Covalent bonding - overview

Electrons have a negative charge and the nuclei they surround have a positive charge. When the two particles approach one another the electron clouds can overlap. When this happens, under certain circumstances, the electrons that are between two nuclei can be attracted to both nuclei, holding them together.

This force of attraction is known as a chemical bond. When the atoms form a bond they become lower in energy and the system is more stable. The energy saved by moving to a more stable situation is released as heat. For this reason bond formation is always exothermic, i.e. heat energy is released. Conversely, in order to break a chemical bond energy must be used - it is an endothermic process. Covalent bonding occurs between atoms of non-metals

4.2.2: Describe how the covalent bond is formed as a result of electron sharing. Dative covalent bonds are required. Examples include CO, NH4+ and H3O+.

Molecular structures

Oxygen molecule formed by two atoms sharing one pair of electrons. This is a double bond.	Nitrogen molecule showing three shared pairs of electrons making full outer shells with eight electrons in each. This is a triple bond

4.2.3: Deduce the Lewis electron dot) structure of molecules and ions for up to four electron pairs on each atom. A pair of electrons can be represented by dots and crosses or by a line. For example chlorine can be shown as:

The octet rule

On studying many molecular compounds it appears that a full set of electrons in the outer shell is a particularly stable situation and this gives rise to the octet rule, which states that atoms tend to share electrons so as to obtain eight electrons in their outer shells

Electron sharing

To form a single covalent bond one electron is provided by each of the bonding atoms making a shared pair. The simplest situation occurs when two atoms of a non-metal share one pair of electrons.

Chlorine molecule formed by two atoms sharing one pair of electrons

Here, we can see that although originally each chlorine atom has only seven electrons in its outer shell, after bonding it has eight (an octet), making the molecule more stable than the sum of the original atoms.

It is possible for atoms to share two pairs of electrons to achieve the octet (8 electrons).

Carbon dioxide molecule formed from two oxygen atoms each sharing two pairs of electrons with a carbon atom

When two pairs of electrons are shared between two atoms this is called a double bond. In the diagram above the carbon dioxide molecule has double bonds between the carbon and each of the oxygen atoms. Triple bonds are also possible (three shared pairs of electrons)

covalent bonding animation

Lewis Structures

4.2.4: State and explain the relationship between the number of bonds, bond length and bond strength. The comparison should include bond lengths and bond strengths of: two carbon atoms joined by single, double and triple bonds, the carbon atom and the two oxygen atoms in the carboxyl group of a carboxylic acid.

Bond strength

A single bond is a shared pair of electrons that is attracted to both nuclei of the bonded atoms.

The bonding electrons draw the atoms closer together but as the atoms get very close they experience repulsive forces from the other non-bonding electrons and from the two nuclei themselves. This means that there is an optimum distance for the two atoms. This is called the bond length.

The single bonds hold atoms together by the forces of attraction between the electron pair (bonding pair) and the two nuclei. As the nuclei of different atoms are obviously different from one another then this force of attraction and hence the bond strength varies between different pair of atoms.

The strength of bonds can be measured by several techniques, usually by seeing how much energy is needed to break the bond.

Some bond energy values

Although it is difficult to make a direct relationship between bond length and strength there are some inferences that can be obtained. It can be seen that as the bond length of the carbon - halogen bonds increases so the bond energy decreases. As the atoms get larger they are held further apart by inter-electron repulsions. The attractive force between the bonding electron pair and the nuclei is consequently weaker.

If single, double, and triple bonds are compared a distinct pattern emerges:

As the bond strength INcreases, so the bond length DEcreases. This follows from a consideration of the force of attraction between the greater number of pairs of electrons and the two nuclei. Four electrons (two pairs) can pull the two nuclei closer together than two electrons (one pair).

The bond strength INcreases as the bond length DEcreases. This follows from a consideration of the force of attraction between the greater number of pairs of electrons and the two nuclei. Four electrons (two pairs) can pull the two nuclei closer together than two electrons (one pair). The same argument explains why a triple bond is even stronger than a double bond. The carbon = carbon triple bond is much stronger than the C=C double bond which is stronger in turn than the C-C single bond

The carboxlate group

The carboxylate group of atoms occurs in the carboxylic acids, such as ethanoic acid CH3COOH or methanoic acid HCOOH. In these acids there is a carbon atom bonded to two different oxygen atoms, one using a single bond and the other with a double bond. These bonds have different strengths and lengths.

