The Covalent Bond, Chemistry tutorial


Matter is made up of atoms and ions that knowledge both attractive and repulsive forces. It is the balance of such forces that consequence in chemical bonding within molecules, in metals, and in ionic compounds. In this chapter, we shall explain the forces between particles and how such forces effect in the configuration of covalent bonds within molecules.

Bonding between non-metals consists of 2 electrons which are shared between 2 atoms. Using the Wave Theory, the covalent bond includes an overlap of the electron clouds from each atom. The electrons are concentrated in the region between the 2 atoms. In covalent bonding, the 2 electrons shared through the atoms are magnetized to the nucleus of both atoms. Neither atom entirely loses or gains electrons as in ionic bonding.

The covalent bond

In chapter 1 of this module, we conversed how atoms could attain an entire shell of electrons through losing or expanding one or more electrons, to form ions. There is another way atoms can satisfy the octet rule: they can share electrons. For instance, 2 hydrogen atoms can share their electrons, as following below. Since each of the shared electrons then 'belongs' to both atoms, both atoms then have a full shell, through 2 electrons. The pair of shared electrons is symbolized via the heavy line between the atoms.

2353_Formation of the Hydrogen Molecule.jpg

Fig: Formation of the Hydrogen Molecule by Simple Covalent Bonding

In terms of charge-charge interactions, what happens is that the shared electrons are located between the two bonded atoms. The force attracting them to both nuclei is stronger than the repulsive force between the nuclei.

The methane (CH4) molecule demonstrates a more complex instance. Each of the 4 electrons in the outermost ('valence') shell of carbon is shared through each of the hydrogen. In turn, each of the hydrogen as well shares one electron by carbon. Largely, carbon 'owns' 10 electrons - to satisfy the octet rule - and each hydrogen has 2 electrons. This is shown in figure:

1617_Formation of the Methane Molecule by.jpg

Fig: Formation of the Methane Molecule by Simple Covalent Bonding

The experimental studies of methane molecule reveals that the four hydrogen atoms spread out evenly around the carbon atom, leading to the 3-dimensional structure here:

As we would expect given that the electrons are shared, if we plot the region where the electrons sit, this isn't localized on one atom, as it was for the ionic compounds, but is all over the molecule:

489_Structure of Methane.jpg

Fig: Structure of Methane

Covalent Bonding and Isomers

As we have seen above, atoms can share electrons through others to form chemical bonds. This can also occur between 2 carbon atoms, to form a molecule such as ethane (C2H6):

1234_Three-Dimensional Structure of Ethane.jpg

Fig: Three-Dimensional Structure of Ethane

When we add 2 more carbon atoms and 4 more hydrogen atoms, to build butane (C4H10), an interesting situation arises: There are 2 different ways of bonding the carbons mutually, to form 2 different molecules, or isomers. For one of the isomers, the 1st carbon is bonded to three hydrogen atoms, and to the second carbon, which is itself bonded to another 2 hydrogen atoms and to the 3rd carbon, which is linked to the 4th carbon. In the other isomer, one of the carbons shapes a bond to all three carbon atoms:

1208_Three-Dimensional Structures of Butane isomers.jpg

Superior compounds can also be shaped, and they will have even more isomers. For instance, this compound by 8 carbons is termed isooctane, and is one of the major components of petrol for cars

Can we check that the formula for this compound is C8H18? Sketch another compound with the same formula? Because covalent bonds can be formed in many different ways, it is possible to write down, and make many different molecules. Many of these are natural compounds, made by living animals or plants within their cells. This instance illustrates one such molecule, cholesterol (C27H46O), which can add to heart disease in people whose diet is too rich in fats:

1476_Three-Dimensional Structure of Cholesterol.jpg

Fig: Three-Dimensional Structure of Cholesterol

As we know that in this structure, 2 neighboring carbon atoms appear to form only 3 bonds, which would go against the octet rule. In fact, such atoms bond via sharing 2 electrons each (a total of 4 electrons). In this way, they complete their electron shell like the others.

This situation is referred to as a double bond, and is following in the pop-up window as a thicker stick between those 2 atoms.

