.76V

Overall cell reaction:Cu2+(aq)+ Zn(s) Cu(s) +Zn2+(aq)

E0 = 0.34 – (-0.76) V

E0 = 1.10 V

(c) The equilibrium constant:

K = 1.80 x 1037

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Application of standard potentials

Application of standard potentials

  • The electrochemical series list the metallic elements in the order of their reducing power as measured by their standard potentials in aqueous solution

  • The cell potential is used to measure the activity coefficient of electroactive ions

  • The standard cell potential is used to infer the equilibrium constant of the cell reaction

  • Species selective electrodes contribute the potential that is characteristics of certain ions in solution

  • The temperature coefficient of the cell potential is used to determine the standard entropy and enthalpy of reaction.

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Impact on biochemistry energy conversion in biological cells

Impact on Biochemistry: Energy Conversion in Biological Cells.

  • The whole of life’s activities depends on the coulping of exergonic & endergonic reactions, for the oxidation of food drives other reactions forward.

  • In biological cells, the energy released by the oxidation of foods is stored in adenosine triphosphate (ATP).

  • the essence of the action of ATP is its ability to lose its terminal phosphate group by hydrolysis & to form adenosine diphosphate (ADP).

    where Pi denotes an inorganic phosphate group e.g. H2SO4.

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Impact on biochemistry energy conversion in biological cells1

Impact on Biochemistry: Energy Conversion in Biological Cells.

  • Examples:

  • Glycolysis – the oxidation of glucose to CO2 and H2O by O2 (the breakdown of foods is coupled to the formation of ATP in the cell).

  • Glycolysis is the main source of energy during anaerobic metabolism, a form of metabolism in which inhaled O2 does not play a role.

  • The citric acid cycle & oxidative phosphorylationare the main mechanisms for the extraction of energy from carbohydrates during aerobic metabolism (in which inhaled O2 does play a role).

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The end

THE END…

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