STUDY NOTES ON CHEMICAL BONDING AND MOLECULAR STRUCTURE FOR JEE-NEET

By SUPRATIM DAS DURGAPUR-WEST BENGAL #9434008713

Example 1:  Covalent bonding diagram for OXYGEN covalent molecule, molecular formula O2

 

Two atoms of oxygen (2.6) combine to form the molecules of the element oxygen O2 (only the outer shell of oxygen’s electrons is shown).

Each oxygen atom is two electrons short of a full outer shell, so each oxygen atom shares two of its electrons with the other atom, so both oxygen atoms have a full outer shell.

The simplified 'dot and cross' electronic diagram for the covalently bonded oxygen molecule. The molecule

has one O=O double covalent bond (displayed formula), oxygen valency 2. The oxygen molecule is held together by the strong O=O oxygen–oxygen double covalent bond by sharing electrons.

Electronically, by sharing two electrons, both oxygen atoms attain a pseudo neon structure (2.8), so the oxygen atoms effectively have full outer shells in forming the covalent bonds when the atoms share their outer electrons.

The electronic dot & cross Lewis diagrams for covalent bonding in oxygen

Example 2: Covalent bonding diagram for CARBON DIOXIDE covalent molecule, molecular formula CO2

   Linear molecule with zero dipole moment

One atom of carbon (2.4) combines with two atoms of oxygen (2.6) to form the compound carbon dioxide CO2 (only the outer shell of carbon’s electrons is shown).

Carbon is four electrons short of a full outer shell (8 electrons) and oxygen is two electrons short of a full outer shell (8 electrons), so one carbon atom shares its four outer electrons with two outer electrons from

each of the oxygen atoms, so all three atoms now have a full outer shell of 8 electrons in the formation of two double bonds (O=C=O).

Electronically, carbon (2.4) becomes like neon (2.8) and oxygen (2.6), also becomes like neon (2.8), so the hydrogen and carbon atoms effectively have full outer shells in forming the covalent bonds when the atoms share their outer electrons.

The simplified 'dot and cross' electronic diagram for the covalently bonded CO2 molecule is the full 'dot and cross' electronic diagram for the covalent bonding in the carbon dioxide molecule.

The molecule can be shown as  (displayed formula) with two carbon = oxygen double covalent bonds (called a linear shape, the O=C=O bond angle is 180o). Valencies of C and O are 4 and 2 respectively.

The carbon dioxide molecule is held together by the strong C=O carbon–oxygen double covalent bonds by sharing electrons.

Example 3: Covalent bonding diagram for ETHENE covalent molecule, molecular formula C2H4

Two atoms of carbon (2.4) combine with four atoms of hydrogen (1) to form ethene C2H4 (only the outer shell of carbon’s electrons is shown).

Electronically, hydrogen (1) becomes like helium (2) and carbon (2.4) becomes like neon (2.8), so ALL the hydrogen and carbon atoms effectively have full outer shells in forming the covalent bonds when the atoms share their outer electrons.

The molecule can be shown as  (displayed formula) with one carbon = carbon double bond and four carbon – hydrogen single covalent bonds (called a planar shape, its completely flat!, the H–C=C and H–C–H bond angles are 120o). The valency of carbon is still 4.

The ethane molecule is held together by the four strong C–H carbon–hydrogen single covalent bonds and one C=C carbon–carbon covalent double bond. formula

Alkanes are relatively small molecules in which all the chemical bonds are covalent bonds. All the bonds in alkane molecules are single bonds i.e. C–C carbon – carbon or C–H carbon – hydrogen. Each carbon atom forms four single covalent bonds and hydrogen atoms form one single covalent bond. All single bonds are formed by sharing a pair of electrons e.g. one from each of a carbon atom and a hydrogen atom, or two carbon atoms contributing (sharing) an electron each to the covalent bond.

Notes on shapes and bond angles:

o Ethane has a linked double tetrahedral shape with sp3 and all H–C–H and H–C–C bond angles are 109o

o Chloromethane has tetrahedral shape with  H–C–H and H–C–Cl bond angles of approximately 109o having same sp3 hybridization

o Methanol, the four bonds around the central carbon are tetrahedrally arranged with a H 'wiggle' on the oxygen. All the H–C–H, H–C–O and C–O–H bond angles are approximately 109o

o The given compounds represent an octahedral shape (e.g. SF6, complex transition metal ions like [Cu(H2O)6]2+ and the bond angles are either 90o or 180o

o Simple molecules with a triple bond are often linear e.g. H–C C–H ethyne or H–C N hydrogen cyanide (methanenitrile)

Covalently bonded molecules and properties

Why do simple covalent molecules typically have low melting and low boiling points?o Typical examples are water (ice & liquid), petrol, butter etc.o The particles in the above diagram represent whole molecules, but the general picture of particles in

the three states of matter help to understand the properties of simple molecular substances. Why do simple molecules NOT usually conduct electricity even when liquid/molten/dissolved.

