Electronegativity, bond polarity and dipoles

· Electronegativity = the power of an atom to attract electrons to itself in a bond.
· If bonded atoms have different electronegativities, the bonding pair is pulled closer to the more electronegative atom.
· This creates a polar bond with partial charges: δ− on the more electronegative atom and δ+ on the less electronegative atom.
· A dipole moment is caused by separation of charge in a polar bond or molecule.
· A molecule has an overall dipole if its bond dipoles do not cancel due to molecular shape.
· A molecule may contain polar bonds but be non-polar overall if the dipoles cancel symmetrically.

Water is polar because oxygen is more electronegative than hydrogen, making O δ− and H δ+. The bent shape means the O–H bond dipoles do not cancel, producing an overall molecular dipole. Source

Van der Waals’ forces: overview

· Van der Waals’ forces = intermolecular forces between molecular entities other than those due to bond formation.
· In CIE, van der Waals’ forces can be used as a generic term for all intermolecular forces.
· Main types required: instantaneous dipole–induced dipole forces, permanent dipole–permanent dipole forces, and hydrogen bonding.
· Intermolecular forces act between molecules, not within covalent molecules.
· Stronger intermolecular forces usually mean higher melting point and higher boiling point because more energy is needed to separate molecules.

This page compares the main intermolecular forces found between molecules. It is useful for seeing how dispersion, dipole–dipole attraction and hydrogen bonding differ in origin and strength. Source

Instantaneous dipole–induced dipole forces / London dispersion forces

· Instantaneous dipole–induced dipole forces are also called London dispersion forces.
· They occur due to temporary uneven electron distribution creating an instantaneous dipole.
· This instantaneous dipole can induce a dipole in a nearby molecule.
· These forces exist between all atoms and molecules, including non-polar molecules.
· They are generally weak but become stronger when molecules have more electrons, larger electron clouds, or greater surface contact.
· Exam phrasing: larger molecule → more electrons → stronger id-id forces → higher boiling point.

Permanent dipole–permanent dipole forces

· Permanent dipole–permanent dipole forces occur between polar molecules with permanent dipoles.
· The δ+ end of one molecule attracts the δ− end of another molecule.
· They are usually stronger than id-id forces for molecules of similar size.
· A molecule must be polar overall, not just have polar bonds.
· Exam tip: always consider both bond polarity and molecular shape when deciding if pd-pd forces are present.

Hydrogen bonding

· Hydrogen bonding is a special case of permanent dipole–permanent dipole force.
· In this syllabus, hydrogen bonding is limited to molecules containing N–H and O–H groups.
· Required examples: ammonia, NH₃, and water, H₂O.
· Hydrogen bonding occurs when H is bonded to a highly electronegative atom and is attracted to a lone pair on N or O in a neighbouring molecule.
· Represent hydrogen bonds using dotted lines between molecules, not as normal covalent bonds.
· Hydrogen bonds are stronger than ordinary pd-pd forces but still usually weaker than ionic, covalent and metallic bonding.

The diagram shows hydrogen bonding between water molecules. The δ+ hydrogen of one water molecule is attracted to the δ− oxygen of another water molecule. Source

Anomalous properties of water explained by hydrogen bonding

· Water has a relatively high melting point and boiling point because extensive hydrogen bonding requires extra energy to overcome.
· Water has relatively high surface tension because hydrogen bonds hold surface water molecules together strongly.
· Ice is less dense than liquid water because hydrogen bonds hold water molecules in a more open structure in the solid state.
· In liquid water, hydrogen bonds are constantly breaking and reforming, allowing molecules to pack closer than in ice.
· Exam phrasing: ice floats because its open hydrogen-bonded lattice makes it less dense than liquid water.

Ice has a more open hydrogen-bonded structure than liquid water. This open lattice explains why solid ice is less dense than liquid water and therefore floats. Source

Relative strength of bonding and intermolecular forces

· In general, ionic bonding, covalent bonding and metallic bonding are stronger than intermolecular forces.
· Covalent bonds are strong attractions within molecules; intermolecular forces act between molecules.
· Melting or boiling simple molecular substances usually involves overcoming intermolecular forces, not breaking covalent bonds.
· Common strength order for exam explanations: covalent / ionic / metallic bonding > hydrogen bonding > pd-pd forces > id-id forces.
· Do not say “covalent bonds are broken during boiling” for simple molecular substances unless the molecule chemically decomposes.

Exam comparison phrases

· Higher boiling point: stronger intermolecular forces require more energy to overcome.
· Hydrogen bonding present: molecule contains O–H or N–H groups and has lone pairs on electronegative atoms.
· Polar molecule: contains polar bonds and has a shape where dipoles do not cancel.
· Non-polar molecule: bond dipoles are absent or cancel due to symmetry.
· Larger non-polar molecule: stronger id-id forces due to more electrons and a more polarisable electron cloud.