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The chemical formula for water is H2O.
Carbon Dioxide is CO2. Why does oxygen combine in different ratios, in
different compounds? Do Chemistry students need to memorize the chemical formulas
for each of the millions of known compounds? Is there a way to predict the ratio by
which elements will combine in a given situation? Fortunately, that is what oxidation
numbers are for.
You probably recall learning about ions
in Biology. An ion is a charged particle formed when a neutral atom
or group of atoms gain or lose one or more electrons. When a single atom forms an
ion, as in the case of Al+3, it is called a monatomic
ion. When of group of atoms that are covalently bonded together form an
ion, as in the case of NH4+, it is called a polyatomic
ion.
Sometimes ions with opposite charges are
attracted together and will form ionic compounds. Table Salt, NaCl
is such a compound formed from Na+ ions and Cl-
ions. Neutral atoms can also form compounds when they join together, as in the case
of water (H2O). However, since these compounds are not composed of ions,
they are called molecular compounds. You will learn more about
these types of compounds in lesson 5-3.
Regardless as to whether a compound is
made up of ions or not, each atom in the compound has an apparent charge. This
apparent charge, called the oxidation number, represents the charge that
an atom would have if electrons were transferred completely to the atom with the greater
attraction for them in a given situation. These oxidation numbers can be used to
predict the ratio by which atoms will combine when they form compounds.
The following rules help us assign the oxidation number of
elements:
Table 5-2a -
Predicting Oxidation Numbers |
| 1. In free elements (that is, in uncombined state), each
atom has an oxidation number of zero. Ex. In O2, the
oxidation number of each oxygen atom is zero. |
| 2. For ions composed of only one atom, the oxidation number
is equal to the charge on the ion. Ex. The oxidation
number of Ca2+ is +2. |
| 3. All alkali metals (elements in column 1of the periodic
table, with the exception of hydrogen) have an oxidation number of +1. Ex. The oxidation numbers of Li, K, and Na will always be +1. |
| 4. All alkaline earth metals (elements in column 2 of the
periodic table) have an oxidation number of +2. Ex. The
oxidation number of Ba is +2. |
| 5. The oxidation number of Aluminum (Al) is always +3. |
| 6. The oxidation number of oxygen in most compounds (such
as H2O and CO2) is -2. In hydrogen peroxide (H2O2)
and peroxide (O22-) oxygen shows a -1 oxidation number. |
| 7. The oxidation number of hydrogen is +1, except when in
is bonded to a metal as a negative ion, in which case it is -1. Ex. H2O shows hydrogen as +1. NaH shows hydrogen as
-1. |
| 8. When halogens (elements in column 17 on the periodic
table) form negative ions, they will have an oxidation number of -1. Ex.
NaCl and CaCl2 both show chlorine with a -1 oxidation number. |
| 9. In a neutral molecule, the sum of the oxidation numbers
of all of the atoms must be zero. Ex. In H2O,
each hydrogen is +1 and the oxygen is -2. So, (2 x +1) + (-2) = 0. |
| 10. In a polyatomic ion, the sum of oxidation numbers of
all the elements in the ion must be equal to the net charge of the ion. Ex. In the polyatomic ion known as hydroxide (OH-), the
oxygen is -2 and the hydrogen is +1. So, (-2) + (+1) = -1, the same as the charge on
the hydroxide ion (OH-) |
Now, in time you will find it easy to
predict many oxidation numbers, as you become more familiar with the periodic table and
the rules above. Until that time, you should make use of reference tables that list
the oxidation numbers of common ions. Depending on your teacher, he or she may allow
you to make use of such tables for quizzes and exams. For your convenience, I
will provide examples of these tables below. Keep in mind that the table that your
teacher uses may differ from the ones provided below.