The first point to appreciate is that the chemical bonding forces between atoms in a molecule are strong BUT the bonding forces between small simple molecules are weak. These weak electrical attractive forces are known as 'intermolecular forces' or 'intermolecular bonding'.

o DO NOT CONFUSE THESE TWO FORCES or it makes the following discussion on the physical properties of simple molecules difficult to follow.

o The contrast between the strong bonds between atoms in a molecule and the weak bonds between individual molecules is really important to know understand the consequences.

o It will also help you to understand why covalent giant molecular structures have very different physical properties.

The electrical forces of attraction, that is the chemical bond, between atoms in a molecule are usually very strong, so,  most covalent molecules do not change chemically on moderate heating.

o e.g. although a covalent molecule like iodine, I2, is readily vaporized on heating, it does NOT break up into iodine atoms I. The purple vapour you see on heating iodine is entirely composed of the diatomic I2 molecules.

o The I–I covalent bond is strong enough to withstand the heating and so the purple vapour still consists of the same I2 molecules as the dark coloured solid is made up of.

o In other words, on heating a simple molecular material like iodine, heating weakens the forces between the molecules BUT not the forces between the atoms in the molecule.

Chemical bonds between atoms are generally only broken if a substance is heated to a VERY high temperature like in the cracking break–down reactions of alkanes from crude oil.

So why the ease of vaporisation on heating? o Although the bonding between the atoms within a molecules is very strong the electrical attractive

force between individual molecules is very weak, so the bulk material is not very strong physically and this has consequences for the melting points and boiling points.

o

If you take the Group 7 Halogen molecules, the F–F, Cl–Cl, Br–Br, I–I covalent bonds (–) are very strong,

but the F–F....F–F, Cl–Cl....Cl–Cl, Br–Br....Br–Br and I–I....I–I intermolecular bonds are weak (weak intermolecular bonding shown as ....),

resulting in low melting/boiling points e.g. at room temperature fluorine and chlorine are gases, bromine a low boiling liquid and iodine an easily vaporised solid on gentle heating.

Note: The bigger the molecule, the stronger the intermolecular forces, which is why the melting/boiling points increase down group 7.

Similarly for hydrocarbons like alkanes, the longer the molecule, the higher the boiling point.

These weak electrical attractions are known as intermolecular forces (or intermolecular bonding) and are readily weakened further on heating. In a solid, the effect of absorbing heat energy results in increased the thermal vibration of the molecules which weakens the intermolecular forces. In liquids the increase in the average particle kinetic energy makes it easier for molecules to overcome the intermolecular forces and change into a gas or vapour.  Consequently, small covalent molecules tend to be volatile liquids with low boiling points, so easily vaporized, or low melting point solids.

o So, on heating simple molecular substances (small molecules) the inter–molecular forces are easily overcome with the increased kinetic energy of the particles, giving the material a low melting or boiling point because a relatively low value of kinetic energy is needed to effect these state changes.

o This contrasts with the high melting points of giant covalent structures with their strong 3D network.o Reminder: The weak electrical attractive forces between molecules, the so called intermolecular

forces should be clearly distinguished between the strong covalent bonding between atoms in molecules (small or giant), and these are sometimes referred to as intramolecular forces (i.e. internal to the molecule).

Covalent structures are usually poor conductors of electricity because there are no free electrons or ions in any state to carry electric charge.

Most small molecules will dissolve in some solvent to form a solution. o This again contrasts with giant covalent structures where the strong bond network stops solvent

molecules interacting with the particles making up the material.o Hydrocarbon molecules like hexane or paraffin wax dissolve in organic solvents but not water, but

sugars are also low melting small covalent molecules but do dissolve in water, insoluble in hydrocarbon solvents.

The properties of these simple small molecules should be compared and contrasted with those molecules of a giant covalent nature (next section).

o Apart from points on the strong bonds between the atoms in the molecule and the lack of electrical conduction, all the other properties are significantly different!

Intermolecular forces – intermolecular bonding

Using water as an example

Between all particles, but with particular reference to covalently bonded molecules, there always exists some very weak electrical attractive forces known as intermolecular forces or intermolecular bonding.

These constantly acting attractive forces or intermolecular bonds are very much weaker than covalent or ionic chemical bonds (approximately 1/30 to 1/20th in comparative attractive force).

For example, although the oxygen and hydrogen atoms are very strongly bonded in water to make a VERY stable molecule, BUT this does NOT account for the existence of liquid water and ice!

It is the weak intermolecular forces that induce condensation below 100oC and freezing–solidification to form ice crystals below 0oC.

In the reverse process, when ice is warmed, the intermolecular forces are weakened and at 0oC the intermolecular bonds are weakened enough to allow melting to take place.

Above 0oC (evaporation), and particularly at 100oC (boiling), the intermolecular forces are weak enough for 'intact water molecules' to escape from the surface of the liquid water.

It is VERY important to realise that the chemical hydrogen–oxygen covalent bonds (O–H) in water are NOT broken and the state changes ...

solid <== freezing/melting ==> liquid <== condensing/boiling ==> gas ...

are due to the weakening of the intermolecular forces/bonds with increase in temperature OR the strengthening of the intermolecular bonds/forces decrease in temperature.

The same arguments apply to all the other small covalent molecules you will come across on your course eg methane, iodine, carbon dioxide, alkanes like hexane in petrol etc. etc.