# Electronegativity Considerations in Assigning Oxidation States

I have never seen anything other than a set of rules like these when textbooks present how to assign oxidation numbers. Such as these:

However, if we keep in mind that oxidation numbers are simply imaginary numbers which suppose all bonding to be ionic - i.e. not electron sharing - and if we keep in mind simply relative electronegativity, we can easily work out the oxidation state of any element in any compound.

For example, take water. Bonding order: $\ce{H-O-H}$. Oxygen has two lone pairs. Oxygen is more electronegative than hydrogen. So we suppose that oxygen takes both electrons in both bonding pairs. Oxygen has 8 electrons. Its valence is 6. 6-8 is -2; oxygen has two more electrons assigned to it than what it has in its valence, so naturally its oxidation state is negative 2. No rules needed.

Now hydrogen peroxide. $\ce{H-O-O-H}$. Here we have an oxygen-oxygen bond so in this case, neither element wins the electronegativity battle. Electrons are split between the oxygen and the oxygen. However, since the oxygens are much more electronegative than the hydrogens, we assign both electrons in both the $\ce{O-H}$ bonding pairs to oxygen. Oxygen has 7 electrons assigned to it; oxygen has a valence of 6; oxidation state: -1. No need to memorize exceptions as stated in the table above.

We could go on. However, I am curious:

1) Has anyone been taught to assign oxidation states this way? 2) If not, what do you think of this method?

• I have, the table came afterwards as an aid to quickly assign alkali and alkaline earth metals, oxygen and halides. – LDC3 May 20 '14 at 3:16
• I don't think I was taught that way but it's the way I use. I think oxidation numbers are a lot easier than how they are taught in high school. – canadianer May 20 '14 at 4:41
• I agree. Memorization is crap. EN isn't that hard and I remember I was taught to memorize the EN trend in high school. However, we were still taught how to assign oxidation numbers through memorizing a bunch of rules. Unfortunate that teachers can't put two and two together. – Dissenter May 20 '14 at 4:53

The electronegativity battle scheme is most helpful for all kinds of compounds since it is the most generic way to derive oxidation states. The table represents just a cheat sheet that might be very helpful in the beginning. If you are spending most of your time with chemistry, this table will be present as some sort of muscle memory - usually referred to as chemical intuition.

The IUPAC defines oxidation states in their goldbook as follows:

A measure of the degree of oxidation of an atom in a substance. It is defined as the charge an atom might be imagined to have when electrons are counted according to an agreed-upon set of rules:

1. the oxidation state of a free element (uncombined element) is zero;
2. for a simple (monatomic) ion, the oxidation state is equal to the net charge on the ion;
3. hydrogen has an oxidation state of 1 and oxygen has an oxidation state of -2 when they are present in most compounds. (Exceptions to this are that hydrogen has an oxidation state of -1 in hydrides of active metals, e.g. $\ce{LiH}$, and oxygen has an oxidation state of -1 in peroxides, e.g. $\ce{H2O2}$;
4. the algebraic sum of oxidation states of all atoms in a neutral molecule must be zero, while in ions the algebraic sum of the oxidation states of the constituent atoms must be equal to the charge on the ion. For example, the oxidation states of sulfur in $\ce{H2S}$, $\ce{S8}$ (elementary sulfur), $\ce{SO2}$, $\ce{SO3}$, and $\ce{H2SO4}$ are, respectively: -2, 0, +4, +6 and +6. The higher the oxidation state of a given atom, the greater is its degree of oxidation; the lower the oxidation state, the greater is its degree of reduction.

What we see is, that your table actually reflects some of these rules. However, this set is not generic at all and it lacks a definition for compounds that do not contain oxygen or hydrogen, are elemental or monoatomic ions. With only these rules it is impossible do determine the oxidation states for $\ce{BF3}$ and [many]$_{n~(n\to\infty)}$ compounds.

