Additional Electrochemistry

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The second part of this higher option looks at electrochemistry in much more depth. ^[1] Hopefully, you will already be familiar with your chapter on electrochemistry and, in particular, the electrochemical series. Here, we take a look at its relationship with corrosion and the electrolysis of molten salts.

Contents 1 Electrolysis of Molten Salts 1.1 Electrolysis of Lead Bromide 2 Electrochemical Series 3 Corrosion 3.1 Formation of Rust 3.2 Prevention 3.3 Sacrificial Anodes 4 References

Electrolysis of Molten Salts

• Electrolysis, electrode and electrolyte were all coined by Michael Faraday.
• From your study of electrochemistry, you should be familiar with various examples of the electrolysis of different solutions.
• The electrolysis of a molten salt is much easier to predict.

Electrolysis of Lead Bromide

• If carrying out this experiment, do so in a well ventilated area (or in a fume cupboard).
• Once heated, lead forms at the cathode and bromine at the anode.
• Lead metal has been extracted from the PbBr₂ salt.
  • Negative electrode: Pb²⁺ + 2e^- gives Pb
  • Positive electrode: 2Br ^- gives Br₂ + 2e^-

Electrochemical Series

• A galvanometer is a device which detects very small currents.
• Alessandro Volta made the first battery by placing a brine solution between a zinc and copper plate.
• The electrochemical series is a list of metals sorted by their tendency to lose electrons.

The electrochemical series ^[2]

Elements (most reactive on top)
Potassium
Calcium
Sodium
Magnesium
Aluminium
Zinc
Iron
Lead
Hydrogen
Copper
Silver
Gold

Corrosion

• Metals at top of electrochemical series corrode easily, due to their tendency to lose electrons.
• Corrosion is any undesirable process where a metal is converted to one of its compounds.

Formation of Rust

• A common occurrence of corrosion is the conversion of iron to rust.
• Rust is an oxide of iron (Fe₂O₃.xH₂O).
• Aluminium protects itself from rusting by the formation of aluminium oxide on its surface.
• When iron comes in contact with water and oxygen, a redox reaction occurs.
• 2Fe gives 2Fe²⁺ + 4e^-
• O₂ + 2H₂O + 4e^- gives 4OH^-
• This reaction is accelerated in the presence of NaCl or acid rain.

Prevention

There are a number of ways of preventing rusting and corrosion in metals.

1. Painting and greasing
2. Galvanising
3. Surface coating with aluminium or chromium
4. Using sacrificial anodes

Sacrificial Anodes

• To protect a certain metal object against corrosion, a more reactive metal is placed in contact with it.
• When in electrolytic solutions, the more reactive metal will be worn away first.
• Examples include the use of zinc to protect steel boats.
• The zinc will have to be replaced periodically.

References

1. ↑ http://www.curriculumonline.ie/uploadedfiles/PDF/lc_chemistry_sy.pdf
2. ↑ http://www.chemguide.co.uk/physical/redoxeqia/ecs.html

Chemistry Topics
Periodic Table and Atomic Structure · Chemical Bonding · Stoichiometry, Formulas and Equations · Volumetric Analysis · Fuels and Heats of Reaction · Rates of Reaction · Organic Chemistry · Chemical Equilibrium · Environmental Chemistry: Water · Additional Industrial Chemistry · Atmospheric Chemistry · Materials · Additional Electrochemistry · Extraction of Metals
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