Other compounds are synthetic, they are made by chemists. Chemists can also make the natural compounds, starting from only simple things as methane and water. The 'natural' molecules made in this way are identical to the "real" natural compounds. Other synthetic molecules do not exist in nature. They can have wanted properties, for instance, many medicines are made in this way. An instance of a 'small' medicinal molecule is aspirin, C9H8O4, shown below. In this molecule, 2 bonds between carbon and oxygen are double bonds, and are following as thicker sticks in the model.

1265_Three-Dimensional Structure of Aspirin.jpg

Fig: Three-Dimensional Structure of Aspirin (Acetyl salicylic acid).

Covalent solids

Covalent solids are shaped through piling sets of covalent molecules mutually, and relying on their slight "stickiness" to hold the solid together, one can also form solids entirely bound together by covalent bonds. A brilliant instance is diamond that is pure carbon, through each carbon atom bonding to four others, to form a huge 'molecule' enclosing many millions of atoms. Figure illustrates a part of a diamond molecule:

1273_Part of a Diamond Molecule.jpg

Fig: Part of a Diamond Molecule

In diamond, all the carbon atoms share one electron by each of their 4 neighboring carbon atoms. There is another form in which pure carbon can be formed: graphite. This is the major component of the 'lead' in pencils. Here, instead of each carbon having four neighbours, it only has three. Each carbon shares one electron through 2 of its neighbours, and 2 electrons by the third neighbor. In this way, one C-C bond out of three is a double bond. The atoms all bond mutually in planes and the planes stack on top of one an additional as following in Figure:

1464_Planes of Atoms in the Structure of Graphite.jpg

Fig: Planes of Atoms in the Structure of Graphite

In graphite, the C-C bonds in the planes are extremely strong, but the force between the different planes is quite weak, and they can slip over one another. This clarifies the 'soft' feel of graphite, and the fact that it is utilized as a lubricant, for instance in motor oil. Other 'Big' Covalent Molecules In solids as diamond and graphite, the different atoms all bond to one another to form one extremely huge molecule. The atoms are bonded to each other in all directions in diamond, and in two directions (within the planes) in graphite, by no bonding in the other direction. Some important covalent molecules engage atoms bonding to each other repeatedly along just one direction, through no bonds in the others. Such are termed polymers, and one simple illustration if polyethene (as well termed polythene, or polyethylene). The structure of polythene is exposed in Figure (the dangling bonds at each end specify how the bonding should actually carry on for thousands of atoms on each side):

826_Three-Dimensional Structure of Polythene.jpg

Fig: Three-Dimensional Structure of Polythene

Polythene is what most plastic bags are made of. Other polymers contain molecules such as nylon, Teflon (such as polythene, are man-made), or cellulose (the stuff that builds wood hard), a biological polymer.

Co-ordinate (Dative Covalent) Bonding

A covalent bond is shaped via 2 atoms sharing a pair of electrons. The atoms are held mutually since the electron pair is attracted through both of the nuclei.

In the formation of an easy covalent bond, each atom supplies one electron to the bond - but that doesn't needed have to be the case. A co-ordinate bond (as well termed a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons appear from the similar atom.

For the rest of this chapter, we shall utilize the term co-ordinate bond - but if we prefer to call it a dative covalent bond, it is acceptable.

The reaction between ammonia and hydrogen chloride if such colourless gases are permitted to mix, a thick white smoke of solid ammonium chloride is formed.

Ammonium ions, NH4+, are formed through the transfer of a hydrogen ion from the hydrogen chloride to the lone pair of electrons on the ammonia molecule.

NH4   +   HCl    →      NH4Cl

1561_Formation of Ammonium Chloride by Coordinate.jpg

Fig: Formation of Ammonium Chloride by Coordinate Bonding

Whenever the ammonium ion, NH4+, is shaped, the 4th hydrogen is attached through a dative covalent bond, since only the hydrogen's nucleus is shifted from the chlorine to the nitrogen. The hydrogen's electron is left behind on the chlorine to form a negative chloride ion.

Once the ammonium ion has been formed it isn't possible to tell any diverse between the dative covalent and the ordinary covalent bonds. Even though the electrons are exposed differently in the diagram, there is no difference between them in reality.

