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Reduction and oxidation |
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Oxidation is the loss of electrons, reduction is the gain of electrons Oxidation half equation... Mg Reduction half equation... O + 2e- Calculation of oxidation number There are a few rules to remember
It is then simply a matter of adding together all of the oxidation numbers of the elements in a compound and making sure that the total is = 0 Example:
When naming the oxidation number of an atom within a compound it is normal to use a Roman Numeral. This becomes necessary when there is some ambiguity in using just the name. Example Phosphoric acid is a term that could apply to acids containing phosphorous with different oxidation states. To specify the acid the oxidation state of the phosphorus must be stated. H3PO3 - phosphoric III acid H3PO4 - phosphoric V acid Similarly Copper oxide could be either CuO or Cu2O , therefore it is important to differentiate by including the oxidation state of the copper in the name CuO - Copper II oxide Cu2O - Copper I oxide If an element is oxidized, its oxidation number will go up (get more positive). If an element is reduced, its oxidation number will go down. To find out, simply write down the oxidation numbers for each element within the compounds as explained previously. Examples Iron II hydroxide As the oxidation state of the iron increases from II to III then it has been oxidised
The oxidation state of the chromium changes from VI to III and so it has been reduced An oxidizing agent is an element which causes oxidation (and is reduced as a result) by removing electrons from another species. A reducing agent is an element which causes reduction (and is oxidized as a result) by giving electrons to another species. In the previous example the Ethanol is the reducing agent as it has brought about the reduction of Chromium (VI) to Chromium (III). Why can't the sulphuric acid be the reducing agent? answer A more reactive metal will displace a less reactive one from a compound and a more reactive halogen will displace a less reactive one from a compound. This can be generalized to say a stronger reducing agent will displace a weaker one from a compound, and a stronger oxidizing agent will displace a weaker one from a compound. Thus, if a metal displaces another, we know it must be more reactive and ditto for halogens (which are the given examples). Feasibility of reaction If an oxidising agent is placed with a reducing agent then they are likely to react. The feasibility of a reaction can be estimated from the substances positions in a reactivity series. For example: Zinc will displace Copper from a solution of Copper sulphate and therefore it should be placed above Copper in a reactivity series. Magnesium is above Zinc in the same series and so by inspection it should be expected to also displace Copper from its salt solution Electricity from chemical reactions As a redox reaction involves transfer of electrons form one species to another it can be used to produce an electric current if the electrons are made to pass around an external circuit to get from the reducing agent to the oxidising agent. This set-up is known as an electrochemical cell. Electrochemical cell diagram
In this cell the Zinc anode dissolves and releases electrons which pass around the external wires to the Copper electrode where they are given to the Copper ions (2+) which are then deposited as Copper atoms on the electrode. At the anode: Zn At the cathode: Cu2+ + 2e The overall cell reaction is found by adding the two equations together: Zn + Cu2+ The salt bridge is there to allow ions to pass through from one half cell to the other so that the half cells do not develop an electrical charge. Electrolysis This is, fundamentally, the reverse of an electrochemical cell. In this case a greater electromotive force is applied from the external circuit using a battery or power source and this forces the species within the cell to perform the reverse reaction to that which they would normally tend to do. It is important not to confuse electrochemical cells which generate electricity by means of a redox reaction and electrolytic cells which use electricity to perform chemical reactions. Electrolysis of molten ionic liquids The ionic substance is forced to reverse the ionisation that produced the original ions. Hence electrolysis of molten lead II bromide will give lead and bromine PbBr2 consists of Pb2+ ions and Br- ions in a giant ionic lattice. When melted (molten) these ions are free to move. The Pb2+ ions go to the cathode where they pick up electrons and become Lead atoms Pb2+ + 2e The Br- ions go to the anode where they lose electrons and become Bromine molecules 2Br- Conduction of current Electricity is a flow of electrons around a circuit. The electrolyte (liquid that conducts electricity) does so by means of ions depositing and picking up electrons from each electrode. Electrons are picked up by positive ions at the cathode and electrons are deposited by negative ions at the anode. Although the electrons have not technically been passed through the conducting liquid, from the point of view of the battery they have gone into the solution via the cathode and are exiting via the anode. As far as the battery is concerned there is a flow of electrons through the cell. Stages:
Net result: The battery sends out 1 electron from the negative side and receives 1 electron at the positive side. Electroplating The ability of electricity to force chemical change makes the deposition of metals possible in an electrolytic cell. If copper sulphate solution is electrolysed then the copper ions present in the solution will migrate to the cathode where they pick up electrons are get deposited as copper atoms. A layer of copper builds up on the cathode. Use can be made of this to plate conducting (metal) items by inserting them into the circuit instead of a normal cathode. If they are used to pass the elctrons into solution them they themselves get a deposit of metal. We call this electroplating. It may be carried out for any metal low down in the electrochemical (reactivity) series such as Copper, Nickel, Silver etc. Electroplating silver
Notice that in the diagram the anode is made of silver. This allows it to dissolve during the process (releasing electrons to the circuit) and replace the silver ions being removed from the solution at the cathode. The anode reaction: Ag(s) The cathode reaction: Ag+(aq)
+ 1e
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Answers 1. Because all of the atoms from sulphuric acid have the same oxidation number in the products as they did in the beginning |
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Copyright: 2003 Isis Publication
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Mg2+ + 2e-. 
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