| Chem 107 |
Fundamentals of Chemistry |
Fall 2008 |
| Lecture Notes: 29 September |
© R. Paselk 2005 |
|
| |
|
|
| PREVIOUS |
|
NEXT |
Polyatomic ions
In these ions a group of atoms are covalently bound to each other and functions as a single charged particle. You should memorize the following:
- ammonium ion:
- cyanide ion:
- carbonate ion:
- nitrate ion:
- nitrite ion:
- phosphate ion:
- sulfite ion:
- oxalate ion:
- hydroxide ion:
- acetate ion:
- bicarbonate ion:
- sulfate ion:
Compounds
Salts (Ionic Compounds) and Bases
Ions of opposite charge may combine to form neutral compounds. Thus the ions must combine in ratios such that the charges cancel. When the negative ion is hydroxide, the compound is considered a base. Examples:
- potassium chloride:
- sodium hydrogen sulfate (sodium bisulfate):
- aluminum hydroxide:
- iron(III) carbonate:
- iron(III) hydrogen carbonate (iron(III) bicarbonate):
- ammonium phosphate:
- copper(I) oxalate:
- calcium hydroxide:
Acids
Acids are compounds which give hydrogen ions (protons) in solution. There are two common inorganic acid types in terms of nomenclature:
- Hydro-( )-ic acids: If hydrogen combines with a nonmetallic element the resulting acid is named by adding the prefix hydro- and replacing -ide by -ic. Examples:
- Oxo acids: When non-metallic elements react with oxygen the resulting products often react with water or form ions which can react with protons to from acids. These oxo acids are named by replacing the -ate suffix with -ic acid or -ite suffix with -ous acid. Examples:
- carbonic acid:
- sulfuric acid:
- acetic acid:
- nitric:
- Chloroacids (Hydrochlorous acid - Perchloric acid)
- The Chemistry Department Table of Common Acids is a useful summary of the acids you should be familiar with.
Special names
These compounds don't follow the rules, but have been in common use so long they keep their traditional names. Examples:
- Water:
- methane:
- ammonia:
- hydrogen peroxide:
Oxidation Numbers
For simple elemental ions it is easy to determine the charge on an atom, but in many other circumstances this is not the case. In order to name compounds we frequently need this information which is obtained from oxidation numbers.
Oxidation numbers are in essence an electronic accounting method in which electrons are assigned to a particular atom in a bond or interaction. As such they give an approximate picture of where electrons actually reside in compounds. We will find this information very useful later when we look at particular types of chemical reactions. Oxidation numbers are essential for nomenclature.
Oxidation numbers are most readily assigned using a simple set of rules:
Rules for Assigning Oxidation Numbers
- In the formula for any substance the sum of the oxidations numbers of all the atoms in the formula is equal to the charge shown. Thus:
- For elements, such as Ar, O2, S8, etc. in the uncombined state the oxidation number for each atom must be 0, since no charge is shown and the atoms are equal to each other.
- For monoatomic ions the oxidation number equals the charge.
- For a compound the sum of the oxidation numbers of the atoms equals 0.
- For a polyatomic ion the sum of the oxidations numbers of the atoms equals the charge on the ion.
- In compounds fluorine is always assigned an oxidation number of -1.
- In compounds oxygen is usually assigned an oxidation number of -2.
- Exception 1: in peroxides it is -1 while in superoxides it is -1/2. These will generally be obvious due to other rules (or the names).
- Exception 2: in combination with fluorine oxygen can be positive due to Rule 2 above, thus for OF2 oxygen is assigned an oxidation number of +2.
- In compounds hydrogen is usually assigned an oxidation number of +1
- Exception: in metallic hydrides hydrogen is assigned an oxidation number of -1. These exceptions will be fairly obvious: NaH, CaH2, etc.
- Alkali metals in compounds will always (for our class) be assigned an oxidation number of +1.
- Alkaline-earth metals in compounds will always (for our class) be assigned an oxidation number of +2
- Aluminum will always (for our class) be assigned an oxidation number of +3, other elements in this Group will usually be assigned an oxidation number of +3.
|
Examples:
- NO3-
- NO2-
- PO43-
- SO32-
- SO42-
- C2O42-
- C2H3O2-
- HCO3-
- MnO4-
- MnO2
- HClO4
- HClO3
- HClO2
- HClO
Finally, note that in writing formulae, the element with the more positive oxidation number comes first. There are, of course, a few exceptions, the most well known being ammonia: NH3 (by the rules it should be H3N).
© R A Paselk
Last modified 30 September 2008
