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Oxidation and Reduction

Oxidation and Reduction. Historically. Oxidation was defined as the addition of oxygen to a substance Eg . when coal was burned C + O 2 CO 2 or the rusting of iron 4Fe + 3O 2 2Fe 2 O 3.

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Oxidation and Reduction

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  1. Oxidation and Reduction

  2. Historically.... • Oxidation was defined as the addition of oxygen to a substance Eg. when coal was burned C + O2 CO2 or the rusting of iron 4Fe + 3O2 2Fe2 O3

  3. Chemists discovered it was possible to remove oxygen from some substances • Eg. When hydrogen gas is passed over heated copper oxide, copper metal is obtained CuO + H2 Cu + H2O

  4. This process was called reduction as it was often used to extract metals from their ores thus getting a reduced amount of metals out of a larger amount • Reduction became known as the removal of oxygen from a substance or since hydrogen was often used the addition of hydrogen to a substance

  5. Oxidation and Reduction in terms of Electron transfer • Many chemical reactions involve the transfer of electrons • In the oxidation of Magnesium to magnesium oxide 2Mg + O2 2Mg2+O2- Mg loses 2 electrons oxidation

  6. In this reaction Magnesium is losing two electrons and oxygen is gaining the two electrons • Oxidation of an element takes place when it loses electrons • In the equation we have just studied what is oxidised ?

  7. Consider the reaction of copper oxide to copper metal Cu2+O2- + H2 Cu + H2O Cu2+ gains 2 electrons reduction

  8. We can see Cu2+ gains 2 electrons to become a copper atom • Reduction of an element takes place when it gains electrons • Thinks of OIL RIG to help you remember • (oxidation is loss of electrons, reduction is gain of electrons)

  9. Redox Reactions • Oxidation and Reduction must always occur together, if one substance loses electrons there must be another substance there to gain those electrons • Think of the reaction between sodium and chlorine to form sodium chloride 2Na + Cl2 2Na+Cl- Can you identify what is reduced and what is oxidised?

  10. These types of reactions are called oxidation-reduction reactions or Redox reactions for short • It is clear from the equation between sodium and chlorine that neither hydrogen or oxygen are present showing that oxidation and reduction is much more broadly defined in terms of electron transfer

  11. Reaction Between Zinc Metal and Copper Ions • When zinc metal is left to stand in copper sulfate solution it is found that the zinc becomes covered in copper deposits

  12. As this reaction takes place the blue colour of the solution fades indicating the Cu2+ ions are being used up • On analysis of the liquid Zinc ions are found • It appears the following reaction has taken place Zn + Cu2+ Zn2+ + Cu

  13. The Zn is oxidised and the Cu is reduced • The Zn could not lose its electrons unless there was something there to accept them (in this case the Cu2+ ions) • We say that the Cu2+ ion is an oxidising agent

  14. Oxidising Agents • An oxidising agent is a substance that brings about oxidation in other substances • In oxidising the Zn the Cu2+ ion gains electrons and is itself reduced • The oxidising agent is always reduced

  15. Reducing Agents • Since Zn is the substance that causes the Cu2+ ion to be reduced we call it the reducing agent • A reducing agent brings about reduction in other substances • In giving electrons to the Cu2+ ion the zinc loses electrons and thus is oxidised • The reducing agent is always oxidised

  16. Identifying Oxidising and Reducing agents in Equations • One of the most common oxidising agents is Hydrogen Peroxide (H2O2) used to bleach hair • It converts coloured hair pigment to colourless hair pigment which appears blonde Coloured Hair + H2O2Colourless Hair + H2O pigment pigment

  17. Iodine used to treat cuts and chlorine used to disinfect swimming pool water work by oxidising chemicals in the cells of germs Live Germ + I2 2I- + Dead Germ Live Germ + Cl2 2Cl- + Dead Germ

  18. Carbon monoxide is a very useful reducing agent in industry • It is used to remove the oxygen from iron ore in order to convert it to pure iron Fe2O3 + CO 2Fe + 3CO2

  19. By using the group number of the elements involved in the reactions it can be determined whether the atoms will want to lose or gain electrons • For example group 1 elements will tend to want to lose one electron hence will be oxidised and are reducing agents • We will study more detailed rules for this later on • Try p190 Q14.2

  20. Oxidation Numbers • In order to keep track of electrons during chemical reactions involving covalent compounds chemists introduced the idea of oxidation numbers/states. • The Oxidation number of an atom is the charge that an atom appears to have when the electrons are distributed according to certain rules

  21. Rules for Assigning oxidation Numbers 1.The Oxidation No of any uncombined element is zeroEg. In Zn the Zn has an ON of 0, in O2 each O has an ON of 0 2.The Oxidation No of an ion of an element is the same as its charge Eg. The ON of Br in the Br- ion is -1 the ON of Na in Na+Cl- is +1 and of Mg in Mg2+O2- is +2 etc

  22. 3.The sum of all the elements in a compound must add up to zero Eg. in H2O each H is +1 and O is -2 therefore 2(+1) + 1 (+2) = 0 4.Oxygen has an Oxidation No of – 2 except in a peroxide or if joined with fluorine In Mg2+O2- the ON of oxygen is -2, however in H2O2 (hydrogen peroxide) it has an ON of -1 so that the sum of all the ON in the compound will be 0 2(+1) + 2(-1) = 0. In Oxygen difluoride OF2 Fluorine is more electronegative than oxygen and attracts electrons to itself each F has a charge of -1 so oxygen has a charge of +2 to ensure the ON’s add up to 0 (2(-1) + 1(+2) = 0

  23. 5.Hydrogen is assigned + 1 except if joined with metals when it will have an ON of -1 In NH3 each H atom is assigned the number of +1 , this would mean N must be -3 as 3(+1) + 1(-3) = 0 However in Sodium Hydride NaH because sodium is a metal H has an ON of -1

  24. 6.Halogens are – 1 unless bonded to a more electronegative atom Fluorine will always have an ON of -1 as it is the most electronegative atom in the periodic table, Chlorine usually has an ON of -1 except when it is bonded to oxygen or fluorine (as these are more electronegative than chlorine) in Cl2O each Chlorine has an ON of +1 in Cl2O7 each chlorine has an ON of +7 7. The sum of all the oxidation numbers in a complex ion must add up to the charge on the ion Examples of complex ions are NO3- , SO42-, CO32- etc. (ions like Cl- and Mg2+are called simple ions)

  25. Oxidation and Reduction in terms of oxidation numbers • When an element is oxidised its oxidation number increases, ie. oxidation is an increase in oxidation number • When an element is reduced its oxidation number decreases, ie. Reduction is a decrease in oxidation number

  26. Balancing Redox Equations • We will refer to examples p187

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