In chemistry, a compound is a substance that results from a combination of two or more different chemical element s, in such a way that the atom s of the different elements are held together by chemical bonds that are difficult to break. These bonds form as a result of the sharing or exchange of electron s among the atoms. The smallest unbreakable unit of a compound is called a molecule
Examples of compounds:
- water (H2O)
- table salt (NaCl)
- sucrose (table sugar, C12H22O11
The relationship is simple.
Atoms are what all matter are ultimately made up of. Atoms are the smallest units of an element.
Elements are substances composed of all the same type of atoms, and have specific chemical properties. Aluminum for example contains only Aluminum atoms, and no other, and has chemical properties specific to Aluminum.
Molecules are combinations of atoms that are not necessarily all the same element. Sometimes they are the same element, like air molecules. Air molecules are a mix of pairs of Nitrogen, and pairs of Oxygen. Although the pairs of atoms are the same element, they are more than one atom so they are molecules. Water molecules are made of Hydrogen atoms and Oxygen atoms, i.e. different elements.
Compounds are combinations of elements into new substances, like water. Water combines the elements of Hydrogen and Oxygen and has chemical properties distinct from the elements it’s made of.
Long before chemists knew the formulas for chemical compounds, they developed a system of nomenclature that gave each compound a unique name. Today we often use chemical formulas, such as NaCl, C12H22O11, and Co(NH3)6(ClO4)3, to describe chemical compounds. But we still need unique names that unambiguously identify each compound.
Common Names
Some compounds have been known for so long that a systematic nomenclature cannot compete with well-established common names. Examples of compounds for which common names are used include water (H2O), ammonia (NH3), and methane (CH4).
Naming Ionic Compounds
(Metals with Non-metals)
The names of ionic compounds are written by listing the name of the positive ion followed by the name of the negative ion.
| NaCl | sodium chloride | ||
| (NH4)2SO4 | ammonium sulfate | ||
| NaHCO3 | sodium bicarbonate |
We therefore need a series of rules that allow us to unambiguously name positive and negative ions before we can name the salts these ions form.
Naming Positive Ions
Monatomic positive ions have the name of the element from which they are formed.
| Na+ | sodium | Zn2+ | zinc | |
| Ca2+ | calcium | H+ | hydrogen | |
| K+ | potassium | Sr2+ | strontium | |
Some metals form positive ions in more than one oxidation state. One of the earliest methods of distinguishing between these ions used the suffixes -ous and -ic added to the Latin name of the element to represent the lower and higher oxidation states, respectively.
| Fe2+ | ferrous | Fe3+ | ferric | |
| Sn2+ | stannous | Sn4+ | stannic | |
| Cu+ | cuprous | Cu2+ | cupric |
Chemists now use a simpler method, in which the charge on the ion is indicated by a Roman numeral in parentheses immediately after the name of the element.
| Fe2+ | iron(II) | Fe3+ | iron (III) | |
| Sn2+ | tin(II) | Sn4+ | tin(IV) | |
| Cu+ | copper(I) | Cu2+ | copper(II) |
Polyatomic positive ions often have common names ending with the suffix -onium.
| H3O+ | hydronium |
| NH4+ | ammonium |
Naming Negative Ions
Negative ions that consist of a single atom are named by adding the suffix -ide to the stem of the name of the element.
Common Polyatomic Negative Ions
| -1 ions | ||||
| HCO3– | bicarbonate | HSO4– | hydrogen sulfate (bisulfate) | |
| CH3CO2– | acetate | ClO4– | perchlorate | |
| NO3– | nitrate | ClO3– | chlorate | |
| NO2– | nitrite | ClO2– | chlorite | |
| MnO4– | permanganate | ClO– | hypochlorite | |
| CN– | cyanide | OH– | hydroxide | |
| -2 ions | ||||
| CO32- | carbonate | O22- | peroxide | |
| SO42- | sulfate | CrO42- | chromate | |
| SO32- | sulfite | Cr2O72- | dichromate | |
| S2O32- | thiosulfate | HPO42- | hydrogen phosphate | |
| -3 ions | ||||
| PO43- | phosphate | AsO43- | arsenate | |
| BO33- | borate |
Naming Polyatomic Ions
At first glance, the nomenclature of the polyatomic negative ions in the table above seems hopeless. There are several general rules, however, that can bring some order out of this apparent chaos.
The name of the ion usually ends in either -ite or -ate. The -ite ending indicates a low oxidation state. Thus,the NO2– ion is the nitrite ion.
The -ate ending indicates a high oxidation state. The NO3– ion, for example, is the nitrate ion.
The prefix hypo– is used to indicate the very lowest oxidation state. The ClO- ion, for example, is the hypochlorite ion.
The prefix per– (as in hyper-) is used to indicate the very highest oxidation state. The ClO4– ion is therefore the perchlorate ion.
There are only a handful of exceptions to these generalizations. The names of the hydroxide (OH–), cyanide (CN–), and peroxide (O22-) ions, for example, have the -ide ending because they were once thought to be monatomic ions.
