Course Outline


Binary Compounds



 

After completing this unit you should be able to:
  • understand the classification of elements as metals, non-metals, or metalloids
  • recognize when a pair of elements will form an ionic compound, and name it
  • recognize when a pair of elements will form a molecular compound

Most elements, including all of those in the s, d, and f-blocks, are classified as metals.  Under standard conditions (25 oC and 1 atm pressure) they are shiny, conduct heat and electricity, and all except mercury are malleable and ductile solids (i.e. they can be shaped into sheets and wires).  Many p-block elements do not have these properties, so they are classified as non-metals.  A few elements have intermediate properties and are called metalloids.  This periodic table shows where these elements are located.

 

1A
(1)

2A
(2)

(3 - 12)

3A
(13)

4A
(14)

5A
(15)

6A
(16)

7A
(17)

8A
(18)

1

METALS METALLOIDS NON-METALS 1
H
2
He

2

3
Li
4
Be
  5
B
6
C
7
N
8
O
9
F
10
Ne

3

11
Na
12
Mg
  13
Al
14
Si
15
P
16
S
17
Cl
18
Ar

4

19
K
20
Ca
21 - 30 31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr

5

37
Rb
38
Sr
39 - 48 49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe

6

55
Cs
56
Ba
57 58 - 71 72 - 80 81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn

7

87
Fr
88
Ra
89 90 - 103 104 - 112            

Ionic Compounds

Binary compounds are those composed of the atoms of two elements.  When one of the elements is a metal and the other is a non-metal the resulting compound is composed of electrically charged atoms called ions.  Positively charged ions are called cations, and negatively charged ions are anions.

Atoms of main group metallic elements have only a small number of electrons in their valence shell.  When these electrons are given up, the remaining electrons have a stable configuration like the preceding noble gas, and the ion has a positive charge (due to the excess protons) equal to the number of electrons lost.  Group 1A elements form cations with charge 1+, group 2A cations are 2+, and aluminum in group 3A forms Al3+.  The charge is written as a superscript after the symbol.

Atoms of non-metallic elements are one or a few electrons short of the electron configuration of the next noble gas in the periodic table.  Ionic compounds are formed when atoms of a metallic element give up electrons to atoms of a non-metallic element.  Group 7A elements (and hydrogen) gain one electron to form anions with charge 1-, Cl- (chloride) for example.  Anions are named by replacing the last few letters of the name of the element with the ending -ide.  Group 6A elements gain two electrons (oxide, O2- and sulfide, S2-), and nitrogen gains three electrons to form the anion nitride, N3-Ionic compounds are named using the name of the metal followed by the name of the anion, e. g. calcium chloride.

The formula for an ionic compound is the lowest whole number ratio of ions that gives a total charge of zero.  Calcium chloride has the formula CaCl2, i.e. the charge on one Ca2+ ion is balanced by the charges on two Cl- ions.  Aluminum oxide has the formula Al2O3, with the charge on two Al3+ ions being exactly balanced by the charge on three O2- ions. 

Some transition metals form just one cation (Ag+, Zn2+), while many others can form two or more different cations.  Copper (Cu+ and Cu2+) and iron (Fe2+ and Fe3+), as well as the main group elements tin (Sn2+ and Sn4+) and lead (Pb2+ and Pb4+), are common examples.  To name compounds of these metals, the charge on the cation must be included as a roman numeral immediately following the name of the metal.  For example, the compound FeCl2 must include the cation Fe2+ to balance the charges of two chloride ions, so the compound is called iron(II) chloride.  There is also a compound iron(III) chloride, with formula FeCl3.  (These two compounds are commonly called ferrous chloride and ferric chloride respectively).

Ionic compounds tend to be crystalline solids with high melting points.  Many rocks and minerals are ionic compounds, and other well-known examples include sodium chloride (table salt) and tin(II) fluoride (which was for many years added to drinking water and toothpaste to prevent tooth decay).

Molecular Compounds

When atoms of non-metallic elements combine to form compounds, neither element can achieve a stable electron configuration by giving up electrons, so ions are not formed.  Instead, the atoms share electrons so that they have at least a share of eight electrons in their valence shell (or two for hydrogen).  Hydrogen and the halogens need to share one pair of electrons, group 6A elements need to share two pairs, group 5A elements share three pairs, and the group 4A elements carbon and silicon share four pairs.

A shared pair of electrons is called a covalent bond, and groups of atoms held together by covalent bonds are called molecules.

The formulas of binary compounds formed from non-metallic elements are written with the element nearest fluorine second, and in the name this element is given the ending -ide, unless (as frequently occurs for this kind of compound) a common name is used.  Methane (CH4), ammonia (NH3), water (H2O), and hydrogen fluoride (HF) are typical.  The number of atoms of each element present in a molecule is indicated in the formula by subscripts.  When a pair of non-metallic elements can form two or more compounds, Greek prefixes (mono-, di-, tri-, tetra-, penta-, hexa-, etc.) are used to indicate the number of atoms of each element in one molecule.  Well-known examples include carbon monoxide, CO, and carbon dioxide, CO2.  Another example is dinitrogen monoxide, N2O, commonly called nitrous oxide or laughing gas.  Notice that the prefix mono- is not used when there is only one atom of the first element in the name, but it is used when there is one atom of the second element. 

Formulas of molecular compounds will not necessarily be the lowest whole number ratio of the elements.  Hydrogen peroxide, a corrosive liquid used in bleach and rocket fuel, consists of molecules having two atoms of each element, so the formula is H2O2.

Formulas for molecular compounds are also limited in that they do not show how the atoms are attached to each other in the molecule.  This can be a serious problem when two different compounds have the same formula, a situation that arises commonly in the chemistry of carbon compounds (organic chemistry).  When necessary, molecules can be drawn out more completely, using lines to represent covalent bonds between atoms, and dots to represent unshared electron pairs in the valence shell.  These are called Lewis dot structures.  Unshared electron pairs are sometimes left out for simplicity.  Here are some examples:

Simple molecular compounds are gases or liquids at room temperature.  Even those that are solids have melting points much lower than typical ionic compounds.

Elements

The atoms of elements in their standard state also combine in ways that allow them to achieve a stable electron configuration.  Hydrogen, the halogens, oxygen, and nitrogen, all exist as diatomic molecules, X2, with one or more covalent bonds between the two atoms.  One naturally occurring form of carbon, diamond, has every one of its atoms covalently bonded to its four nearest neighbors.   A diamond is essentially one huge molecule containing billions and billions of atoms!  In between these two extremes are elements like phosphorus and sulfur that most commonly form molecules with four and eight atoms respectively.

Atoms of metallic elements cannot form stable electron configurations by sharing electrons in covalent bonds, so their electrons are shared in a different way.  Essentially all of their valence electrons are contributed to a 'sea' (the valence or conduction band) and shared among all of the atoms.  This is what gives a metal those uniquely metallic properties.  Metals do not form compounds with other metals.  Instead they can be mixed in any proportions (forming alloys), with all of the atoms contributing electrons to the communal 'sea'.

 

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