Chemical Bonding

Overview

Chemical bonds are formed when some force of attraction exists between atoms, ions or molecules. If the bond involves two or more different elements, the result is a chemical compound. There are many different types of chemical bond. Some bonds are very strong, while others are relatively weak. The properties of a particular bond, together with the manner in which it is created, will depend upon the precise electronic configuration of the atoms or molecules involved in the bonding process.

Regardless of the type of chemical bond that is formed, however, all chemical bonds, have one thing in common. They all owe their existence, directly or indirectly, to the electrostatic attractive forces that exist between the positively charged protons in the nucleus of an atom and the negatively charged electrons surrounding the atoms nearest to it.

The noble gases

During the nineteenth century, chemists organised the elements known at the time according to how they bonded chemically with other elements. They found that one group of elements (the elements we know today as the noble gases) tended to be found in nature in their elemental form.

The atoms of the noble gas elements did not seem inclined to combine with the atoms of other elements - or, indeed, with each other. It was eventually discovered that these same elements all had one thing in common - they all, apart from helium, have eight electrons in their outermost shell (helium is the exception because it only has two electrons in total).

The role of valence electrons

The electrons in the outer electron shell of an atom are known as valence electrons. They have higher energy levels, and are further away from the nucleus, than other electrons. As a result, they are only weakly bound to the nucleus, and it takes relatively little energy to separate them from the parent atom. It is these valence electrons that are involved in any chemical bonding that occurs between atoms.

Atoms with eight valence electrons have an outer electron shell consisting of one s orbital and three p orbitals, each of which contains two electrons. In this configuration, individual electrons have less energy than they would otherwise have, and are relatively tightly bound to the nucleus of the atom. This gives the atom a very stable electronic configuration.

The chemical bonding that takes place between atoms invariably involves the exchange or sharing of electrons. The atom's electronic configuration - which is defined by the number of valence electrons it has - will therefore determine how likely it is that the atom will bond with other atoms. As a general rule, the atoms of all elements apart from the noble gases have the potential to bond with other atoms.

Atoms with only one or two valence electrons have a tendency to give up those electrons when they form a chemical bond with another atom in order to achieve a stable outer shell. Conversely, atoms that require only one or two valence electrons in order to achieve a noble gas configuration tend to form bonds with other atoms that allow them to acquire the missing electrons.

A typical chemical reaction involves a process of bonding between two or more atoms in which each atom achieves a stable electronic configuration. The resulting molecule or formula unit also has a stable electronic configuration, and the energy levels in the participating atoms is lower than it would be for the atoms in their unbound state.

Covalent bonding

Covalent bonding occurs when two or more atoms share electrons in order for each atom to achieve a stable electronic configuration. Covalent bonds can be formed between atoms of the same type, or between atoms of different types. The simplest example of covalent bonding occurs when two hydrogen atoms combine to form a hydrogen molecule.

In a hydrogen (H₂) molecule, the negatively charged electron in each hydrogen atom is attracted to the positively charged proton in the nucleus of the other hydrogen atom. The hydrogen atoms are sufficiently close together that their electron clouds overlap. Each hydrogen atom thus effectively gains the second electron it needs in order to achieve a full outer shell.

Covalent bonding occurs in a number of elemental substances. In some elements, the covalent bonds can have a number of different configurations. As a result, different forms of the same element are created that can have very different properties. These variants are known as allotropes. Allotropism is the the property of some chemical elements to exist in two or more different forms in the same physical state (i.e. as a solid, liquid or gas).

The element carbon has four electrons in its valence shell. In order to achieve a stable configuration, it must somehow either gain four electrons or lose four electrons. Losing four electrons completely, or gaining four additional electrons, would leave the carbon atom with a very strong negative or positive charge, either of which would leave it in an unstable state, despite having a complete outer electron shell. Sharing electrons through covalent bonding eliminates this problem.

Well known allotropes of the element carbon include diamond and graphite. These allotropes have very different properties. In diamond, each carbon atom shares its four outer electrons with four other carbon atoms to create a crystalline structure like the one illustrated below.

The animation above was created using VESTA (Visualization for Electronic and STructural Analysis), a software package developed by Koichi Momma and Fujio Izumi that is distributed free of charge for academic, scientific, educational, and non-commercial use.

Covalent bonds can only form if the atoms involved are close enough together to share the electrons in their outer shell. Diamonds are created when a number of carbon atoms each form four covalent bonds with four neighbouring carbon atoms - something that can only occur at extremely high temperature and pressure. Most naturally occurring diamonds were formed deep within the earth's mantle.

