Has NH3 a charge
Oxidation numbers / oxidation states
The oxidation numbers are pure formal auxiliary variables and should not be confused with real cargoes. To understand the derivation of the oxidation numbers, the definition is the Electronegativity important.
Note: Electronegativity (EN) is a relative measure of attracting electrons in a bond.
The electronegativity according to Pauling is assigned to each element symbol and can usually be found in the periodic table. Let's look at the connection HCl (Hydrochloric acid). The electronegativity of chlorine amounts to 3,16 and those of hydrogen2,2. This results in an electronegativity difference (ΔEN) between the attachment partners. This is indicated by the partial charges (δ + or δ-). Thus, chlorine has a higher tendency to attract the binding electrons than hydrogen. To determine the oxidation number, the binding electrons are completely assigned to the more electronegative binding partner, in this case chlorine (see Figure 3).
By doing this, the chlorine atom in this compound formally receives an electron (e-) more than it had before, and the hydrogen atom is missing an electron (e-). Therefore the hydrogen atom has the oxidation number + I and the chlorine atom is the oxidation number –I. The sum of the oxidation numbers is 0 because the HCl molecule is a neutral particle to the outside world. The oxidation numbers can be written as Roman or Arabic numerals. Anyone can handle this as he wants. Many compounds have been viewed in this way and the following rules for determining oxidation numbers emerged, which are listed in Table 1.
Of course, there are many compounds made up of more elements than those mentioned in Table 1. Nevertheless, one can determine the oxidation number for each binding partner of a compound if one keeps in mind that the sum of the oxidation numbers must give the charge of the entire molecule. If there is a neutral molecule, then of course no “charge” must remain.
Note: Oxidation numbers are fictitious charges and represent an auxiliary variable in redox reactions. The basis of the oxidation numbers is electronegativity (EN).
Let's look at another molecule and determine the oxidation numbers of all binding partners (see Figure 4).
We know from Table 1 that bound hydrogen usually has the oxidation number + I. We can apply this knowledge here. Since there are three hydrogen atoms in the ammonia molecule, this results in a formal charge of + III. Since ammonia is a neutral molecule, the nitrogen atom (N) has to compensate for this formal charge of + III with its oxidation number. Therefore N has the oxidation number -III. In total, this results in an oxidation number of 0.
Let us analyze another example in Figure 5. The ion is called permanganate (MnO4-) and the aqueous solution of this ion is deep purple.
Table 1 shows that bound oxygen usually has the oxidation number –II. However, since 4 oxygen atoms are bound in the permanganation, we get a formal charge of -VIII. This charge has to go through the manganese (Mn) a negative charge can be compensated. Thus the Mn has the oxidation number + VII and in total the molecule is simply negatively charged to the outside.
Video: Oxidation Numbers / Oxidation Levels
There are, however Exceptions of the rules in Table 1.Let's look at the compounds hydrogen peroxide (H2O2) and lithium hydride (LiH) at.
We can see in Fig. 6 that the bound hydrogen here has the oxidation number + I, as is usually the case. The oxygen atoms, however, have the oxidation number –I (not –II as usual). The reason for this is the oxygen-oxygen bond. There we have no electronegativity difference (ΔEN) and thus we cannot assign the binding electrons to any binding partner.
As the next exception, let's look at lithium hydride in Fig. 7. The electronegativity of lithium is 0.98, that of hydrogen 2.2. hydrogen is in this case more electronegative than its binding partner. As a result, it has the oxidation number –I in this connection (not + I as usual).
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