Table
5-2b - Oxidation Numbers of Some Common Monatomic Ions |
|
CHARGE |
ION |
+1 |
+2 |
+3 |
+4 |
NONE |
-1 |
-2 |
| Aluminum (Al) |
|
|
X |
|
|
|
|
| Argon (Ar) |
|
|
|
|
X |
|
|
| Barium (Ba) |
|
X |
|
|
|
|
|
| Bromide (Br) |
|
|
|
|
|
X |
|
| Cadmium (Cd) |
|
X |
|
|
|
|
|
| Calcium (Ca) |
|
X |
|
|
|
|
|
| Cesium (Cs) |
X |
|
|
|
|
|
|
| Chloride (Cl) |
|
|
|
|
|
X |
|
| Fluoride (F) |
|
|
|
|
|
X |
|
| Hydride (H) |
|
|
|
|
|
X |
|
| Hydrogen (H) |
X |
|
|
|
|
|
|
| Iodide (I) |
|
|
|
|
|
X |
|
| Lithium (Li) |
X |
|
|
|
|
|
|
| Magnesium (Mg) |
|
X |
|
|
|
|
|
| Neon (Ne) |
|
|
|
|
X |
|
|
| Oxide (O) |
|
|
|
|
|
|
X |
| Potassium (K) |
X |
|
|
|
|
|
|
| Sodium (Na) |
X |
|
|
|
|
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|
| Silver (Ag) |
X |
|
|
|
|
|
|
| Strontium (Sr) |
|
X |
|
|
|
|
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| Sulfide (S) |
|
|
|
|
|
|
X |
| Zinc (Zn) |
|
X |
|
|
|
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|
Now, some
elements show different positive oxidation numbers, in different situations. The
stock system, which you will learn more about in lessons 4-3 and 4-4, uses Roman numerals
to show the oxidation number of the element. For example, Lead(II) is lead with an
oxidation number of +2. Chromium(III) is Chromium with an oxidation number of
+3. The oxidation number for these types of elements will always be positive.
I provide a table below, but once you understand the stock system you will not need the
table any longer.
Table
5-2c - Stock System
Oxidation Numbers of Metals with Multiple Oxidation States |
|
CHARGE |
ION |
+1 |
+2 |
+3 |
+4 |
| Chromium(III) |
|
|
X |
|
| Cobalt(II) |
|
X |
|
|
| Copper(I) |
X |
|
|
|
| Copper(II) |
|
X |
|
|
| Iron(II) |
|
X |
|
|
| Iron(III) |
|
|
X |
|
| Lead(II) |
|
X |
|
|
| Lead(IV) |
|
|
|
X |
| Manganese(II) |
|
X |
|
|
| Mercury(II) |
|
X |
|
|
| Nickel(II) |
|
X |
|
|
| Tin(II) |
|
X |
|
|
Table
5-2d - Oxidation Numbers of Some Common Polyatomic Ions |
|
CHARGE |
ION |
+1 |
+2 |
-1 |
-2 |
-3 |
|
|
|
|
|
|
| Acetate, (CH3COO-) |
|
|
X |
|
|
| Ammonium (NH4+) |
X |
|
|
|
|
| Carbonate (CO32-) |
|
|
|
X |
|
| Chlorate (ClO3-) |
|
|
X |
|
|
| Chromate (CrO42-) |
|
|
|
X |
|
| Cyanide (CN-) |
|
|
X |
|
|
| Dichromate (Cr2O72-) |
|
|
|
X |
|
| Hydroxide (OH-) |
|
|
X |
|
|
| Hypoclorite (ClO-) |
|
|
X |
|
|
| Iodate (IO3-) |
|
|
X |
|
|
| Nitrate (NO3-) |
|
|
X |
|
|
| Nitrite (NO2-) |
|
|
X |
|
|
| Oxalate (C2O42-) |
|
|
|
X |
|
| Perchlorate (ClO4-) |
|
|
X |
|
|
| Permanganate (MnO4-) |
|
|
X |
|
|
| Peroxide (O22-) |
|
|
|
X |
|
| Phosphate (PO43-) |
|
|
|
|
X |
| Silicate (SiO32-) |
|
|
|
X |
|
| Sulfate (SO42-) |
|
|
|
X |
|
| Sulfite (SO32-) |
|
|
|
X |
|
| Tartrate (C4H4O62-) |
|
|
|
X |
|
| Tetraborate (B4O72-) |
|
|
|
X |
|
| Thiosulfate (S2O32-) |
|
|
|
X |
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Please forward all questions, comments and criticisms to Gregory L. Curran.
© Copyright 2004 Fordham Preparatory School, All Rights Reserved.
Last Modified February 07, 2008 |