Symbolizing co-ordinate bonds in easy diagrams, a co-ordinate bond is exposed via an arrow. The arrow points from the atom donating the lone pair to the atom admitting it.

Dissolving hydrogen chloride in water to create hydrochloric acid something alike occurs. A hydrogen ion (H+) is shifted from the chlorine to one of the lone pairs on the oxygen atom.

1015_Dissolution of Hydrogen Chloride in Water.jpg

Fig: Dissolution of Hydrogen Chloride in Water

H2O + HCl   →   H3O+ +   Cl-

The H3O+ ion is variously termed the hydroxonium ion, the hydronium ion or the oxonium ion.

If we write the hydrogen ion as H+ (aq), the '(aq)' symbolizes the water molecule that the hydrogen ion is attached to. When it reacts through something (an alkali, for example), the hydrogen ion simply becomes detached from the water molecule again.

Note that once the co-ordinate bond has been set up, all the hydrogen atoms attached to the oxygen are precisely the same. When a hydrogen ion splits away again, it could be any of the three.

The reaction between ammonia and boron trifluoride, BF3

If we have recently read the page on covalent bonding, we might remember boron trifluoride as a compound that doesn't have a noble gas structure around the boron atom. The boron only has 3 pairs of electrons in its bonding level, whereas there would be room for 4 pairs. BF3 is explained as being electron deficient.

The lone pair on the nitrogen of an ammonia molecule can be utilized to overcome that deficiency, and a compound is shaped involving a co-ordinate bond.

1728_Reaction between Ammonia.jpg

Fig: Reaction between Ammonia and Boron Trifluoride

Using lines to represent the bonds, this could be drawn more simply as:

2444_second diagram.jpg

The second diagram shows another way that we might discover co-ordinate bonds drawn. The nitrogen end of the bond has happen to positive since the electron pair has shifted away from the nitrogen towards the boron - which has therefore become negative. 

The structure of aluminium chloride

Aluminium chloride sublimes (turns straight from a solid to a gas) at about 180°C. If it simply enclosed ions, it would have an extremely high melting and boiling point since of the strong attractions between the positive and negative ions. The implication is that when it sublimes at this relatively low temperature, it must be covalent. The dots-and-crosses diagram illustrates only the outer electrons.

AlCl3, like BF3, is electron deficient. There is likely to be a similarity, since aluminium and boron are in the similar group in the Periodic Table, as are fluorine and chlorine. Measurements of the relative mass of aluminium chloride show that its formula in the vapour at the sublimation temperature isn't AlCl3, but Al2Cl6. This implies that it exists as a dimer (two molecules joined together). The bonding between the 2 molecules is co-ordinate, using lone pairs on the chlorine atoms. Each chlorine atom has 3 lone pairs, but only the 2 significant ones are given away in the line diagram.

1263_The Structure of Aluminium Chloride.jpg

Fig: The Structure of Aluminium Chloride

Energy is liberated when the 2 co-ordinate bonds are formed, and so the dimer is more stable than 2 divide AlCl3 molecules.

The bonding in hydrated metal ions

Water molecules are powerfully attracted to ions in solution - the water molecules clustering around the positive or negative ions. In many cases, the attractions are so huge that formal bonds are made, and this is true of roughly all positive metal ions. Ions through water molecules attached are explained as hydrated ions.

Although aluminium chloride is covalent, when it dissolves in water, ions are produced. Six water molecules bond to the aluminium to provide an ion through the formula Al (H2O)63+. It is termed the hexa aquaaluminium ion - that translates as six ("hexa") water molecules ("aqua") wrapped around an aluminium ion.

The bonding in this (and the similar ions formed by the great majority of other metals) is co-ordinate (dative covalent) using lone pairs on the water molecules.

Aluminium is 1s22s22p63s23px1. When it forms an Al3+ ion, it loses the 3-level electrons to leave 1s22s22p6.

This implieds that all the 3-level orbitals are now empty. The aluminium re-organizes (hybridises) six of these (the 3s, three 3p, and two 3d) to produce 6 new orbitals all with the same energy. Such 6 hybrid orbitals admit lone pairs from 6 water molecules.

