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Electrochemistry. Year 11 – Stage 1 SACE. In 1793, Alessandro Volta discovered that electricity could be produced by placing two dissimilar metals on opposite sides of a moistened paper.
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Electrochemistry Year 11 – Stage 1 SACE
In 1793, Alessandro Volta discovered that electricity could be produced by placing two dissimilar metals on opposite sides of a moistened paper. • In 1800, Nicholson and Carlisle, using Volta‘s primitive battery as a source, showed that an electric current could decompose water into oxygen and hydrogen. • By 1812, the Swedish chemist Berzelius could posit that all atoms are electrified, hydrogen and the metals being positive, the non- metals negativecalling them ions
When a piece of zinc metal is immersed in pure water, the following occur: • Zn ions go in the solution leaving electrons in the metal • This continues till the level of electrons increase too much and the ions are unable to leave the metal anymore, hence, the process comes to a halt. Zn(s)Zn2+ + 2e–
When Zinc metal is immersed in Copper sulphate solution, the following things occur: • Zinc metal goes in the solution • Copper metal deposits on the Zinc metal • The blue colour of the solution starts disappearing. Zn(s) + CuSO4 ZnSO4 + Cu(s)
Activity Series of metals lists metals in descending order of its activity. A reactive metal displaces any metal in its ion form.
terminology • Electrochemical Cell (EC): A cell which converts chemical energy into electrical energy or vice –a –versa • Electrolytic cell: An EC which converts electrical energy in a chemical energy. • Galvanic (Voltaic) Cell: An EC cell used to convert chemical energy into electrical energy • Electrolyte: A compound when molten or dissolved has the capacity to conduct electricity and inturn gets decomposed. Depending upon the level of ionization, electrolytes can be called strong, weak or a non-electrolyte which does not ionize when dissolved in water.
Electrode: A metal or Carbon rod which connects the electrolyte with a circuit. • Anode: The electrode where oxidation takes place • Cathode: The electrode where reduction takes place. • Electrolysis: A process in which electrical energy is used to produce a chemical reaction.
Galvanic Cell: http://www.youtube.com/watch?v=nNG5PMlHSoA
Voltaic (Galvanic) Cell Construction: • 2 half EC cells – in each cell a half reaction takes place. • Each half cell is made up of dipping a metal strip in its salt solution. E.g: Zinc metal strip dipping in a solution of Zinc salt. Copper metal strip dipping in a solution of Copper salt • The two half cells are connected with a salt bridge • The redox reaction is spontaneous but the electrons travel through an external circuit.
Points to remember • Oxidation occurs at the Anode (-) Zn(s)Zn2+ + 2e– • Reduction occurs at the Cathode (+) Cu2+ + 2e– Cu(s) • Electrons generated by the reaction move through the external circuit. These electrons are used by the cathode for the reduction reaction • To maintain electron neutrality, cations move towards cathode and anions move towards anode. • The electrons always flow from the negative to the positive (i.e. anode to cathode)
Notations of a voltaic cell • The anode is always written on the left, the cathode on the right. • The salt bridge is denoted by 2 vertical lines. • A single vertical line represents a phase change • E.g. Zn(s) Zn2+(aq) Cu2+(aq) Cu(s)
When a gas is involved, then, an inert metal like platinum is used as an electrode, which just catalyses the reaction, does not actively participate in the reaction. • E.g. for cathode: H+(aq) H2(g), Pt • For Anode: Pt, H2(g) H+(aq)
Construction of an Electrolytic cell • Electrolytic cell: An EC which uses electric current to cause a non – spontaneous chemical reaction • Consists of: • A battery (source of electron) • Electrolyte • electrodes
http://group.chem.iastate.edu/Greenbowe/sections/projectfolder/flashfiles/electroChem/electrolysis10.htmlhttp://group.chem.iastate.edu/Greenbowe/sections/projectfolder/flashfiles/electroChem/electrolysis10.html
Points to remember • When the charge is supplied electrons are drawn away from the Anode hence, anode becomes positively charged, attracting negatively charged ions • Electrons are supplied to the cathode hence becoming negatively charged, attracting positively charged ions. • Anode receives electrons, oxidises anions or water. • Cathode provides electrons, reduces cations or water
Applications: • Extraction of metals from their compounds, e.g. zinc • Recharging a lead – acid battery. • Electroplating, e.g. nickel, chromium • Refining of metals, e.g. copper • Production of non – metallic elements, e.g. chlorine, hydrogen
Molten electrolytes • Molten Electrolyte is defined as when the electrolyte is kept at a temperature higher than its boiling point. • For e.g. : Lead (II) Bromide. • Reaction at anode: 2Br-(l) Br2 (g) + 2e- • Reaction at Cathode: Pb2+ (l) + 2e- Pb(s) • Another example :NaCl
Aqueous electrolysis • Reactions carried out in aqueous solution provide an added complication. Water ionises to produce hydrogen (H+) ions and hydroxide (OH¯) ions. These are often displaced from solution to give hydrogen gas and oxygen gas. Which ions are released depends upon the following factors: • the position of the ions in the reactivity series. • Concentration of ions in the electrolytes. • Nature of the electrodes.
