3q. Describing and explaining the properties of small covalently bonded molecules using a molecular model

e.g. relatively low melting point and boiling point, physically weak if crystalline  at room temperature and lack of electrical conduction

Typical properties of simple molecular substances

Pd

metals

Part of the modern Periodic Table Pd = period, Gp = group

metals => non–metals

Gp1

Gp2

Gp3

Gp4

Gp5

Gp6

Gp7

Gp0

1

1H  Note that hydrogen does not readily fit into any group but is a non-metal

2He

2

3Li

4Be

atomic number Chemical Symbol eg 4Be

5B

6C

7N

8O

9F

10Ne

3

11Na

12Mg

13Al

14Si

15P

16S

17Cl

18Ar

4

19K

20Ca

21Sc

22Ti

23V

24Cr

25Mn

26Fe

27Co

28Ni

29Cu

30Zn

31Ga

32Ge

33As

34Se

35Br

36Kr

5

37Rb

38Sr

39Y

40Zr

41Nb

42Mo

43Tc

44Ru

45Rh

46Pd

47Ag

48Cd

49In

50Sn

51Sb

52Te

53I

54Xe

6

55Cs

56Ba

Transition Metals

81Tl

82Pb

83Bi

84Po

85At

86Rn

The non-metals in the Periodic Table highlighted in white, tend to form covalent bonds with each other. This might be as an element or a compound, usually forming small molecules, but sometimes a giant structure This combination of two non-metallic elements usually results in relatively small covalently bonded molecules

Why do simple covalent molecules typically have low melting and low boiling points?

• The forces between the individual molecules are weak and easily overcome on heating.

• These weak electrical attractive forces are called 'intermolecular forces' or 'intermolecular bonding'.

  • It is these weak forces that determine the bulk properties of this type of covalent compound.

  • There are more detailed discussions about these forces later on this section.

• Typical examples of low melting simple covalent molecules are water (ice & liquid), petrol, butter etc.

• 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.

• 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'.

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

  • 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.

  • 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.

• e.g. although a covalent molecule like iodine, I₂, is readily vapourised 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.

• 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.

• 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?

  • 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.

    halogen molecules

• 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 (the weak intermolecular forces 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 vapourised 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 arguments apply to all relatively small molecules BUT not to giant covalent structures like carbon as diamond or silica.

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 vapourised, or low melting point solids.

• 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.

• Energy changes for the physical changes of state of melting and boiling for a range of differently bonded substances are compared in a section of the Energetics Notes.

• This contrasts with the high melting points of giant covalent structures with their strong 3D network of bonds.

• 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).

Why do simple covalent molecules NOT usually conduct electricity even when liquid, molten or dissolved?

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

• The molecules do NOT have an overall electric charge, they are electrically neutral and cannot form part of an electric current, which is a flow of charged particles.

• Most small molecules will dissolve in some solvent to form a solution.

  • This again contrasts with giant covalent structures where the strong bond network stops solvent molecules interacting with the particles making up the material.

  • 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).

  • 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!

• Note that bigger molecules have bigger intermolecular forces of attraction, so higher kinetic energies, at higher temperature are needed to overcome these intermolecular bonding - so they have much higher boiling points.

A little more on 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 induces condensation below 100^oC and freezing–solidification to form ice crystals below 0^oC.

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

Above 0^oC (evaporation), and particularly at 100^oC (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.