4.2.5: Predict whether a compound of two elements would be covalent from the relative electronegativity values or from their positions in the periodic table

Prediction of bonding type

Compounds in which the bonded atoms have a large electronegativity difference for ionic compounds. Where the difference is slight, they are covalent. There is no hard and fast value at which the change occurs. Rather there is a greater and greater degree of covalency as the values become closer together.

Perhaps the closest to a 'cut off' is the compound formed between aluminium and chlorine. In the solid state at 0ºC there is considerable evidence that it is ionic, but at room temperature it seems to be covalent. At higher temperatures it sublimes as a dimer with the formula Al₂Cl₆.

Aluminium has an electronegativity of 1.5 and chlorine 3.0. That makes the difference in electronegativity = 1.5 units on the Pauling scale. This is a good value to use as a 'rule of thumb'.

Greater than 1.5 units = ionic

Less than 1.5 units = covalent.

Ionic or covalent bond formation

It was stated above that metals reacting with non-metals form ionic bonds. This is the GENERAL rule. It is really the ease with which a metal can lose electrons coupled with the attraction that a non-metal has for more electrons.

When an element loses electrons easily and another element has a high attraction for electrons then this encourages ionic bond formation. As metals lose electrons easily they are said to be electropositive, or to put it another way, to have low electronegativity values.

Non-metals tend to attract electrons and the non-metals closest to group 7 have higher electronegativity values. Electronegativity increases from bottom left to top right in the periodic table.

When the DIFFERENCE between the electronegativity values is large, ionic bonds are formed.

It is usually the case that the difference in electronegativity values between metals and non-metals is large enough to cause ionic bond formation (transfer of electrons etc etc), however this is not the case for all metal:non-metal combinations.

4.2.6: Predict the relative polarity of bonds based on electronegativity values. AIM 7: Simulations may be used here

Comparison of electronegativity difference

From the previous section you can see that a direct comparison of bonds may be done using the electronegativity difference between the atoms in the bonds.

It is safe to say that the C-F bond (electronegativity difference = 1.5) has a greater dipole than the C-Br bond (electronegativity difference = 0.3)

4.2.7: Predict the shape and bond angles for species wth four, three, and two negative charge centres on the central atom using the valence shell electron pair repulsion theory VSEPR. Example should include CH4, NH3, H2O, NH4+, H3O+, BF3, C2H4, SO2, C2H2, and CO2. Simulations are available to study the three dimensional structure of these and the structures in 4.2.9 and 4.2.10

water

The central oxygen atom has its electron arranged into four pairs in four distinct regions (the orbitals are sp3 hybridised)

These repel one another and adopt a tetrahedral arrangement.

However only two of the electron pairs are used in bonding and the other two pair are 'lone' (ie cannot be seen). The shape of the molecule is therefore 'angular' or 'bent'

Tetrahedral electronically but the molecular shape is angular

The valence shell electron pair repulsion theory now looks at the relative strength of the repulsions between the lone pairs and the bonding pairs. As the lone pairs are not drawn further away from the central atom by another atomic nucleus then they exert a greater repulsion on each other than the bonding pairs do on each other. Intermediate is the repulsion felt between a bonding pair and a one pair.

Order of repulsion strength:

lone pair- lone pair >> lone pair - bonding pair >> bonding pair - bonding pair

This causes the tetrahedral electronic shape to distort and squeezes the bonding pairs together. The bond angle then closes sightly from 109,5º to 104,5º

H-O-H bond angle 104,5º

Water

(click to see in 3D - if it doesn't work you have to download chime)

4.2.8: Predict whether a molecule is polar from its molecular shape and bond polarities.

Bond polarity

This is caused by a difference in electronegativity between two bonded atoms. Most bonds are polar, but in reality only those with an electronegativity difference of at least 1 unit on the Pauling scale shows the effect.

