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Electrochemistry

Electric current is the flow of electric charge [A = C/s]

metallic conduction (electrons)

electrolytic conduction (ions)

Main topics of electrochemistry

Galvanic cells Electrolysis

A cell that uses a chemical redox-reactionto produce electrical energy

The use of electric current to bringabout chemical changes

Electrical conduction

The spontaneous reaction between zinc and copper(II) ion.

A Galvanic cell (voltaic cell) based on the zinc-copper reaction.

Oxidation half-reactionZn(s) → Zn2+(aq) + 2e-

Reduction half-reactionCu2+(aq) + 2e- → Cu(s)

Overall (cell) reactionZn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

Oxidation (loss of e-) occurs at the anode, which is therefore the source of e-.

Zn(s) → Zn2+(aq) + 2e-

Over time, the Zn(s) anode decreases in mass and the [Zn2+] in the electrolyte solution increases.

Reduction (gain of e-) occurs at the cathode, where the e-

are used up.

Cu2+(aq) + 2e- → Cu(s)

Over time, the [Cu2+] in this half-cell decreases and the mass of the Cu(s) cathode increases.

Operation of the Voltaic Cell

A voltaic cell based on the zinc-copper reaction.

Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

Oxidation half-reactionZn(s) → Zn2+(aq) + 2e-

After several hours, the Zn anode weighs less as Zn is oxidized to Zn2+.

Reduction half-reactionCu2+(aq) + 2e- → Cu(s)

The Cu cathode gains mass over time as Cu2+

ions are reduced to Cu.

The anode produces e- by the oxidation of Zn(s). The anode is the negative electrode in a voltaic cell.

Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

Electrons flow through the external wire from the anode to the cathode, where they are used to reduce Cu2+ ions.

The cathode is the positive electrode in a voltaic cell.

Charges of the Electrodes

The salt bridge completes the electrical circuit and allows ions to flow through both half-cells.

As Zn is oxidized at the anode, Zn2+ ions are formed and enter the solution.

Cu2+ ions leave solution to be reduced at the cathode.

The salt bridge maintains electrical neutrality by allowing excess Zn2+ ions to enter from the anode, and excess negative ions to enter from the cathode.

A salt bridge contains nonreacting cations and anions, often K+ and NO3

-, dissolved in a gel.

The Salt Bridge

Zn(s) → Zn2+(aq) + 2e- Cu2+(aq) + 2e- → Cu(s)

Electrons flow through the wire from anode to cathode.

Cations move through the salt bridge from the anode solution to the cathode solution.

Zn2+

Anions move through the salt bridge from the cathode solution to the anode solution.

SO42-

By convention, a voltaic cell is shown with the anode on the left and the cathode on the right.

Flow of Charge in a Voltaic Cell

Figure 21.6 Measuring the standard cell potential of a zinc-copper cell.

The standard cell potential is designated E°cell (or electromotive force, emf) and is measured at a specified temperature with no current flowing and all components in their standard states.

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17.3. Electrode potentials

electrical double layer electrode potential, inV

It is impossible to determine the absolute value of the potential for a single electrode!

Metal-electrode

Mez+

Solution

ze

+++++

_____

anode ()

half-cell

Half-reaction (oxidation): Me Mez+ + ze

The standard electrode potential (E°half-cell) is the potential of a given half-reaction when all components are in their standard states.

By convention, all standard electrode potentials refer to the half-reaction written as a reduction.

The standard cell potential depends on the differencebetween the abilities of the two electrodes to act as reducing agents.

E°cell = E°cathode (reduction) - E°anode (oxidation)

Standard Electrode Potentials

Half-cell potentials are measured relative to a standard reference half-cell.

The standard hydrogen electrode has a standard electrode potential defined as zero (E°reference = 0.00 V).

This standard electrode consists of a Pt electrode with H2gas at 1 atm bubbling through it. The Pt electrode is immersed in 1 M strong acid.

2H+(aq; 1 M) + 2e- H2(g; 1 atm) E°ref = 0.00V

The Standard Hydrogen Electrode

Determining an unknown E°half-cell with the standard reference (hydrogen) electrode.

Oxidation half-reactionZn(s) → Zn2+(aq) + 2e−

Reduction half-reaction2H3O+(aq) + 2e- → H2(g) + 2H2O(l)

Overall (cell) reactionZn(s) + 2H3O+(aq) → Zn2+(aq) + H2(g) + 2H2O(l)

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Standard hydrogen electrode

p(H2) = 101.3 kPa (atm. pressure)[H+] = 1 mol/dm3 ; T = 298.15 K (25 oC)0 = 0.000 V (arbitrarily assigned)

The potential of an electrode (half-cell) is expressed relative to the hydrogen-electrode.

