Bonding (sl)

4.1 - Ionic Bonding

Ionic bond - +ve (cations) and -ve (anions) ions are attracted to each other and form a continuous giant ionic lattice.

Ions are formed from elements by loss (metals) or gain (non-metals) of electrons to attain a full outer shell.

group number 1 2 transition metals 3 4 5 6 7 0
ions formed +1 +2 many positive ions +3 no ions -3 -2 -1 no ions
examples Li+ Mg2+ Fe2+, Fe3+ Al3+ - N3- O-2 Cl- -

Greater ease of ionisation Li->Cs is due to the increased electron shielding of the nuclear attraction caused by additional inner shells of electrons. This may also be considered from the point of view of more electrons means increased inter-electron repulsion and consequent lower ionisation energy.

The easier atoms are to ionise, the more reactive they will be because less energy is required to ionise them, and so they react more easily.

The transitions metals (elements from Ti to Cu, ignore Sc and Zn) can form multiple ions due to close proximity (energetically) of 4s and 3d shells which means that they can lose both 's' electrons but they can also lose variable numbers of 'd' electrons

Factors affecting lattice energy

Naming simple ions

Positive ions take the same name as the metals they come from, but negative ions use the name of the element as a basis and change the ending to -ide (for binary compounds) .

examples: fluoride, chloride, bromide, iodide, oxide, sulphide, nitride and phosphide

Physical properties

Property Ionic material characteristics Explanation
m.p. High melting points Many strong electrostatic attractions to be broken
Conductivity Conductors only when molten or in solution The ions can carry the electric charge when they are free to move
Hardness Hard but brittle Ions are held rigidly in position in the lattice. Stress brings ions of the same charge in close proximity and the structure breaks along cleavage planes
Solubility Mostly soluble The charged ions can be carried off by the polar water molecules unless the electrostatic attractions within the lattice are too large
Structure Giant lattice of repeating ions in three dimensions Electrostatic attraction between oppositely charged ions causes the negative ions to surround the positive ions and vice versa.
Bonding section 4.1  

4.2 -Covalent Bonding

Covalent bonds are where two atoms each donate 1 electron to form a pair held between the two atoms.

In terms of minimising chemical potential energy

Non-metal atoms bonded to other non-metal atoms

Difference in electronegativity between the two bonded atoms causes the bond to be polar and the electron pair will be held closer to the more electronegative atom (there is charge separation producing partial positive and negative charges with the partial negative beinbg on the more electronegative atom)

The shape of molecules with 4 electron pairs depends on the number of lone pairs.

number of lone pairs 2 lone pairs 1 lone pair No lone pairs
shape bent or angular trigonal pyramid tetrahedral
example (view) water ammonia methane

The polarity of a molecule depends on both the shape and the polarity of the bonds...

1) if there are no polar bonds, it's not polar.

2) if there are polar bonds, but the shape is symmetrical, it's not polar (think about it like 3D vector addition...if they add to zero, then it's not polar).

3) if there are polar bonds, and it's not symmetric, then the molecule is polar.

Bonding section 4.2  

4.3 - Intermolecular forces

Van der Waal's forces -- Electrons will not be evenly spread around an atom/molecule at any given time, meaning the molecule will have a slight +ve charge on one end, and a -ve at the other. this temporary state may cause attraction between two molecules, pulling them together (also known as London dispersion forces). The magnitude pof Van der Waals force depends on the relative molecular mass, high mass produces a larger force.

Boiling points of the alkanes data and alkenes data

These very clearly illustrate the effect of increasing Van der Waals attractions as the relative molecular mass increases. The influence of branching in the alkanes can also illustrate the effect of different surface areas on the Van der Waala forces (the more branching the lower the b.p)

Dipole-dipole forces -- Polar molecules, when properly oriented, will attract each other as a result of this. Stronger than van der Waal's forces.

Hydrogen bonding -- When hydrogen is bonded to nitrogen, oxygen or fluorine, a very strong dipole is formed, making the hydrogen very strongly positive. This hydrogen is then attracted to the lone pairs on other similar molecules (nitrogen, oxygen and fluorine all have lone pairs) forming a hydrogen bond, which is stronger than van der Waal's or dipole-dipole, but weaker than covalent bonding.

The effect of hydrogen bonding on intermolecular forces can be demonstrated very well by studying the boiling points of the group 6 hydrides

b.p. comparison of main group hydrides

Order of priority

  1. Hydrogen bonding strongest
  2. Dipole -dipole interactions
  3. Van der Waals forces

Hydrogen bonds result from hydrogen bonded as described above. This results in molecules with hydrogen bonding exhibiting stronger intermolecular forces, ie higher boiling/melting points etc. eg H2O has a higher bp then H2S due to hydrogen bonding, and so on down the strength list.


Water has a very high melting and boline point due to extensive hydrogen bonding. Having two hydrogen it can form two H- bonds per molecule allowing a lattice diamond-like structure to be built up, as shown below in ice:

Ice structure showing the intermolecular hydrogen bonding in red, holding the molecules of water into a 'diamond' like ring structure. The blue bonds represent the covalent bonds within the water molecules (intramolecular)

Polar molecules: There are dipole-dipole forces that arise from polar bonds and asymmetry in molecules.

Non-polar molecules: These have only Van der Waal's forces (induced dipole dipole interactions) which are also present in all other molecules, though the strength of the intermolecular force may be insignificant compared to the other forces although this depends on the relative molecular mass.

Bonding section 4.3  

4.4 -Metallic bonding

Metallic bonding...the metal atoms "lose" their outer electrons which then become delocalised, and free to move throughout the entire metal. These negative delocalised electrons hold the metal cations together strongly. Since these electrons can flow, atoms with metallic bonding exhibit high electrical conductivity. Unlike ionic bonding, distorting the atoms does not cause repulsion so metallic substances are ductile (can be stretched into wires) and malleable (can be formed into shapes). The free moving electrons also allow for high thermal conductivity, and the electrons can carry the heat energy rather than it being transferred slowly through atoms vibrating.

Bonding section 4.4  

4.5 - Physical Properties

Melting + Boiling point...High with Ionic and metallic bonding (and network covalent), Low with covalent molecular bonding.

Expt: Determination of boiling point


Covalent molecular substances are volatile, others aren't


Metallic substances conduct. Molecular substances do not conduct (graphite is an exception). Ionic substances don't conduct when solid, do conduct when molten or dissolved in water.


Ionic substances generally dissolve in polar solvents (like water). Non-polar molecules are generally soluble in non-polar solvents, and polar in polar. "Like dissolves like"

For a full appreciation of the reasons why "like dissolves like" just consider that it is a balance between the forces holding a solid together and the force of attraction between the liquid particles and the solid particles. If the force of attraction of the liquid particles for the solid particles is greater than the force of attraction felt by the solid particles for each other then the solid dissolves. More...

Organic molecules with a polar head : Short chain molecules are soluble in polar solvents, long chains can eventually outweigh the polar 'head' and will dissolve in non-polar solvents.

Bonding section 4.5  

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