CHEM1612 - PharmacyWeek 10: Corrosion/Batteries

Dr. Siegbert Schmid

School of Chemistry, Rm 223

Phone: 9351 4196

E-mail: siegbert.schmid@sydney.edu.au

Unless otherwise stated, all images in this file have been reproduced from:

Blackman, Bottle, Schmid, Mocerino and Wille,     Chemistry, John Wiley & Sons Australia, Ltd. 2008

     ISBN: 9 78047081 0866

Lecture 29 - 3

Electrochemistry Blackman, Bottle, Schmid, Mocerino & Wille:

Chapter 12, Sections 4.8 and 4.9

Key chemical concepts: Redox and half reactions Cell potential Voltaic and electrolytic cells Concentration cells

Key Calculations: Calculating cell potential Calculating amount of product for given current Using the Nernst equation for concentration cells

NaCl

Lecture 29 - 4

Al is an expensive metal because of the stability of its oxide Al2O3.

Al cannot be electrolysed from solution because H2O is preferentially reduced (E0

Al = -1.66 V; EH20= -0.42 V).

Al cannot be electrolysed from the pure oxide because it melts at too high a temperature (2045 ºC).

In 1886, Hall and Herault independently developed a method for electrolytic production of Al metal, that is still used today.

Hall-Herault process: dissolve Al2O3 in hot cryolite, Na3AlF6, which reduces the melting point to about 900 ºC.

Production of Aluminium

Lecture 29 - 5

At the anode graphite is oxidised to CO2 (as a result the electrodes are rapidly used up), and fluoro-oxy ions are transformed in Al fluorides. Very high currents are used (~250,000 A) on an industrial scale.

Production of Aluminium

Hall-Herault process

Graphite-lined furnace

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Lecture 29 - 6

Refining of CuElectro-refining is the principle method by which Cu is refined to high purity.

Less easily reduced metals remain in solution.

Noble metals are not oxidised, so fall to the bottom as “mud”.

http://electrochem.cwru.edu/ed/encycl/fig/m02/m02-f06b.jpg

Lecture 29 - 7

Corrosion: Unwanted voltaic cells

The reduction of a metal oxide to a metal requires a lot of energy. This means that the reverse, oxidation of a metal to its oxide will be

exothermic, and likely to be spontaneous.

Metal

oxide

Metal

Reduction + energy

Oxidation, spontaneous

Economically, the most important corrosion process is that of iron or steel.

Lecture 29 - 8

Corrosion

Corrosion is the process by which metals are oxidised in the atmosphere.

In corrosion, a metal can act as both an anode and a cathode.

The electrons released at the anode travel through the metal to the cathode.

Eo for the reaction is positive (a spontaneous process, product favoured).

Al, Ti, Cr, Ni and Zn do not corrode (much) because they form an impervious oxide layer.

Corrosion results in loss of structural strength.

Iron roof

Lecture 29 - 9

The Mechanism of Corrosion1) Oxidation of Fe at active anode forms a pit and yields e- which travel through the metal

2) Electrons at the Fe (inactive) cathode reduce O2 to OH-

3) Fe2+ migrates through the drop and reacts with OH- and then O2 to form rust.

Lecture 29 - 10

Redox chemistry of corrosion The rusting of iron involves two (or more) redox reactions:

Anode: 2 x {Fe Fe2+ + 2e-} E ox

0 = 0.44 V

Cathode: O2 + 4H+ + 4e- 2H2O E 0 = 1.23 V

The Fe2+ is further oxidised at the edges of the droplet, where [O2] is highest:

Anode: 2 x {Fe2+ Fe3+ + e-} Eox 0 = -0.77 V

Cathode: ½ x {O2 + 4H+ + 4e- 2H2O} E 0 = 1.23 V

Iron (III) forms a very insoluble oxide (rust) which is deposited at the edge:

2Fe3+(aq) + (3+n) H2O(l) Fe2O3•n H2O(s) + 6H+(aq)

Lecture 29 - 11

You should now be able to explain some of the known features of rusting:

Why does iron not rust in dry air?

