1. Chpt. 20: Electrochemistry I

2. Humphry Davy (early 1800’s) Used electricity to separate compounds
into their elements
He discovered
potassium, sodium,
calcium, barium, strontium and magnesium
Start of
electrochemistry
Humphry Davy
(early 1800’s)

3. Electrochemistry Electrolysis Cells/Batteries
(use of electric current to bring about chemical reactions)
(use of chemical reactions to bring about electric current)

4. Electrolysis

5. Electrolysis
Electrolysis is defined as the use of electricity to bring about a chemical reaction in an electrolyte, i.e. a chemical reaction is produced when an electric current is passed through an electrolyte.
ELECTROLYSIS - 1st coined by Michael Faraday – Davy’s greatest discovery!
Electrolysis involves passing an electric current through an electrolyte in contact with electrodes.

6. An electrolyte is a compound which when molten or dissolved in water conducts electricity as it has ions that can move.
Electrodes – rods (two) that dip into the electrolyte and make electrical contact with it.

7. Inert Electrodes
DO NOT react with the electrolyte into which they are dipping
merely carry current in and out of electrolyte
most common - graphite & platinum
Active Electrodes
DO react with the electrolyte into which they are dipping
copper and iron

8. In general:
- one electrode is positively charged - ANODE
- other electrode is negatively charged - CATHODE

9. The science behind electrolysis!!!
Essentially chemical reactions are occurring at each electrode.
The battery ‘pumps’ electrons to the NEGATIVE electrode (cathode) where they are gained by some part of the electrolyte – REDUCTION
Electrons are lost at the POSITIVE electrode (anode) by some species in the electrolyte – OXIDATION
OPRN – Oxidation at Positive, Reduction at Negative

10. Five examples of electrolysis reactions involving solutions
Electrolysis of (acidified) H2O (Hofmann’s voltameter)
Electrolysis of Na2SO4 (Higher Level)
Electrolysis of KI (Higher Level)
Electrolysis of CuSO4 using Cu electrodes
To demonstrate the movement of ions under the influence
of an electric field

11. Please see handout for notes on each of these demonstrations.

12. Electroplating
Electroplating is the process where electrolysis is used to put a layer of one metal, usually expensive, on the surface of another metal, usually cheaper
Examples include gold plating (jewellery), chromium plating (taps), tin plating (food storage) and EPNS – Electro Plated Nickel Silver (cutlery)

13. Rules for Electroplating
The object to be plated must be attached to the negative terminal of the battery
The electrolyte must be a salt of the metal being plated onto the object
The positive electrode must be made of the same metal that is being plated onto the object

14. Example of electroplating – spoon being electroplated with
Example of electroplating – spoon being electroplated with silver metal
Object being plated MUST BE connected to –ive terminal - SPOON
Electrolyte MUST BE a salt of silver metal. Electrolyte SILVER NITRATE
Positive electrode MUST BE made of silver – bar of SILVER METAL

15. Electroplating - reactions involved
Negative Electrode (spoon):
Ag+ + e- → Ag↓
Ag+ ions from silver nitrate attracted to –ive electrode where they accept an e- to form Ag atom
Thus, silver metal is plated on to negative electrode – SPOON
’DRIVING FORCE’
Positive Electrode (Silver bar):
Ag → Ag+ + e-
Silver atoms in the +ive electrode lose an e- and go into solution as Ag+ ions
’Driving force’ is the need to replace the Ag+ being removed from solution at the –ive electrode

16. Example of Electroplating – copper plating a key
What is happening at negative electrode???
What is happening at positive electrode???

17. Electroplating Other uses include:
- protecting objects against corrosion:
- bath fittings (taps & towel rails), handlebars of
bicycle, motorbikes etc. - steel coated
with chromium
- food cans (tin cans) – made of steel coated with
tin (tin very unreactive i.e. does not react with food
stored in can)
Silver, chromium, zinc and tin are the metals most commonly electroplated on to cheaper metals

18. Cells/Batteries

19. Having studied how electric current can be used to bring about a chemical reaction the following section involves a study of the reverse i.e. how a chemical reaction can be used to bring about an electric current

20. Major contributors to the discovery/production of the cell/battery were Luigi Galvani and Alessandro Volta
A cell in which a chemical reaction results in the production of an electric current is called a galvanic cell or voltaic cell

21. Generating Electric Current in a simple cell
Draw diagram here

22. Generating Electric Current in a simple cell
Zinc Electrode:
Zn atoms (from electrode) have strong tendency to lose 2e-
Thus, Zn2+ ions formed
Oxidation rxn has taken place:
Zn → Zn e-
Copper Electrode:
Cu ions (from soln.) accept 2e- forming Cu atoms
Reduction rxn has taken place:
Cu e- → Cu↓
2e- flow around wire to copper electrode
Voltage generated = 1.1V
*Note: Salt bridge – helps to maintain a balance of ions in each part of the cell.
Simple salt bridge – piece of filter paper soaked in potassium chloride or potassium nitrate

23. Why does Zinc and not Copper lose its electrons????
Some metals have a greater tendency to lose their outer electrons than other metals.
If Zn is replaced by Mg then a much higher voltage is observed this is due to the fact that Mg has a far greater tendency to lose electrons than Zn. (Higher voltage indicates greater tendency to loose e-)
The electrode potential is the term used to describe the tendency of a metal to lose electrons i.e. tendency for metal to be oxidised.
Electrode potential values for each metal are obtained by measuring the voltage when the half-cell of that metal is connected to reference electrode – the hydrogen half cell. Thus, giving rise to the ELECTROCHEMICAL SERIES

24. Electrochemical Series
Electrochemical Series: is a list of elements in order of their standard electrode potentials

25. Electrochemical Series
Based on values of standard electrode potentials:
Potassium Hydrogen
Calcium Copper
Sodium Mercury
Magnesium Silver
Aluminium Gold
Zinc
Iron
Lead
Must know order of these metals!!!

26. Metals near the top of the series have the greatest tendency to lose electrons (oxidised) forming positive ions, hence they are VERY REACTIVE
Metals near the bottom of the series have only a small tendency to lose electrons, hence they are VERY UNREACTIVE (can occur in nature as the free metal e.g. gold)

27. Activity Series
Based on reactions with acid, air and water i.e. chemical reactivity:
Potassium (Pretty) Hydrogen (Honolulu)
Sodium (Sally) Copper (Causing)
Calcium (Could) Mercury (Many)
Magnesium (Marry) Silver (Strange)
Aluminium (A) Gold (Gazes)
Zinc (Zulu)
Iron (In)
Lead (Lovely)

28. Uses of Electrochemical Series
Electrochemical series allows us to predict the results in displacement reactions – a metal will displace a metal that is below it in the electrochemical series from a solution of its ions.
Example:
1) If a piece of Mg metal is dipped into a solution of copper ions
(CuSO4) the Mg metal becomes coated in Cu:
Reaction: Mg + Cu2+ → Mg2+ + Cu↓
2) Displacement of Cu metal from a solution of copper salts using Fe Fe + Cu2+ → Fe2+ + Cu↓
Application - Scrap Fe is used to extract Cu metal from Cu salts

29. In general, an element higher up the electrochemical series will displace one lower down in the series from a solution of its salt

30. Example 1: (Higher Level Only)
Complete and balance the equation for the chemical reaction that occurs when a piece of aluminum is placed in a solution of copper (II) ions:
Cu Al →