For example carbon has an electronegativity of 2.5 and hydrogen 2.1. In principle they are polarised and the bond has a dipole, but the two values are close enough together as to be insignificant. Carbon - hydrogen bonds are not said to be polar.

Resolution of dipoles

Dipoles are vector quantities in that they have both magnitude and direction. They can be resolved in the same way as forces.

If the resolution of all vector dipoles in a molecue leaves an overall dipole in any specific direction then the molecule is polar.

Example: Water

Water has two dipoles on the oxygen - hydrogen bonds. If the molecule is orientated with the oxygen at the top these cancel out in a horizontal direction but add up in a vertical direction.

Water is consequently a polar molecule.

Example: Carbon dioxide

In this case the carbon - oxygen bonds are also polarised, but they cancel out in the horizontal axis, while there is no component in the vertical axis. Carbon dioxide is therefore non-polar.

Polar molecules

Polar molecules attract one another increasing the degree of intermolecular force experienced. This means that they have higher boiling ponts than expected from purely van der Waaal's forces.

4.2.9: Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite and C60 fullerene)

Allotropy

Allotropes are different physical forms of the same element. All elements are made up uniquely of their own atoms and therefore any physical differences must be a consequence of how the atoms are joined together - their arrangement within the bulk structure.

Many elements exhibit allotropy as there are often varous ways in which the atoms can be linked together into molecules and also different ways in which the molecules can be arranged to make larger structures.

In the case of carbon, the atoms form either giant macromolecular structures (diamond and graphite) in which all of the atoms in the bulk structure are joined together by covalent bonds making giant molecules, or smaller molecules (buckminster fullerene) in which there are only discrete molecules made up of 60 carbons in a structure resembling a football (hence the nickname 'bucky balls')

Diamond

Each carbon in a diamond crystal is bonded to four other carbon atoms making a giant macromolecular array (lattice). As each carbon has four single bonds it is sp³ hybridised and has tetrahedral bond angles of 109º 28'

Properties of diamond

• hardest substance known to man
• brittle (not malleable)
• insulator (non-conductor)
• insoluble in water
• very high melting point

Physical properties of diamond explained by considering the structure and bonding

Graphite

Again the carbon atoms are bonded together to make a giant structure but in this case all of the carbons are bonded to only three neighbour and are sp² hybridised. As the sp² hybridisation results in planar structures, there are giant 2 dimensional layers of carbon atoms and each layer is only weakly linked to the next layer by Van der Waal's forces.

Physical properties of graphite explained by considering the structure and bonding

Fullerenes

These are small molecules of carbon in which the giant structure is closed over into spheres of atoms (bucky balls) or tubes (sometimes caled nano-tubes). The smallest fullerene has 60 carbon atoms arranged in pentagons and hexagons like a football. This is called Buckminsterfullerene.

The name 'buckminster fullerene' comes from the inventor of the geodhesic dome (Richard Buckminster Fuller) which has a similar structure to a fullerene. Fullerenes were first isolated from the soot of chimineys and extracted from solvents as red crystals.

The bonding has delocalised pi molecular orbitals extending throughout the structure and the carbon atoms are a mixture of sp2 and sp3 hybridised systems.

Fullerenes are insoluble in water but soluble in methyl benzene. They are non- conductors as the individual molecules are only held to each other by weak van der Waal's forces.

Buckminster fullerene	Structure
	As the molecule is totally symmetrical with all bond lengths and angles being equal, it is likely/inevitable that the hybridisation of the carbon atoms is somewhere between that of sp2 and sp3. Another example of a theory (hybridisation in this case) having to be modified to accomodate the observed experimental data.

Physical properties of fullerenes explained by considering the structure and bonding

4.2.10: Describe and structure and bonding in silicon and silicon dioxide

These are giant covalent structures, with the bonding covalent from atom to atom in a never ending array. The bond angles at each silicon atom is 109º 28'.

The oxygen atoms act as bridges between silicon atoms in silicon dioxide.

Resources

covalent bonding animation

covalent bonding in oxygen

covalent bonding in hydrogen

bonding in water