We might wonder why it chooses to utilize six orbitals rather than 4 or 8 or whatever. Six is the maximum number of water molecules that is possible to fit around an aluminium ion (and most other metal ions). By making the maximum number of bonds, it liberates most energy and so happens to most energetically stable. 

1953_lone pair.jpg

Only one lone pair is following on each water molecule. The other lone pair is pointing away from the aluminium and hence not engaged in the bonding. The consequential ion looks like this:

204_centre of the ion.jpg

Since of the movement of electrons towards the centre of the ion, the 3+ charge is no longer located entirely on the aluminium, but is now spread over the whole of the ion.

Carbon monoxide, CO

Carbon monoxide can be thought of as having two ordinary covalent bonds between the carbon and the oxygen plus a co-ordinate bond using a lone pair on the oxygen atom.

2379_Structure of Carbon.jpg

Fig: Structure of Carbon (II) Oxide

Nitric acid, HNO3

In this case, one of the oxygen atoms can be thought of as attaching to the nitrogen via a co-ordinate bond using the lone pair on the nitrogen atom.

1747_Structure of Trioxonitrate.jpg

Fig: Structure of Trioxonitrates (V) Acid

In fact, this structure is misleading since it suggests that the 2 oxygen atoms on the right-hand side of the diagram are linked to the nitrogen in different ways. Both bonds are really identical in length and strength, and so the arrangement of the electrons must be identical.

There is no way of demonstrating this using a dots-and-crosses picture. The bonding engages delocalization.

Properties of Covalent Molecules

The covalent bonds between atoms in a given molecule are extremely tough; as strong as ionic bonds. Though, unlike ionic bonds, there is a bound to the number of covalent bonds to other atoms that a following atom can form. For instance carbon can build 4 bonds - not more while oxygen can shape only 2 bonds. As a consequence, once each atom has made all the bonds it can make, as revealed in all the molecules above, the atoms can no longer interact through other ones. For this reason, 2 covalent molecules barely stick mutually. Light molecules are hence gases, these as methane or ethane, above, hydrogen, H2, nitrogen, N2 (the main component of the air we breathe, etc. Heavier molecules, these as the isooctane molecule, are liquids at room temperature, while others, for instance cholesterol, are solids.

Typical properties of simple covalent substances

1. The electrical forces of attraction, which is the chemical bond, between atoms in a molecule are frequently extremely tough, consequently; most covalent molecules don't transform chemically on moderate heating. For example, although a covalent molecule as iodine, I2, is readily vaporized on heating, it doesn't break up into iodine atoms I. The I-I covalent bond is strong sufficient to hold out the heating and the purple vapour still consists of the similar I2 molecules as the dark colored solid is made up of. 

2. Ease of vaporization on heating- The electrical attractive forces between individual molecules are weak, so the bulk material isn't extremely strong physically and there are as well consequences for the melting and boiling points. 

3. Such weak electrical attractions are identified as intermolecular forces and are eagerly weakened further on heating. The consequence of absorbing heat energy consequences in raised the thermal vibration of the molecules that weakens the intermolecular forces. In liquids, the amplify in the average particle kinetic energy builds it easier for molecules to defeat the intermolecular forces and change into a gas or vapour. Consequently, little covalent molecules tend to be volatile liquids through low boiling points, simply vaporized or low melting point solids. 

4. On heating the inter-molecular forces are simply defeat through the raised kinetic energy of the particles giving the material a low melting or boiling point and a comparatively small amount of energy is required to consequence such state transforms. This contrasts via the high melting points of giant covalent structures by their strong 3D network. 

As we know the weak electrical attractive forces between molecules, the so termed intermolecular forces should be visibly differentiated between the strong covalent bonding between atoms in molecules (little or giant), and such are sometimes termed to as intramolecular forces (that is internal to the molecule). 

5. Covalent structures are generally poor conductors of electricity since there are no free electrons or ions in any state to bear electric charge. 

6. Most small molecules will dissolve in some solvent to form a solution. This once more contrasts through giant covalent structures where the well-built bond network stops solvent molecules interacting by the particles making up the material. 

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