The lack of the definition is very well known - but not surprisingly the IUPAC has not changed the official set. Hans-Peter Loock proposed a much simpler concept in Expanded Definition of the Oxidation State:

The oxidation state of an atom in a compound is given by the hypothetical charge of the corresponding atomic ion that is obtained by heterolytically cleaving its bonds such that the atom with the higher electronegativity in a bond is allocated all electrons in this bond. Bonds between like atoms (having the same formal charge) are cleaved homolytically.

This is not a new definition, but it predates the IUPAC rules by several decades. For example, Linus Pauling provided a similar definition of the oxidation state in his 1947 edition of General Chemistry (3).

(3):Pauling, L. General Chemistry; Freeman: San Francisco, CA, 1947.republished by Courier Dover Publications, 2012

So to me it is completely unclear, why IUPAC has chosen a different definition.

• I reused that for the tag-wiki entry oxidation-state. – Martin - マーチン May 20 '14 at 6:43
• Thank you. It is unclear to me too why the IUPAC definition is this vague as well. – Dissenter May 20 '14 at 17:40

My apologies for bumping an old question; however I'm not sure that the accepted answer provides an adequate answer to the stated question:

1) Has anyone been taught to assign oxidation states this way? 2) If not, what do you think of this method?

For part 1, I can say personally that I use this method in teaching a 3rd year Inorganic Chemistry course. It is covered in Rayner-Canham's Descriptive Inorganic Chemistry textbook (Chapter 8 of the 5th edition). After teaching this method for a number of years, I have anecdotal evidence that students get a better understanding of what the oxidation state is, and when it is suitable to use this type of electron bookkeeping instead of formal charge (for example). In addition to relying on a periodic trend as opposed to a list of rules, this method helps to explain the differences in the oxidation numbers of inequivalent atoms in ions/molecules (such as sulfur in $\ce{S2O3^2-}$), which can't be done with the list-of-rules approach.

For part 2 of the question, the biggest drawback of using the electronegativity approach to determining oxidation numbers is that it requires a knowledge of lewis dot structures. With the current (American) General Chemistry track, students learn about oxidation numbers prior to drawing structures, and therefore must be taught using the list-of-rules approach. I suspect a shuffling of the content to allow students to use the electronegativity approach would have some pedagogical benefits; instead of sitting at the bottom of Bloom's Taxonomy with remembering a list of rules, we can move up to applying the electronegativity trends to analyze the degree of electron deficiency of an atom has in an ion or molecule.

• I do not understand your last paragraph. Are you saying that students learn how to assign oxidation numbers before learning about the structure of the molecule which they are assigning the oxidation numbers to? How can that approach even remotely work? – Martin - マーチン Jul 16 '15 at 9:00
• @Martin-マーチン students typically learn how to determine the oxidation number of atoms in polyatomic ions based on the "rules". They know the chemical formula of the ion, but they have not encountered Lewis dot structures. So they can use the rules to determine that the N in nitrate has an oxidation state of +5 before they learn how to determine whether it is trigonal pyramidal or trigonal planar. – bobthechemist Jul 16 '15 at 11:56
• But then I don't see how considering the electronegativity instead of the rules would change anything. The procedure mostly stays the same. – Martin - マーチン Jul 16 '15 at 12:24
• @Martin-マーチン That's my point, gen chem students learn all that is needed to apply electronegativity to ox numder assignments, just in the wrong order. Therefore, it is only useful (with the current standard US curriculum) later on in the chemistry program. – bobthechemist Jul 16 '15 at 14:29
• I don't want to go into a deep conversation about the topic in the comments, but I really do not understand. What is the wrong order? Do students learn the set of rules to assign ox numbers prior to electronegativity? I just want to understand why there is a problem teaching the far superior approach instead of a fixed and incomplete set of rules. I am not even starting to question that they are taught a concept of bookkeeping of electrons before learning about molecular structures. If you want to talk about this, consider creating a chat room where we can talk. – Martin - マーチン Jul 16 '15 at 14:53