Topic 4.2: Covalent Bonding

4.2.1 Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei.

When atoms of two non-metals react together, each is seeking to gain electrons in order to achieve the stable electron structure of a noble gas. By sharing an electron pair, they are effectively able to achieve this. The shared pair of electrons is concentrated in the region between the two nuclei and is attracted to them both. It therefore holds the atoms together by electro-static attraction and is known as a covalent bond.

4.2.2 Describe how the covalent bond is formed as a result of electron sharing

The covalent bond is often described as electron sharing as those electrons will be shared between the two positive nuclei. Sometimes there are not enough electrons to achieve the octet rule, stable arrangement of eight electrons in their outer shell, on all the atoms in the molecule. Atoms will have to share more than one electron pair; in other words a multiple bond. Double bonds forms when two electron pairs and a triple bond forms when three electron pairs are shared.

4.2.3 Deduce the Lewis (electron dot) structures of molecules and ions for up to four electron pairs on each atom.

Lewis structures look something like this.

They can be drawn in either form but they must include a few simple notations.

1. Calculate the total number of valence electrons in the molecule by multiplying the group number of each element by the number of atoms of the element in the formula and find the sum.
2. Draw the skeletal strucutre of the molecule to show how the atoms are linked to each other
3. Use a pair of crosses, two dots or a single line to show one electron pair and put it between the atoms
4. Add more electron pair to complete the octets (8 electrons) around the atoms (other than hydrogen)
5. If there are not enough electrons to complete the octets, form double bonds or triple bonds.
6. Check the total number of electrons

Note: When drawing the lewis structure for ions, calculate the valence electron and add or subtract by the number of electrons based on the charge. Also place a square bracket with the charge shown.

Dative Bonds

Sometimes the bond forms by both the electrons in the pair originating from the same atom. This means that the other atom accepts and gains a share in a donated electron pair. Such bonds are called dative bonds or coordinate bond. In the Lewis structure, it is represented using an arrow.

4.2.4 State and explain the relationship between the number of bonds, bond length and bond strength.

Short Bonds are strong bonds. As it progresses from single bonds to double bonds to triple bonds, it decreases in length but also increases in strength.

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

Polar bonds result from unequal sharing of electrons because not all sharing is equal. This occurs when there is a difference in the electronegativities of the bonded atoms, as the more electronegative atom exerts a greater pulling power on the shared electrons. The bond that is unsymmetrical with respect to electron distribution and is said to be polar. The term dipole is often used to indicate the fact that this type of bond has two separated opposite electric charges.

Note that the symbol δ delta is used to represent a partial charge, but is always less than the unit charge associated with ions. 

Majority of the time, elements in a covalent molecule will come from the right hand side of the periodic table, all non-metals. This is because they have electrons to share and not to lose. If the two atoms have a low electronegativity value difference, then the bond is a covalent.


4.2.6 Predict the relative polarity of bonds from electronegativity values

By finding the difference between the two atoms, we can determine the relative polarity compared with 0, which are two atoms of the same element.

4.2.7 Predict the shape and bond angles for species with four, three and two negative charge centres on the central atom using the valence shell electron pair repulsion theory (VSEPR)

Valence shell electron pair repulsion (VSEPR) theory states that electron pairs found in the outer energy level or valence shell of atoms repel each other and thus position themselves as far apart as possible.

To predict the shape and bond angles of a molecule, the following points might help.

• The repulsion applies to both bonding and non-bonding pairs of electrons
• Double and triple bonded electron pairs are orientated together and so behave in terms of repulsion as a single unit known as a negative charge center.
• The total number of charge centers around the central atom determines the geometrical arrangement of the electrons
• The shape of the molecule is determined by the angles between the bonded atoms.
• Non-bonding pairs of electrons (lone pairs) have a higher concentration of charge than a bonding pair because they are not shared between two atoms and so they cause more repulsion than bonding pairs.

The repulsion is highest between two lone pairs, then a lone pair and a bonding pair, then finally two bonding pairs.

Molecules with two charge centres will position them at 180 degrees to each other, therefore the molecule will have a linear shape.

Molecules with three charge centres will position them at 120 degree to each other, giving a planar triangular shape to the distribution of electrons.

However, if there are lone pairs at the center atom rather than a bond, then the angle will be <120 degree. The shape of this molecule is called bent.

Molecules with four charge centres will position themselves at 109.5 degrees to each other, giving a tetrahedral shape to the electron pairs.

However, if one or more of the charge centers is a non-bonding pair, this will again influence the final shape of the bonding pairs that determines the positions of the atoms. This is called the trigonal pyramidal. The angle is <109.5

When there is a bent molecule, remember that the bent molecule with two pairs of lone pairs creates an even lower angle. 2 lone pairs < 1 lone pair < 120


4.2.8 Predict whether or not a molecule is polar from its molecular shape and bond polarities

The polarity of a molecule depends on the polar bonds that it contains and the way in which such polar bonds are orientated with respect to each other, in other words, on the shape of the molecule.

If the bonds are of equal polarity and are arranged symmetrically with respect to each other, their charge separations will oppose each other and so will effectively cancel each other out.

The molecules carbon dioxide, bromine fluoride and methane are all non-polar because the dipoles cancel out.

However if the bonds are not symmetrically arranged then the polarities will not cancel out. Any molecules with lone pairs will be polar unless the lone pairs cancel out with another lone pair. For example, hydrogen chloride, chloro-methane and ammonia


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

There are substances which have a crystalline structure in which all atoms are linked together by covalent bonds. They are often referred to as a giant molecular structure or a macro-molecule.

Graphite

Each C atom is sp2 hybridized covalently bonded to 3 others, forming hexagons in parallel layers with bond angles of 120 degree. The layers are held only by weak van der Waal's forces so they can slide over each other.

Each carbon atom, contains one non-bonded, delocalized electron. Conducts electricity due to the mobility of these electrons. Uses as a lubricant and in pencils.


Diamond

Each C atom is sp3 hybridized covalently bonded to 4 others tetrahedrally arranged in a regular repetitive pattern with bond angles of 109.5 degree. It is also the hardest known natural substance.

Fullerene C60

Each C atom is sp2 hybrized, bonded in a sphere of 60 carbon atoms, consisting of 12 pentagons and 20 hexagons. Structure is a closed spherical cage in which each carbon is bonded to 3 others. Shape of a football.

It easily accepts electrons to form negative ions; a semiconductor at normal temperature and pressure due to some electron mobility. Reacts with potassium to make superconducting crystalline material. Used to make nanotubes for electronic industries

4.2.10 Describe the structure of and bonding in silicon and silicon dioxide

Silicon is a Group 4 element and so its atoms will have four valence shell electrons. In the elemental state, each silicon atom is covalently bonded to four others in a tetrahedral arrangement. This results in a giant lattice structure much like diamond.

SiO2 is commonly known as silica which forms a giant covalent structure. This is a similar tetrahedrally bonded structure. Each silicon atom is covalently bonded to four oxygen atoms and each oxygen atom is bonded to two silicon atoms.

SiO2 refers to the ratio of atoms within the giant molecule. The structure is strong, insoluble in water, high melting point and does not conduct heat or electricity. Glass and sand are different forms of silica.