Explain with example, why there cannot be an oxidation reaction without a corresponding reduction reaction

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Explanation:

By definition, 'oxidation' is the loss of electrons from an element during a reaction process. The phrase 'loss of electrons' is a misnomer in that the electrons are actually not 'lost' but transferred to another element willing to gain the electrons. The process of gaining electrons is by definition 'reduction'.

A reaction that is an Oxidation-Reduction Reaction is basically two 'half-reactions'; one being the oxidation half reaction and the other the reduction half reaction. For example, consider an aqueous reaction of Copper(II) Sulfate with Zinc solid in the following reaction scheme.

Zn(s) + CuSO₄(aq) => ZnSO₄(aq) + Cu(s)

In this reaction Zinc undergoes oxidation losing (transferring) 2 electrons to Copper(II) ion gaining the 2 electrons in a reduction process. This can be represented by showing two half-reactions; one for oxidation and one for reduction as follows:

Zn(s) => Zn⁺²(aq) + 2e⁻  => Oxidation shows electrons on right of arrow.

Cu⁺²(aq) + 2e⁻ => Cu(s) => Reduction shows electrons on left of arrow.

When these two reactions are added together, the electron numbers should cancel leaving the net oxidation reduction reaction ...

Zn(s) + Cu⁺²(aq) => Zn⁺²(aq) + Cu(s)

The SO₄²⁻(aq) ion is a spectator ion and does not undergo reaction. It is there only as a 'required' counter-ion for the elements that do undergo oxidation-reduction.

c. 75.7

i think that is the right