My teacher told me that elements get oxidised when:

1)There is a loss of electron

2) There is an addition of oxygen

3)There is a loss of hydrogen

And that they get reduced when:

1)There is a gain of electron

2) There is a loss of oxygen

3)There is an addition of hydrogen

Now I don't understand these things:

1) Most elements will lose an electron (and get oxidised) on the addition of oxygen due to its high electronegativity. But what about OF2? In this case shouldn't oxygen get oxidised and fluorine get reduced due to the higher electronegativity of fluorine? (I know its a covalent bond but I am talking in terms of on which side the electron will get attracted)

2) How does loss of hydrogen (or addition of hydrogen) oxidise (or reduce) an element? Like there are elements that have higher electronegativity than hydrogen then how will they get oxidised on the loss of hydrogen? They should get reduced because of the gain of an electron from hydrogen, right?

  • $\begingroup$ Fluorine is an exception, since it is more electronegative than oxygen, it gets reduced. $\endgroup$
    – Yulmart
    Jul 31 '16 at 14:18

The formal definition of an oxidation/reduction is linked to the loss/gain of electrons. Considering the addition or loss of hydrogen and oxygen is not a global rule, but rather a trick generally used in organic chemistry. If you consider organic molecules, thus mainly composed of carbon, "adding" oxygen atoms suggests an oxidation of the carbon, as oxygen is more electronegative than carbon, while "adding" hydrogen atoms suggests a reduction, as hydrogen is less electronegative than carbon.

An example would be the successive oxidation of ethanol to ethanal, then acetic acid : $$\ce{CH3CH2OH ->[-2 e^-][``-2\text{ H}"] CH3CHO ->[-2e^-][``+1\text{ O}"] CH3COOH}$$


Your teacher is correct in these definitions. However, in many reactions there is no oxygen or hydrogen involved as such, so the loss/gain of electrons is the most general definition. It is important to understand that in one reaction a species may be oxidised but reduced in another, which is what you suggest from your question.

For example in the reaction
2Fe$^{3+}$+Zn= 2Fe$^{2+}$+Zn$^{2+}$ the zinc is oxidised to Zn$^{2+}$ and so reduces the Fe$^{3+}$ to Fe$^{2+}$

In the reaction
I$_2$ + H$_2$SO$_3$ +H$_2$O= 2HI + H$_2$SO$_4$
the iodine acts as an oxidiser (so is reduced) to the strong reducer H$_2$SO$_3$, (which is oxidised) but in the reaction
I$_2$ +5Cl$_2$+6H$_2$O=HIO$_3$+10HCl
iodine acts as a reducer as chlorine is a much stronger oxidiser.

In complicated reactions like these it is useful to use oxidation numbers to work out what is oxidised and reduced.

If you find a textbook with a table of redox reactions you will see that species have different potentials which is how we know which are the strongest oxidisers (most oxidising Fluorine positive potential) and reducers (most reducing, lithium, negative potential). There are standard ways of working out whether a reaction will occur between oxidising and reducing species.

There is a description of calculating oxidation numbers in my answer to this question
Oxidation Number Of Phosphorus In PH3

  • $\begingroup$ But your answer doesn't actually answer my question. Does the addition of oxygen always oxidise an element? Then what about compounds like OF2 where oxygen is reducing fluorine. And how does the loss of hydrogen oxidise an element which is in a compound containing hydrogen? $\endgroup$ Jul 26 '16 at 9:59
  • $\begingroup$ I made some general comments which I thought would help you understand that some atoms can be either oxidising or reducing depending on their reaction partner. As fluorine is the most oxidising species it oxidises oxygen and is itself reduced. (It may help to think of oxygen as some species X then statements like oxidising oxygen seem less absurd). If you look at oxidation numbers, oxygen goes from 0 to +2, oxygen acts as reducer, and fluorine from 0 to -1, acts as oxidiser. $\endgroup$
    – porphyrin
    Jul 26 '16 at 12:19
  • $\begingroup$ So my teacher's statements are not always correct? $\endgroup$ Jul 26 '16 at 15:16
  • $\begingroup$ They are correct as far as I can see within the scope of the present knowledge of those that are taking your course. You will realise that as you progress things become more complicated, and hence more interesting, but in teaching its usually necessary to give a simplified version of things to get one started on a new subject. $\endgroup$
    – porphyrin
    Jul 26 '16 at 19:24
  • $\begingroup$ But how does loss of hydrogen oxidise an element? Can you please give me an example? $\endgroup$ Jul 27 '16 at 9:59

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