oxidation and reduction. historically.... oxidation was defined as the addition of oxygen to a...
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Oxidation and Reduction
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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
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• 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
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• 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
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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
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• 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 ?
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• Consider the reaction of copper oxide to copper metal
Cu2+O2- + H2 Cu + H2O
Cu2+
gains 2 electrons
reduction
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• 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)
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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?
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• 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
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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
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• 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
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• 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
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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
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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
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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 + H2O2 Colourless Hair + H2O pigment pigment
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• 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
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• 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
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• 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
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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
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Rules for Assigning oxidation Numbers
1.The Oxidation No of any uncombined element is zero Eg. 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
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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
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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
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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- , SO4
2-,
CO32- etc. (ions like Cl- and Mg2+ are called
simple ions)
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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
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Balancing Redox Equations
• We will refer to examples p187