Naming Simple Covalent Compounds
( Non-metals with non-metals )
Oxidation states also play an important role in naming simple covalent compounds. The name of the atom in the positive oxidation state is listed first. The suffix -ide is then added to the stem of the name of the atom in the negative oxidation state.
| HCl | hydrogen chloride |
| NO | nitrogen oxide |
| BrCl | bromine chloride |
As a rule, chemists write formulas in which the element in the positive oxidation state is written first, followed by the element(s) with negative oxidation numbers.
The number of atoms of an element in simple covalent compounds is indicated by adding one of the following Greek prefixes to the name of the element.
| 1 mono- | 6 hexa- | ||
| 2 di- | 7 hepta- | ||
| 3 tri- | 8 octa- | ||
| 4 tetra- | 9 nona- | ||
| 5 penta- | 10 deca- |
The prefix mono– is seldom used because it is redundant. The principal exception to this rule is carbon monoxide (CO).
Naming Acids
Simple covalent compounds that contain hydrogen, such as HCl, HBr, and HCN, often dissolve in water to produce acids. These solutions are named by adding the prefix hydro– to the name of the compound and then replacing the suffix -ide with -ic. For example, hydrogen chloride (HCl) dissolves in water to form hydrochloric acid; hydrogen bromide (HBr) forms hydrobromic acid; and hydrogen cyanide (HCN) forms hydrocyanic acid.
Many of the oxygen-rich polyatomic negative ions in Table 2.1 form acids that are named by replacing the suffix –ate with -ic and the suffix -ite with -ous.
| Acids containing ions ending with ide often become | hydro -ic acid | |||
| Cl– | chloride | HCl | hydrochloric acid | |
| F– | fluoride | HF | hydrofluoric acid | |
| S2- | sulfide | H2S | hydrosulfuric acid | |
| Acids containing ions ending with ate usually become | -ic acid | |||
| CH3CO2– | acetate | CH3CO2H | acetic acid | |
| CO32- | carbonate | H2CO3 | carbonic acid | |
| BO33- | borate | H3BO3 | boric acid | |
| NO3– | nitrate | HNO3 | nitric acid | |
| SO42- | sulfate | H2SO4 | sulfuric acid | |
| ClO4– | perchlorate | HClO4 | perchloric acid | |
| PO43- | phosphate | H3PO4 | phosphoric acid | |
| MnO4– | permanganate | HMnO4 | permanganic acid | |
| CrO42- | chromate | H2CrO4 | chromic acid | |
| ClO3– | chlorate | HClO3 | chloric acid | |
| Acids containing ions ending with ite usually become | -ous acid | |||
| ClO2– | chlorite | HClO2 | chlorous acid | |
| NO2– | nitrite | HNO2 | nitrous acid | |
| SO32- | sulfite | H2SO 3 | sulfurous acid | |
| ClO– | hypochlorite | HClO | hypochlorous acid |
Complex acids can be named by indicating the presence of an acidic hydrogen as follows.
| NaHCO3 | sodium hydrogen carbonate (also known as sodium bicarbonate) |
| NaHSO3 | sodium hydrogen sulfite (also known as sodium bisulfite) |
| KH2PO4 | potassium dihydrogen phosphate |
Valency
The valency of an atom is the number of single chemical bonds that it can make (in the case of a covalently bonding substance) or the number of electrical charges that it carries (for an ion). Notice that once again the nature of the substance in question requires that the definitions be adapted appropriately. The concept of valence can be used to find the formula of a compound from the valencies of its constituent elements, or to find the valency of an elements within a compound of known formula.
Every atom within a substance is assigned a valency number that is either positive or negative. The total sum of all of the valencies within a formula unit is zero
Using valencies
Once the valencies of a few elements are known it becomes a simple matter to construct the formula of unknown compounds using the valency method. Remember that the sum of the valencies of all of the atoms in the compound must equal zero.
Where an atom may have either positive or negative valency, it is negative if it is the more electronegative element in the compound and positive if not.
| Example: From the water molecule above we know that the valency of hydrogen is +1. If the valency of nitrogen in ammonia is -3 then we can construct the formula of ammonia thus: We need enough hydrogens to cancel out the -3 valency of nitrogen. Each hydrogen = +1 therefore we need three hydrogen atoms. The formula of ammonia = NH3 |
Working with ions
When using valencies to work out the formula of an ion we have to remember the final charge on the ion must equal the sum of the valencies, taking into account whether the valency of each atom is negative or positive.
| Example: Find the formula of the sulfate (2-) ion given that the valency of the sulfur atom is +VI and the valency of the oxygen atom is -II Oxygen always has negative valencies (unless bonded to fluorine) There is one sulfur atom with a valency of +6 and overall the ion has a valency of -2 Therefore +6 +(xO) = -2 Therefore (xO) = -2 -6 = -8 each O =-2 therefore there are four oxygen atoms in the ion Formula of the sulfate ion = SO42- |
EVALUATION
1.write the symbols and the valencies of the following:
i. Iron ii. potassium iii. Oxygen iv. Chlorine
2. What is valency?
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