Voltmeter

anode(-) cathode(+)

1 M emf = o(X) – o(H2)or

emf = o(H2) – o(X)here:

emf = o (Cu) – o(H2) = 0.34 V

emf = o (Cu) = + 0.34 V

Table 21.2 Selected Standard Electrode Potentials (298 K)

Half-Reaction E°(V)+2.87

−3.05

+1.36+1.23+0.96+0.80+0.77+0.40+0.34

0.00−0.23−0.44−0.83−2.71

strength of reducing agent

stre

ngth

of o

xidi

zing

age

nt

F2(g) + 2e− 2F−(aq)Cl2(g) + 2e− 2Cl−(aq)MnO2(g) + 4H+(aq) + 2e− Mn2+(aq) + 2H2O(l)NO3

-(aq) + 4H+(aq) + 3e− NO(g) + 2H2O(l)Ag+(aq) + e− Ag(s)Fe3+(g) + e− Fe2+(aq)O2(g) + 2H2O(l) + 4e− 4OH−(aq)Cu2+(aq) + 2e− Cu(s)

N2(g) + 5H+(aq) + 4e− N2H5+(aq)

Fe2+(aq) + 2e− Fe(s)2H2O(l) + 2e− H2(g) + 2OH−(aq)Na+(aq) + e− Na(s)Li+(aq) + e− Li(s)

2H+(aq) + 2e− H2(g)

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„lead”

golden crown

Electrochemistry and dentistry

The contacting metals form a local galvanic celland the less noble metal (tin) dissolves!

Figure 21.9 A dental “voltaic cell.”

Biting down with a filled tooth on a scrap of aluminum foil will cause pain. The foil acts as an active anode (E°aluminum = -1.66 V), saliva as the electrolyte, and the filling as an inactive cathode as O2 is reduced to H2O.

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o standard electrode potential V[Mez+] concentration of Mez+ mol/dm3

R ideal gas constant, 8.314 J/K·molT temperature KF faraday, 96 485 C/molz number of electrons involved in the half reaction

17.5. Effect of concentration on electrode potentials

Nernst equation for Mez+ + ze = Me, where [Me] = 1 M by definitionand ln (1/[Mez+]) = ln [Mez+]

Electromotive force of the Daniell-cell:

emf = 0Cu 0

Zn +RTzF ln [Cu2+]

[Zn2+]

= o + zFln [Mez+]RT

= o + 0.05916z log [Mez+]

at 25 oC

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Concentration cells

a Voltaic cell constructed from the same substances but that differ in their concentrations

e. g.

(-)Cu Cu2+(0.1 mol/dm3) Cu2+ (1.0 mol/dm3) Cu(+)

a = o + 0.059162 log 0.1 c = o + 0.05916

2 log 1.0

emf = c– a = 0.059162 log

1.00.1 = 0.02958 V

Ecell > 0 as long as the half-cell concentrations are different.The cell is no longer able to do work once the concentrations are equal.

A concentration cell based on the Cu/Cu2+ half-reaction.

Overall (cell) reactionCu2+(aq,1.0 M) → Cu2+(aq, 0.1 M)

Oxidation half-reactionCu(s) → Cu2+(aq, 0.1 M) + 2e-

Reduction half-reactionCu2+(aq, 1.0 M) + 2e- → Cu(s)

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Ksp = [Ag+][Cl]

= o(Ag) + FRT

lnKsp

[Cl]

constant constant

Reference electrodes

The electrode potentials, consequently the emf of a galvanic cell is changingcontinuously during the operation, due to the change of concentrations ofthe solutions.The e. p. of a reference electrode remains constant during the operation of the cell at a given temperature.(e.g. Ag/AgCl-, or the calomel, Hg/Hg2Cl2- electrode).

RT = o(Ag) + Fln[Ag+]

silver (Ag)

KCl solution

reference electrode

0(Ag/AgCl) = 0.2225 V

porous partition

Ag+

= o(Ag/AgCl) – (RT/F) ln [Cl-]

[Cl-] = 3.5 M - practically const.

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Measuring the pH using glass electrode

buffer solution

glass electrode

silver (Ag)

KCl solution

porous partition

(g) = o(g) + 0.05916·log[H+]

(ref) = constant; emf = (g) – (ref)

pH = log[H+]

reference electrode

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Laboratory measurement of pH.

The operation of a pH meter illustrates an important application of concentration cells. The glass electrode monitors the [H+] of the solution relative to its own fixed internal [H+].

An older style of pH meter includes two electrodes.

Modern pH meters use a combination electrode.

Table 21.3 Some Ions Measured with Ion-Specific Electrodes

Species Detected Typical Sample

NH3/NH4+ Industrial wastewater, seawater

CO2/HCO3- Blood, groundwater

F- Drinking water, urine, soil, industrial stack gases

Br- Grain, plant tissueI- Milk, pharmaceuticalsNO3

- Soil, fertilizer, drinking waterK+ Blood serum, soil, wineH+ Laboratory solutions, soil, natural

waters

Minimocroanalysis.

