Shapes, Polarity & Intermolecular Forces (AQA A Level Chemistry): Video

Hi, I'm Eleanor with 3 years of experience teaching chemistry, and this video is about the shapes of simple molecules and ions, bond polarity, and the forces between molecules.

Shape and bond polarity describe what happens inside a molecule; intermolecular forces act between molecules, and it is those intermolecular forces that set a substance's melting and boiling points.

A molecule has a shape, set by the electron pairs around its central atom. And bonds that may be polar, set by the electronegativities of the atoms joined by the bond. Both are features inside the molecule. Separate from them are the forces that act between whole molecules. Melting and boiling separate molecules from one another, so it is the strength of those intermolecular forces that decides a molecular substance's melting and boiling points.

We start with VSEPR and the shapes of simple molecules and ions. Then bond polarity and electronegativity. Then the difference between forces inside a molecule and forces between molecules, and the three types of intermolecular force. We finish with what those forces do to melting point, boiling point, surface tension and density.

Valence shell electron pair repulsion theory, VSEPR, predicts the shape and the bond angles of a molecule. Electron pairs are negatively charged, so they repel each other, and the molecule adopts the shape that minimises those repulsions. Lone pairs repel more strongly than bonding pairs, because their electron clouds are more concentrated and lie closer to the nucleus. So the order of repulsion from most to least runs: lone pair to lone pair, then lone pair to bond pair, then bond pair to bond pair. Four bonding pairs give a tetrahedral shape at 109.5 degrees; replace one bonding pair with a lone pair for a trigonal pyramidal shape at 107 degrees, and two lone pairs for a bent shape at 104.5 degrees. Shape is the first of the two features inside the molecule that the rest of this video builds on.

Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond towards itself. When both atoms have the same electronegativity, the electrons are shared equally and the bond is non-polar. When electronegativity differs, the electrons are drawn towards the more electronegative atom, which takes a partial negative charge, delta minus, while the other takes delta plus. The bigger the difference, the more polar the bond. Electronegativity increases across a period and decreases down a group, so the general trend is an increase towards the top right of the periodic table, with fluorine the most electronegative element. Bond polarity is the second feature inside the molecule, and it is what decides whether permanent dipoles form between molecules.

Intramolecular forces are forces within a molecule, and they are usually covalent bonds, formed when the outer electrons of two atoms are shared. Intermolecular forces are the weaker forces between molecules: a hydrogen bond is about one tenth the strength of a covalent bond. In order of strength, covalent bonding is strongest, then hydrogen bonding, then permanent dipole to permanent dipole, then instantaneous dipole to induced dipole. This is the distinction the rest of the video rests on: melting and boiling separate molecules from one another, so it is the intermolecular forces that have to be overcome.

There are three types of intermolecular force. Induced dipole to dipole forces, also called van der Waals, London or dispersion forces, exist between all atoms and molecules: a constantly moving electron charge cloud can concentrate on one side of a molecule, creating a temporary dipole that induces a dipole in its neighbour.

Permanent dipole to dipole forces act between polar molecules, where the delta plus end of one molecule is attracted to the delta minus end of the next. For molecules with the same number of electrons, permanent dipoles are the stronger of the two: butane is non-polar and boils at 0 degrees celcius, while propanone is polar and boils at 56 degrees celcius.

Hydrogen bonding is the strongest form of intermolecular bonding, and it needs hydrogen covalently bonded to either oxygen, nitrogen or fluorine: that bond is so polarised that the hydrogen can bond to a lone pair on an O, N or F atom in another molecule. Which of these forces act in a substance decides how much energy is needed to separate its molecules.

More energy is needed to break stronger intermolecular forces, so stronger forces mean a higher boiling point. Water shows this clearly. Hydrogen bonds between water molecules are strong, so a lot of energy is needed to separate them, and water has high melting and boiling points for a molecule of its size. On the enthalpy of vaporisation trend for the group 16 hydrides, water should be around 17 kilojoules per mole, but it is almost three times that, and hydrogen bonding is the reason. The same bonds give water its high surface tension, because molecules at the surface are pulled downwards by hydrogen bonds to the molecules below. They also hold ice in an open lattice, which is why ice is about 9 percent less dense than liquid water.

To explain a reduced bond angle, start from the parent geometry, 109.5 degrees for tetrahedral, and deduct 2.5 degrees per lone pair: ammonia 107, water 104.5. Examiners report that students wrongly start from 180 degrees. Always name VSEPR and state the repulsion order, because a common lost mark is not mentioning VSEPR at all. When you draw intermolecular forces, use a dashed straight line, starting at the lone pair and going right up to the delta positive atom.

VSEPR sets the shape and bond angles of a molecule from the repulsion between its electron pairs. A difference in electronegativity makes a bond polar, with delta minus on the more electronegative atom. Between molecules, three forces act: induced dipole, permanent dipole, and hydrogen bonding. And it is the strength of those intermolecular forces that sets the melting and boiling points of the substance, along with water's surface tension and the low density of ice.