Ionic Radii

Before continuing we should consider the radii of the ions (common cations/common anions) formed when electrons are lost or gained by neutral atoms. The cation formed when one or more electrons are lost has a smaller ionic radius compared to the neutral atom. The loss of electrons increases the effective nuclear charge of the electrons remaining on the ion and they experience a greater attraction towards the nucleus. Adding of one or more electrons to a neutral atom, to form an anion, decreases the effective nuclear charge of the valence electrons and the ionic radius is greater compared to the neutral atomic radius. For isoelectronic ions the greater the positive charge the smaller the ionic radius. While for the isoelectronic ions the greater the negative charge the larger the atomic radius.

Same number of valence electrons in Be and B⁺, however the valence electrons in B⁺ 'see' a greater effective nuclear charge and are attracted closer to the nucleus––smaller radius than Be.

Same number of valence electrons in Ne and F^–, however the valence electrons in F^– 'see' a smaller effective nuclear charge and are not held as tightly by the nucleus––larger radius than Ne.

Some elements which form ionic compounds are too far removed from a noble gas to readily achieve the rare gas configuration. Members of the transition metals series would have to lose a large number of electrons to become isoelectronic with the nearest noble gas. Most transition metals form cations of 2+ and 3+. This suggests that stable ions can exist which do not duplicate the electronic arrangements of the noble gases. Some of the ions can be characterized by filled subshells, for example, Zn²⁺ [Ar]3d¹⁰, Ag⁺ [Kr]3d¹⁰ and Pb²⁺ [Xe]6s²5d¹⁰. It is therefore more difficult to predict most stable ions in the transition metals. Instead their chemistry is a collection of a large variety of ionic compounds.

Covalent Compounds

As important and varied as ionic compounds are, there are many compounds which do not demonstrate the high melting points, high water solubility and electrical conductivity of ionic compounds. This other group of compounds consists of examples that are gaseous, liquids and solids at room temperature, and generally they have low melting or boiling points. Frequently they are insoluble in water. They do not conduct electricity in the liquid state, or when soluble in water, do not conduct electricity in aqueous solution. Compounds that do not conduct electricity are called nonelectrolytes. There is no evidence to indicate that the elements in these compounds are ionic. Rather there is considerable experimental evidence to suggest that such compounds are made of discrete molecules, hence compounds in this group are called molecular compounds.

The forces that hold the atoms together in molecular compounds can not be understood on the basis of oppositely charged ions. So it was that Gilbert Lewis proposed that the strong attractive force between two atoms in a molecule reselted from a covalent bond, formed by sharing of a pair of electrons between the atoms in the bond.

In hydrogen, H₂, as the two hydrogen atoms approach one another their spherical 1s orbitals begin to overlap. Each electron occupying the space around the two nuclei. Each electron is attracted simultaneously by each nuclei. The attraction that bonds the electrons to both nuclei is the force holding the atoms together. Thus while ions do not exist in covalent compounds the bond can be regarded as arising from the attraction of oppositely charged particles–nuclei and electrons.

We can represent the formation of the covalent bond in hydrogen by writing the Lewis electron–dot formula for the atoms and the molecule.

The two electrons between the two hydrogen nuclei represent a covalent bond. And the two electrons are referred to as an electron pair. Each hydrogen atom can be thought of as sharing the pair of electrons. When we think in these terms we note that each atom has a 1s² electron configuration, isoelectronic with the next noble gas–helium.

The formation of the covalent bond in gaseous HF can be described in similar terms;

The fluorine atom and the hydrogen atom each require one electron to achieve a configuration isoelectronic with a noble gas. Helium for hydrogen and neon for fluorine. The two atoms share the electron pair to achieve their respective noble gas configurations. The other three pairs of electrons around the fluorine atom are called nonbonding or lone pair electrons.

We can understand the formulas of a large number of covalent compounds by writing or drawing the Lewis electron–dot formulas for the compounds.

(Use lines to denote the covalent bond and dots to denote the electron pairs.)

Notice that in writing our formulas that F, N, O and C all have eight electrons around it. The tendency of atoms in a molecule to share electrons to have a total of eight is called the octet rule. Many compounds follow the octet rule, and some do not.

So far all the examples we've discussed involved atoms sharing two electrons. Two electron bonds are called single bonds. However, it is possible for two atoms to share two pairs of electrons to form double bonds and even three pairs of electrons to form triple bonds. Atoms such as C, N, O and S form double bonds in certain instances, while C and N exhibt examples of triple bonds.

Although the covalent bonds in dihydrogen, H₂, and hydrogen fluoride, HF, each involve a pair of electrons being shared, there are some important differences in the way inwhich the electrons are shared. In hydrogen it can be noted that the electrons spend an equal amount of time near each nucleus. However, it can not be said that the electron pair shared between hydrogen and fluorine in HF is equally shared. In fact, the electrons spend more time near the fluorine nucleus than near the hydrogen nucleus. Resulting in the fluorine appearing to have some extra negative charge and the hydrogen some extra positive charge.

The covalent bond in HF is described as a polar covalent bond because the electrons spend an unequal amount of time on the two nuclei. The bond in hydrogen is called a nonpolar bond. It should be noted that the polar bond can be viewed as an intermediate case between the equal sharing of electrons in hydrogen and the transfer of electrons that arises in ionic bonds.

Atoms form compounds by losing, gaining or sharing enough electrons to achieve the outer electron configuration of a noble gas. The combining capacities of atoms are a consequence of the proportions in which they must associate to achieve noble gas configurations.

ELECTRONEGATIVITY

Electronegativity is a concept which is used to help understand which atom in a chemical bond will have the greatest attraction for the bonding pair of electrons. Electronegativity is a measure of the ability of an atom in a chemical bond to attract electrons to itself.

From the figure one can see the important trend in electronegativity across a period and down a group. In general the electronegativity increases going across a period, and decreases going down a group. The larger the electronegativity value the greater the attraction of the electron by the atom. The absolute value of the difference in electronegativity of two atoms sharing a pair of electrons is a measure of the degree of polarity of the chemical bond. When the difference is small the bond is nonpolar. When it is large the bond is polar, if the difference is very large the bond is ionic.