Bond

I

Introduction

Bond, strong force between atoms that holds them together in molecules. When two or more atoms are close enough, an attractive force between the electrons of individual atoms, and the nuclei of one or more of the other atoms, can result. If this force is large enough to keep the atoms together, a chemical bond is said to have formed. All chemical bonds result from the simultaneous attraction of one or more electrons by more than one nucleus.

When a bond is formed between atoms, energy is normally given out because the resulting compound is at a lower energy level, and therefore more stable, than the original atoms. Similarly, energy must normally be put into a compound, for example by means of heating or an electric current, in order to break existing bonds between atoms. The energy required to break a bond is called its bond energy (or bond enthalpy), and is a useful quantity when the stability of compounds is being studied.

The most convenient model of the atom to use in order to understand chemical bonding is the Bohr model (see Niels Bohr), in which the electrons surrounding the nucleus are considered to occur in distinct shells or energy levels. The first two elements, hydrogen and helium, have one and two electrons respectively, both in the first shell, but the third element, lithium, has two electrons in its first shell, and one in a second shell. The next seven elements in their order in the periodic table (lithium to neon) have two to eight electrons in their second shell, but the next element, sodium, has one electron in its third shell, and the pattern repeats in a similar way thereafter. Note that the number of electrons in the outermost shell of any element corresponds with the number of the group in the periodic table to which the element belongs. This arrangement is usually represented thus:

• Hydrogen (H) 1
• Helium (He) 2
• Lithium (Li) 2.1
• Beryllium (Be) 2.2
• Flourine (F) 2.7
• Neon (Ne) 2.8
• Sodium Na) 2.8.1
• Magnesium (Mg) 2.8.2
• Chlorine (Cl) 2.8.7
• Argon (Ar) 2.8.8
• Potassium (K) 2.8.8.1

In this sequence, the noble gases, the least reactive of all the elements, all have eight electrons in their outermost shell (except for helium, which only has two). It appears that in all forms of chemical bonding there is a strong tendency for atoms to reach this stable electronic state by gaining, losing, or sharing electrons.

II

Covalent Bond

Covalent bonds are formed when atoms share electrons, one from each atom in a single bond, to form electron pairs, usually making their outermost shells up to eight electrons by this means. They are most frequently formed between pairs of non-metallic elements. Non-metallic elements usually have from four to eight electrons in their outermost shells, the so-called valency electrons, which are used for chemical bonding. In any given “full” shell of eight electrons, the electrons occur in four pairs, but in incomplete shells, the electrons exist singly where possible. Thus, the carbon atom, with four electrons in its outermost shell, has four unpaired electrons; nitrogen, with five valency electrons, has one pair and three single electrons; oxygen, with six, has two pairs and two singles, and so on.

In the covalent compound between carbon and hydrogen in methane, each hydrogen atom shares its one electron with one of the four unpaired valency electrons to make a single electron-pair bond. The methane molecule consists of four hydrogen atoms attached by covalent bonds to the central carbon atom, as illustrated in the table. By this means, the carbon atom in methane has a complete shell of eight electrons, and each hydrogen atom similarly achieves the electronic structure of the noble gas, helium, by sharing electrons. The compound between nitrogen, with three unpaired valency electrons, and hydrogen (ammonia, NH₃) consists similarly of three hydrogen atoms linked by three covalent bonds to a single nitrogen atom. Other simple examples are water (H₂O), hydrogen chloride (HCl), and phosphorus trichloride (PCl₃). Covalent bonds are usually represented by a single straight line, for example, H8Cl.

Sometimes, atoms of elements form covalent bonds with other atoms of the same element. Thus two chlorine atoms form the chlorine molecule, Cl₂, by sharing their unpaired electrons. In the case of oxygen (O₂), there are two unpaired valency electrons in each atom, so that two electron-pair bonds are formed between the two atoms to complete their octet of electrons, a double bond.

Covalent bonds can also be formed in such a way as to form a giant molecule, such as happens in diamond. Here, each of the four valency electrons of a carbon atom is shared with one of the valency electrons of another carbon atom, so that every carbon atom in the structure has four different carbon atoms bonded to it, as shown in the table above.

