Curriculum

Oxford AQA International A Level Chemistry

1. Physical Chemistry

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2. Inorganic Chemistry

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3. Organic Chemistry

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Text lesson

1.3 Bonding (Part 3)

1.3 Bonding (Part 3) — OXFORD AQA INTERNATIONAL A LEVEL CHEMISTRY — COMPLETE NOTES notes: Compare the types of intermolecular force, explain boiling points, and describe hydrogen bonding in water and ice.

Force	Where it occurs	Relative strength
van der Waals	Between all atoms and molecules	Weakest
Permanent dipole–dipole	Only between polar molecules	Stronger than van der Waals
Hydrogen bonding	Between molecules with H bonded to N, O or F	Strongest

Note
- intermolecular forces act between molecules and should not be confused with covalent bonds

Mechanism

Mechanism the δ- end of one polar molecule is attracted to the δ+ end of a neighbour

Effect

Effect boiling points are slightly higher than expected from temporary dipoles alone

Comparison

Comparison for small molecules with the same number of electrons, dipole–dipole attraction is stronger than dispersion forces

	Butane	Propanone
Electrons	34	34
Polarity	Non-polar	Polar
Forces	van der Waals	van der Waals and dipole–dipole
Boiling point	-1 °C	56 °C

Where

Where they exist between all atoms and molecules, and are also called London dispersion forces

Origin

The electron cloud moves constantly, so it can be denser on one side, giving a temporary dipole.
This temporary dipole induces a dipole in a neighbour.
The δ+ and δ- ends attract, but the dipoles are only temporary.

Strength depends on

The number of electrons, since more electrons give stronger forces.
The relative molecular mass, since higher mass gives stronger forces.

Example

Example pentane (C₅H₁₂) has a higher boiling point than propane (C₃H₈)

Hydrogen bonding

Hydrogen bond

Hydrogen bond the strongest intermolecular force

Conditions

Conditions a hydrogen atom bonded to a very electronegative O, N or F, which polarises the bond

Mechanism

Mechanism the H becomes so δ+ that it bonds with a lone pair on an O, N or F atom in another molecule

Water

Water each molecule can form up to four hydrogen bonds, because oxygen has two lone pairs and there are two δ+ hydrogens

Between different molecules

Between different molecules hydrogen bonds form between water and ammonia, since each contains O–H or N–H

High melting and boiling points of water

Cause

Cause hydrogen bonding gives water anomalous properties such as high melting and boiling points, needing a lot of energy to separate molecules

Group 16 hydrides

Group 16 hydrides the enthalpy of vaporisation increases from H₂S to H₂Te because more electrons give stronger dispersion forces

Water as the exception

Water as the exception on that trend it should be about 17 kJ mol^-1, but its value is almost three times larger because of hydrogen bonds

Conclusion

Conclusion water has hydrogen bonds as well as dispersion forces

Ice

Structure

Structure when water freezes, the molecules are fixed in a three-dimensional crystal

Spacing

Spacing the long hydrogen bonds hold the molecules slightly further apart than in liquid water

Density

Density ice is less dense than water, so ice floats

Term	Meaning
van der Waals forces	Weak attractions between all molecules from temporary and induced dipoles.
Permanent dipole–dipole force	The attraction between permanent dipoles of polar molecules.
Hydrogen bond	A strong attraction between a δ+ H on N, O or F and a lone pair on another N, O or F.
Dispersion force	Another name for induced dipole–induced dipole force.
Enthalpy of vaporisation	The energy needed to turn a liquid into a gas.

Confusing intermolecular forces with covalent bonds — covalent bonds are within molecules and do not break on boiling a simple molecule.

Saying hydrogen bonds form with any H — the H must be bonded to O, N or F.

Saying water follows the dispersion-force trend for hydrides — hydrogen bonding makes it exceptional.

Saying ice is denser than water — it is less dense because of the open hydrogen-bonded structure.