In diamond, all of the electrons in the carbon atom's outer shell are involved in covalent bonds. There are therefor no free electrons available to facilitate the conduction of electric current (which is, after all, the movement of free electrons). Diamond therefore does not conduct electricity, unlike another more common allotrope of carbon - graphite.

In graphite, a carbon atom forms three covalent bonds with three other carbon atoms. As a result, graphite essentially forms as a two-dimensional network of carbon atoms just one atom thick, totally unlike the crystalline structure of diamond. A sample of graphite consists of many layers of these two-dimensional networks, often referred to as graphene sheets.

Within each graphene sheet, the carbon atoms are held together tightly by strong covalent bonds, but the bonds between individual graphene sheets consist of relatively weak electrostatic (van der Waals) forces. The sheets can therefore slide over one another relatively easily. This property makes graphite useful as a lubricant. Graphite is also the stuff from which pencil lead is made!

In graphite, only three of the four electrons in the carbon atom's outer shell are involved in covalent bonding; the fourth electron is free to move around between the graphene layers. This makes graphite an excellent conductor of electricity, unlike diamond.

Carbon also forms covalent bonds with other elements - in fact, all forms of life on our planet are based on carbon. One very common carbon compound is methane (CH₄), which is formed when one carbon atom forms covalent bonds with four hydrogen atoms. Each hydrogen atom needs one electron to complete its outer shell, while carbon requires four electrons.

Note that the covalent bonds we have seen so far have been single covalent bonds formed by the pairing of a single electron from each atom. Sometimes, two electrons from each atom are shared in a double covalent bond (this is the kind of bond found in oxygen molecules). There are even occurrences of triple covalent bonds being formed, in which three pairs of electrons are shared (for example, in nitrogen molecules). Generally speaking, double covalent bonds are shorter and stronger than single covalent bonds, and triple covalent bonds are shorter and stronger than either single or double covalent bonds.

In the diagrams used to illustrate covalent bonding between two atoms, the shared electrons are often shown occupying the area of overlap between the two atoms (the hydrogen molecule we saw earlier is a good example). In reality, however, the shared electrons are moving at high speed and could be found anywhere around either nuclei at any given moment.

In some covalent bonds, the electrons will be more strongly attracted to the nucleus of one atom than the other, and consequently will spend more time around that atom's nucleus. Covalent bonds in which this kind of behaviour occurs are called polar covalent bonds, because the atom around which the shared electrons spend more of their time has a greater electron affinity (or electronegativity) than the other atom involved in the bond.

The water molecule (H₂O) is a good example of polar covalent bonding. In water, covalent bonds form between each of the two hydrogen atoms and the single oxygen atom, but the positively charged protons in the nucleus of the oxygen atom have a far greater attraction for the shared electrons than the single proton in either of the two hydrogen atoms. The shared electrons therefore spend more time around the nucleus of the oxygen atom than they do around the nuclei of the two hydrogen atoms.

The net effect of this polarisation is to give the oxygen end of the molecule a slightly negative charge and the hydrogen ends a slightly positive charge, even though the molecule is electrically neutral overall. The molecule effectively becomes a dipole (a dipole is a pair of electric charges of equal magnitude but opposite polarity, separated by a small distance).

Ionic bonding

In ionic bonding, as in covalent bonding, electrons have a role to play in giving the participating atoms a complete outer shell. Unlike in covalent bonding, however, the electrons in an ionic bond are not shared between two atoms. Instead, an electron is transferred from one atom to the other.

The atom that gives up the electron becomes a positively charged ion known as a cation. The atom that receives the electron becomes a negatively charged ion called an anion. The resulting chemical bond is due to the electrostatic attraction that the oppositely charged ions have for one another.

Ionic bonding usually occurs between atoms of elements from groups one and eighteen of the perodic table. Elements in group one, such as sodium and potassium, have just one valence electron. Giving up this electron will leave them with a full outer electron shell.

Group eighteen elements like chlorine and fluorine, on the other hand, have seven valence electrons. They must therefore acquire one additional electron in order to achieve a full outer electron shell.

Chemical reactions between the (highly reactive) elements of these two groups result in stable chemical compounds like sodium chloride (common salt), a crystalline compound created when equal numbers of sodium and chlorine atoms combine. The sodium atoms become positively charged sodium ions, while the chlorine atoms become negatively charged chloride ions.

Note that within the resulting crystalline structure, the positive and negative ions are not arranged in pairs. Instead, each ion is surrounded by ions of the opposite charge in a three-dimensional crystal lattice. In sodium chloride, for example, each sodium ion (Na+) is surrounded by six chloride (Cl-) ions. Similarly, each chloride ion is surrounded by six sodium ions.