No water no “salt bridge”

Why does iron not rust in oxygen-free water, such as ocean depths?

No oxygen no oxidant

Why does iron rust more quickly in acidic environments?

H+ is a catalyst

Why does iron rust more quickly at the seaside?

More conductivity in the “salt bridge”

Chemistry of corrosion

Lecture 29 - 12

Protection against corrosion Fe can be protected by preventing O2 and H2O from reaching the metal,

by oiling the surface or coating with a thin film of metal oxide. Anything more readily oxidised than Fe will act as anode and prevent Fe

from oxidising. These sacrificial anodes can be made of any metal that is a stronger

reducing agent than Fe (“Activity Series of Metals”: Zn and Mg). This is called “cathodic protection”, and is used frequently in large iron

structure such as ships, pipes, bridges, etc

Zinc anode

Bronze rudder

Lecture 29 - 13

Mild steel bolts

Stainless steel hanger

Mild steel karabiner

Aluminium karabiner

Images from http://www.theleedswall.co.uk/ymc/boltfund.htm

Galvanic Corrosion

Lecture 29 - 14

Batteries Commercial use of redox reactions 3 classes of batteries:

Primary batteries: Non-rechargeable (e.g. alkaline battery) Secondary batteries: Rechargeable (e.g. lead-acid, Ni-Cd, Li-ion

batteries)

Fuel cells: Fuel (e.g. H2/O2) pass through the cell, which converts

chemical energy into electrical energy.

Lecture 29 - 15

Primary Batteries Alkaline battery

Use a solid alkaline electrolyte paste (KOH) .

Cannot be recharged, it is “dead” when its components reach equilibrium concentrations.

Anode: Zn + 2OH- ZnO + H2O + 2e- Eox0 = 1.25V

Cathode: 2MnO2 + H2O + 2e- Mn2O3 + 2OH- E0 = 0.12V

Overall:

Lecture 29 - 16

Secondary Batteries Lead-acid battery (rechargeable)

Used to start cars.

The battery is recharged (turning it into an electrolytic cell) to re-establish non-equilibrium concentrations.

Anode: Pb + HSO4- PbSO4 + H+ + 2e- Eox

0=0.30 V

Cathode: PbO2 + 3H+ + HSO4- + 2e- PbSO4 + 2H2O E0=1.63 V

Lecture 29 - 17

Fuel Cells A fuel cell is a voltaic cell where the reactants are a combustible fuel, e.g. H2, CH4. The fuel undergoes a normal (overall) combustion reaction, however the two half-reaction are separated and the electrons harnessed. Fuel cells are still in the experimental stage, and their most notable success is probably for production of energy and water in space .

Anode: H2 2H+ + 2e-; E0=0.0 V

Cathode: ½O2 + 2H+ + 2e- H2O; E0=1.23 V

Anode: CH4 + 2H2O CO2 + 8H+ + 8e-; Eox0=-0.3 V

Cathode: 4 x {½O2 + 2H+ + 2e- H2O}; E0=1.23 V

Lecture 29 - 18

Hydrogen fuel cell

Pt catalyst surrounding graphite electrode

←e-

• Efficient

• No pollutants

• Newer designs use polymer electrolyte membrane that ferries H3O+ groups across

Lecture 29 - 19

Li-ion batteries

On discharge Li-ions move from anode to cathode.

On charge Li-ions move from cathode to anode

In the case of LiCoO2 the battery is supplied in its discharged state.

Lecture 29 - 20

Summary of Electrochemistry Concepts

Redox reactions Standard reduction potential, E0

Reference electrodes Galvanic cells, cell notation, and electromotive force Ecell

Electrolytic cells and Faraday’s Law Nernst Equation and concentration cells Examples of biological concentration cells Relationship between E0, ΔG, Q, and K Corrosion Batteries