A microelectrode records electrical impulses of a single neuron in a monkey’s visual cortex. The electrical potential of a nerve cell is due to the difference in concentration of [Na+] and [K+] ions inside and outside the cell.

Alkaline battery.

Anode (oxidation): Zn(s) + 2OH-(aq) → ZnO(s) + H2O(l) + 2e-

Cathode (reduction): MnO2(s) + 2H2O(l) + 2e- → Mn(OH)2(s) + 2OH-(aq)Overall (cell) reaction:

Zn(s) + MnO2(s) + H2O(l) → ZnO(s) + Mn(OH)2(s) Ecell = 1.5 V

In a fuel cell, also called a flow cell, reactants enter the cell and products leave, generating electricity through controlled combustion.

Reaction rates are lower in fuel cells than in other batteries, so an electrocatalyst is used to decrease the activation energy.

Fuel Cells

Hydrogen fuel cell.

Anode (oxidation): 2H2(g) → 4H+(aq) + 4e-

Cathode (reduction): O2(g) + 4H+(aq) + 4e- → 2H2O(g)Overall (cell) reaction: 2H2(g) + O2(g) → 2H2O(g) Ecell = 1.2 V

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Corrosion is the process whereby metals are oxidized to their oxides and sulfides.

The rusting of iron is a common form of corrosion.

- Rusting requires moisture, and occurs more quickly at low pH, in ionic solutions, and when the iron is in contact with a less active metal.

- Rust is not a direct product of the reaction between Fe and O2, but arises through a complex electrochemical process.

Corrosion: an Environmental Galvanic Cell

Fe(s) → Fe2+(aq) + 2e- [anodic region; oxidation]O2(g) + 4H+(aq) + 4e- → 2H2O(l) [cathodic region; reduction]

The loss of iron:

2Fe(s) + O2(g) + 4H+(aq) → 2Fe2+(aq) + 2H2O(l) [overall]

The rusting process:

2Fe2+(aq) + ½O2(g) + (2 + n)H2O(l) → Fe2O3·nH2O(s) + 4H+(aq) Overall reaction:

H+ ions are consumed in the first step, so lowering the pH increases the overall rate of the process. H+ ions act as a catalyst, since they are regenerated in the second part of the process.

The Rusting of Iron

The corrosion of iron.

Enhanced corrosion at sea.

The high ion concentration of seawater enhances the corrosion of iron in hulls and anchors.

The effect of metal-metal contact on the corrosion of iron.

Fe in contact with Cu corrodes faster.

Fe in contact with Zn does not corrode. The process is known as cathodic protection.

The use of sacrificial anodes to prevent iron corrosion.

In cathodic protection, an active metal, such as zinc, magnesium, or aluminum, acts as the anode and is sacrificed instead of the iron.

Cobalt Alloys: ASTM F75• Co-Cr-Mo

The stem and neck shown above are recalled Stryker Rejuvenate hip replacement parts removed from one of our clients. The corrosion is evident on the neck (the middle piece).

Chromium and cobalt ions can cause damaging tissue reactions:fibrinoid necrosis (pseudo-tumors)lymphocytic (lymph cell) infiltration and aggregationvasculitis (inflammation causing thickening and scarring of the walls of the blood vessels).

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Electrolysis. Electrolysis in molten sodium chloride

reduction oxidation

cathode() anode(+)

Direct currentsource

mp.(NaCl) 801oC

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DC source

cathode(-) anode(+)

H2O H+ + OH-

17.1.2. Electrolysis of aqueous sodium sulfate solution

H+ ions discharge easier than Na+ ionsOH ions discharge easier than SO4

2 ions

Na2 SO4 2 Na+ + SO42-

cathode process (reduction):

2 H+ + 2 e- H2

anode process (oxidation):

2 OH ½ O2 + H2O + 2 e-

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Stoichiometry of electrolysis

For cathodic reductions of Mez+ + ze Me e.g. Na+ + e Na; Ca2+ + 2e Ca; Al3+ + 3e Al

The electric charge of 1 mole of electrons is called faraday:F = NA.e = 6.022 x 1023 mol1 x 1.6022 x 1019 C = 96 485 C/mol

(Requires 32 Li-ion (850 mAh) cellphone batteries to store 1 mol of electrons!)

Faraday’s LawQuantity of electricity (Q = I·t) is directly proportional to the chemicalchange (m) in electrolysis. (M. Faraday 1832-1833)

m mass of discharged substance [g]M atomic or molecular mass [g/mol]z number of electrons involved in the processQ quantity of electric charge used [C]F faraday, 96 485 [C/mol]

m = MQzF

2 500 mAh = 2.5 A 3 600 s = 9 000 As = C