Elements that have seven valence electrons are more electronegative than other elements (electronegativity is the ability of an atom to pull electrons away from other atoms). They are therefore more likely to gain an electron and become negatively charged. The opposite is true of elements that only have one valence electron, like sodium. They are less electronegative than other elements, and therefore more likely to lose an electron and become positively charged.

Although ionic bonds are strong, they are not as strong as covalent bonds, and can be easily broken under the right circumstances. Ionic compounds tend to dissociate (split into their constituent ions) in water, for example, because the positive and negative ions are more strongly attracted to the negative and positive poles of the water molecules (water molecules are polar - see below) than they are to each other.

Metallic bonding

Metallic bonding, as the name suggests, is the kind of bonding that usually occurs in metals. Metallic bonding occurs between the atoms of metallic elements to form metallic solids. It is different to both covalent bonding (in which electrons are shared between two atoms), and ionic bonding (in which one atom donates electrons to another atom).

In a metal, one or more electrons in each atom have sufficient energy to break away from their parent atom. This results in the creation of positively charged metal ions that are surrounded by a kind of "sea" of negatively charged electrons. These free, or delocalised, electrons are shared between the metal atoms. They hold the ions together, a bit like glue.

Although the negatively charged delocalised electrons are attracted to the positively charged metal ions, a single delocalised electron is not associated with one particular atom. In fact, at any one time, it can be associated with a significant number of metal ions. This metallic bonding is what gives metals their unique properties, which we will talk about elsewhere in this section.

Weak intermolecular bonding

The strong chemical bonds that form between atoms of the same or different elements to create molecules or formula units of an element or chemical compound are called primary bonds. Covalent bonds, ionic bonds and metallic bonds are all types of primary bond.

The forces that attract two or more molecules to one another are called secondary bonds, or Van der Waals forces after the Dutch theoretical physicist Johannes Diderik van der Waals (1837-1923), who made significant contributions to modern molecular science. Secondary bonds are significantly weaker than primary bonds. Several types of secondary bond have been identified. They include:

• London dispersion forces
• Dipole-dipole interactions
• Hydrogen bonds

London dispersion forces

London dispersion forces are the weakest of the intermolecular forces we will examine. They are attractive forces that occur between the positive and negative ends of instantaneous dipoles. These dipoles arise because the electrons in the valence shell of an atom are not always evenly dispersed throughout the atom's electron cloud.

Over a period of time, the valence electrons are equally dispersed around the nucleus of the atom. At any given instant, however, the distribution of valence electrons throughout the cloud is asymmetric, creating a small charge imbalance. This charge imbalance can induce a dipole in a neighbouring atom if it is close enough, causing a momentary dipole-dipole attraction between the two atoms. If the atoms happen to belong to different molecules, a temporary intermolecular bond is created.

Because the distribution of valence electrons within an atom is constantly changing, the weak intermolecular bonds created between neighbouring molecules due to London dispersion forces are extremely short lived. Even so, they can have significant consequences.

As a gas cools down, its kinetic energy decreases. At some point, there will not be enough kinetic energy to overcome even the weak bonds created by the London dispersion forces, which are then able to bind the molecules together in a liquid state. If the liquid is further cooled, its kinetic energy will continue to decrease, and the dispersion forces will eventually cause it to freeze and become a solid.

Even a noble gas like argon, which has a full outer electron shell, will condense into a liquid if you cool it down sufficiently. The degree to which an atom can be polarised by London dispersion forces (called its polarisability) increases with atomic size because, in larger atoms, the valence electrons are not so tightly bound to the nucleus and are thus better able to form temporary dipoles. That's why argon condenses at a much higher temperature (-186 °C) than helium (-272 °C), even though both elements are noble gases.

London dispersion forces are often the only attractive forces between non-polar molecules like those of the hydrocarbons and, because the forces are weak and of short duration, two molecules must be in very close proximity for them to have any effect. The strength of the intermolecular bonds created by the dispersion forces is thus directly proportional to the surface area over which they can operate, which is dependent on the shape and size of the molecule.

Although the hydrocarbons neopentane and n-pentane both have the same chemical formula (C₅H₁₂), they have different molecular structures, as illustrated below. The n-pentane molecule has a larger surface area, allowing the London dispersion forces to create stronger intermolecular bonds between neighbouring n-pentane molecules. As a result, n-pentane has a significantly higher boiling point then neopentane.