Covalent bonds are directional, that is to say they have specific directions in space relative to one another, and this gives the molecules containing them a particular shape. The molecule of methane, shown in the table, has the shape of a regular tetrahedron, and most covalent molecules are based on this shape, with an angle between the bonds of 109.5°, although the precise bond angles depend on the nature of the atoms in the molecule. In a water molecule, for example, hydrogen atoms occupy two of the corners of the tetrahedron with the oxygen atom at its centre, whereas the “lone pairs” of electrons on the oxygen atom occupy the other two corners; this gives an angle between the two hydrogen atoms in water of 104.5°. Some of the properties of water, noted later in this article, depend on the fact that the water molecule is not formed in a straight line, but is bent at an angle.

The typical energy of a covalent bond is between 100 to 500 kJ/mole.

In some covalent compounds, for example, benzene (C₆H₆), it is not possible to allocate a specific structure to the molecule which correctly represents the properties of the compound. The electrons do not appear to occur where they contribute only to particular bonds, but are delocalized. This is the phenomenon of resonance (see Benzene).

A

Properties

The typical properties of compounds containing covalent bonds depend on the size of their molecules.

In simple molecules, the molecules are held together by strong covalent bonds, but the bonds between different molecules are relatively weak, and therefore easily broken. Many compounds containing covalent bonds, therefore, are liquids or gases at room temperature, although at lower temperatures they form crystalline solids; those that are solids at room temperature have relatively low melting and boiling points.

These compounds are frequently insoluble in water, as the mechanism for dissolving described for ionic compounds cannot occur. However, water is also capable of dissolving covalent compounds which form hydrogen bonds, because instead of the ionic attachment described below, the water molecules can attach themselves to the molecules of the covalent compounds, with a similar result.

Non-polar solvents dissolve some covalent compounds, but usually when their structures are sufficiently similar for bonding to occur between them.

Solutions of covalent compounds do not normally contain ions, and therefore do not conduct electricity.

In giant molecules, the strong covalent bonds occur throughout the crystal, which is really all one molecule. This gives the name, giant molecule. The bonds are directional, giving the solid its specific shape. The crystal is hard, and has very high melting and boiling points because the covalent bonds themselves have to be broken (theirs are not merely weak intermolecular forces).

III

Ionic Bond

Ionic bonds are formed through the electrostatic attraction between two oppositely charged ions. This type of bond normally occurs between metallic and non-metallic elements. Metals typically have few valency electrons, and occur in groups I, II, and III of the periodic table; reactive non-metals typically have more electrons in their outermost shells, and occur in groups IV, V, VI, and VII. An ionic bond occurs when a metal loses all its valency electrons, leaving a positively charged ion with a noble gas structure, and a non-metal gains the number of electrons it requires to give it a noble gas structure, and becomes a negative ion. A sodium atom, for example, has the electronic structure of 2.8.1; when it loses an electron, it becomes a positively charged sodium ion, Na⁺, with the electronic structure of neon, 2.8. A chlorine atom has the structure 2.8.7; when it gains an electron, it becomes a negatively charged ion Cl^-, with the electronic structure of argon, 2.8.8. Similarly, a calcium atom (2.8.8.2) loses two electrons to become a Ca²⁺ ion, while an oxygen atom (2.6) gains two electrons to become an O^2- ion, and so on.

The ionic bond occurs by means of the attraction between these charged particles, but because there are, of course, millions and millions of them in any given sample of sodium chloride, they do not just pair off to make simple “molecules”. Instead, the positive ions cluster as close as they can round the negative ions, and similarly the negative ions cluster round the positive ions. The net result of this process is that a regular lattice of ions is formed, whose structure, and hence the shape of the resulting ionic crystal, depends on the sizes of the ions and their relative charges. There must, in the end, be an equal number of ions of each charge to ensure that the final structure is neutral, so it is justifiable to write the formula, NaCl, for example, for an ionic compound. However, it must be noted that there is no such entity as a single molecule of an ionic compound, in the way that there can be a single molecule of a covalent compound. A useful way to represent a compound formed by an ionic bond is to show the charges on the ions, Na⁺Cl^-.

Lattice energies, the equivalent of bond energies for covalent compounds, are of the order of 600 to 1,000 kJ/mole, for compounds containing singly charged ions, and higher if either or both of the ions has a multiple charge.