Redox - Introduction
- In redox reactions one element is reduced and oxidised.
- Reduction and oxidation must occur together because if one element loses an electron another element must gain them.
- Oxidation is the loss of electrons or increase in oxidation number.
- Reduction is the gain of electrons or decrease in oxidation number.
- Disproportionation a reaction in which an element is simultaneously oxidised and reduced.
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- Oxidation states allow us to work out which elements have been oxidised and reduced in a reaction.
- Oxidation states are "charges" assigned to each element in a reaction .
- The rules for assigning oxidation states are:
- Elements are 0.
- In compounds, H = +1 and O = -2.
- Group 1 = +1, Group 2 = +2, Group 6 = -2, Group 7 = -1.
- In a neutral compound the oxidation states add up to 0.
- In an ion the oxidation states add up to the charge on the ion.
- NO3- = -1
- Oxidation state of O = -2 x 3 = -6.
- Oxidation of N must be +1
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- In reactions, metals generally lose electrons = oxidation.
- In reactions, non - metals generally gain electrons = reduced.
- Metals undergo redox reactions with dilute hydrochloric and dilute sulphuric acids.
Zn + 2HCl --> ZnCl2 + H2
- Zn goes from 0 to +2 = oxidation.
- H goes from +1 to 0 = reduction.
Mg + H2SO4 --> MgSO4 + H2
- Mg goes from 0 to +2 = oxidation
- H goes from +1 to 0 = reduction
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