As a general rule, the strength of the intermolecular bonds created by dispersion forces, together with the boiling point of a substance, increases in proportion to molecular weight. Small molecules like methane (the main component of natural gas) experience weak intermolecular forces and thus have low boiling points. Methane has a boiling point of −161.49 °C. Larger molecules like octane (an important component of petrol), have much higher boiling points. Octane has a boiling point of between 125.1 °C and 126.1 °C.

Dipole-dipole interactions

We saw above that covalently bonded atoms share electrons in order to achieve a full outer electron shell. If the atoms taking part in the bond are not equally electronegative, we have a polar covalent bond in which the shared electrons spend most of their time around the nucleus of the most eletronegative atom. As a result, the molecule created by the bonding process becomes polarised. One end of the molecule has a small negative charge, and the other end has a small positive charge.

Molecules that have polar covalent bonds are called polar molecules. They are electrically neutral overall, but the separation of charge within the molecule effectively turns it into a dipole, like the water molecule we saw above. If two polar molecules are close enough together, the positive end of one of the molecules will be attracted to the negative end of the other molecule.

Hydrogen bonds

A hydrogen bond is essentially a special case of the dipole-dipole interaction. A hydrogen bond occurs when a hydrogen atom that is covalently bonded to a highly electronegative atom such as nitrogen, oxygen or fluorine (this atom is known as the proton donor) is attracted to another highly electronegative atom (known as the proton acceptor) nearby. Hydrogen bonds can occur both between molecules and within a single (complex) molecule.

The illustration above shows a hydrogen bond between two water molecules. In a water molecule, each hydrogen atom's electron is strongly attracted to the highly electronegative oxygen atom. As a result, it spends far more time around the nucleus of the oxygen atom than it does around the nucleus of the hydrogen atom. Each hydrogen atom thus carries a net positive charge, while the oxygen atom carries a net negative charge.

This difference in charge means that water molecules are dipoles, and we could therefore expect polar bonding to take place between them. Remember, however, that the hydrogen nucleus is very small - it has approximately one-sixteenth the mass of the oxygen atom. As a result, the net positive charge on the hydrogen atom has a relatively high charge density.

When two water molecules are sufficiently close together, one of the positively charged hydrogen atoms belonging to one of the water molecules (i.e. the proton donor) is strongly attracted to the negatively charged oxygen atom belonging to the other water molecule (the proton acceptor).

Hydrogen bonds are not as strong as ionic or covalent bonds, but they are much stronger than other types of secondary bond due to the hydrogen atom's high charge density, which increases the strength of the electrostatic attraction between the molecules involved in the bond.

When a water molecule forms, two of its six valence electrons participate in covalent bonds with the two hydrogen atoms, each of which can form hydrogen bonds with other molecules. This leaves the oxygen atom with two free pairs of electrons, both of which can also participate in hydrogen bonds with other molecules. A single water molecule can thus form hydrogen bonds with up to four other molecules.

In a container full of water, each water molecule can form hydrogen bonds with four other water molecules as shown below. Because hydrogen bonds are significantly weaker than covalent bonds, however, they will continually break and reform. One estimate puts the average number of hydrogen bonds per water molecule at any given time at 3.59.

Intermediate types of bonding

There are a number of intermediate kinds of chemical bonding that do not fall precisely into one of the categories discussed above. Covalent bonds, for example, involve two atoms sharing a pair of electrons, whereas in ionic bonds, a metal atom will relinquish one or more electrons to a non-metal atom. In reality, however, all of these bonds have some covalent character in the sense that the electrons are still shared between the participating atoms.

The degree to which a particular bond is ionic or covalent depends on the difference in electronegativity between the two atoms. Electronegativity, remember, is the degree to which an atom can pull electrons away from other atoms. The greater the difference in electronegativity, the more ionic the nature of the bond will be.

The electronegativity of an atom cannot be directly measured; it is calculated from other atomic or molecular properties, using one of a number of available methods. Values are usually determined using the Pauling scale, named after the American chemist Linus Carl Pauling (1901-1994), who first proposed it in 1932.

The scale is dimensionless (it does not have units), and has a minimum value of 0.0 and a maximum value of 4.0. It is based on bond energy calculations for various elements in covalent bonds. One commonly used rule of thumb predicts that bonds created between the atoms of two different elements will be ionic in nature if the difference in electronegativity between them is 1.7 or greater; otherwise it will be covalent in nature.

To take an example, the electronegativity value for sodium has been calculated as 0.93, whereas the electronegativity value for chlorine has been calculated as 3.16. This gives a difference in electronegativity between these two elements of 2.23, indicating that the chemical bonds formed between sodium and chlorine atoms will be strongly ionic in nature.