Shapes of Molecules and Intermolecular Forces (OCR A-Level Chemistry A): Revision Notes

Hydrogen bonding

What is a hydrogen bond?

A hydrogen bond is a special type of permanent dipole-dipole interaction that forms between specific molecules. It is the strongest type of intermolecular force, making it particularly important in chemistry and biology.

For a hydrogen bond to form, two key requirements must be met:

1. One molecule must contain an electronegative atom with a lone pair of electrons (specifically fluorine, oxygen, or nitrogen)
2. Another molecule must have a hydrogen atom covalently bonded to an electronegative atom (forming bonds like $\ce{O-H}$, $\ce{N-H}$, or $\ce{F-H}$)

The hydrogen bond forms between the lone pair of electrons on the electronegative atom in one molecule and the hydrogen atom in a different molecule. Because the electronegative atom pulls electrons away from the hydrogen in a covalent bond, the hydrogen becomes partially positive ($\delta^+$) while the electronegative atom becomes partially negative ($\delta^-$). This creates the attraction that we call a hydrogen bond.

Representing hydrogen bonds

When drawing hydrogen bonds in diagrams, we use specific conventions to make them clear:

• Hydrogen bonds are shown as dashed or dotted lines (not solid lines like covalent bonds)
• Partial charges must be indicated: $\delta^+$ on hydrogen atoms and $\delta^-$ on electronegative atoms
• The arrangement around the hydrogen atom involved in the hydrogen bond is linear (approximately 180°)
• Always show lone pairs of electrons on the electronegative atoms, as these are essential to the hydrogen bond formation

Anomalous properties of water

Hydrogen bonding has a profound influence on the properties of many molecules, but its effects are most dramatic in water ($\ce{H2O}$). Water exhibits several unusual (anomalous) properties that are directly caused by hydrogen bonding, and these properties are essential for supporting life on Earth.

Ice is less dense than water

One of water's most remarkable properties is that solid ice is less dense than liquid water, causing ice to float. This behaviour is highly unusual—for most substances, the solid form is denser than the liquid and therefore sinks.

This phenomenon occurs because of the way hydrogen bonds arrange water molecules in ice:

• In solid ice, hydrogen bonds hold water molecules in a rigid open tetrahedral lattice structure
• Each water molecule forms four hydrogen bonds with neighbouring molecules (using its two lone pairs as acceptors and its two $\ce{O-H}$ bonds as donors)
• These hydrogen bonds hold the molecules slightly apart from each other, creating a structure with large gaps or "holes"
• The bond angle around each hydrogen atom involved in hydrogen bonding is close to 180°, contributing to this open arrangement

When ice melts, the rigid lattice structure collapses. The hydrogen bonds break (though not completely—they reform and break rapidly in liquid water), and the molecules can move closer together. As a result, liquid water is actually denser than solid ice.

Density comparison:

• Ice at 0°C: 0.917 g cm⁻³
• Liquid water at 0°C: 1.029 g cm⁻³

This means that only about $\frac{1}{9}$ of an iceberg's volume sits above the water surface—the rest is hidden below.

Water has relatively high melting and boiling points

Water has much higher melting and boiling points than would be expected for a molecule of its size. This is because hydrogen bonds provide additional intermolecular forces beyond the London forces present in all molecules.

Energy considerations:

• All molecules experience London (dispersion) forces between them
• Water molecules also experience strong hydrogen bonding
• To melt or boil water, energy must be supplied to overcome both London forces and hydrogen bonds
• This means considerably more energy is needed than for similar-sized molecules without hydrogen bonding

What happens during phase changes:

• When the ice lattice melts, the rigid arrangement of hydrogen bonds is broken, but hydrogen bonds between water molecules continue to form and break rapidly in the liquid state
• When water boils, the hydrogen bonds break completely and water molecules separate into the gas phase
• The extra energy required to break these hydrogen bonds results in higher melting and boiling points

Boiling points of hydrides

The effect of hydrogen bonding becomes particularly clear when we examine the boiling points of hydrides across Groups 14-17 in the periodic table. A hydride is simply a compound of hydrogen with another element.

Observed trends:

The graph reveals several important patterns:

1. General trend: For most hydrides, boiling point increases going down each group (from Period 2 to Period 6). This occurs because as the central atom gets larger, there are more electrons present, leading to stronger London forces between molecules.

2. Anomalous behaviour: Three compounds show dramatically higher boiling points than expected:

   • $\ce{H2O}$ (water)
   • $\ce{HF}$ (hydrogen fluoride)
   • $\ce{NH3}$ (ammonia)

3. Explanation: These three compounds contain hydrogen bonded to nitrogen, oxygen, or fluorine—precisely the conditions needed for hydrogen bonding. The additional strength from hydrogen bonding causes their boiling points to be much higher than predicted from the group trend.

4. No hydrogen bonding in other hydrides: Compounds like $\ce{CH4}$, $\ce{SiH4}$, $\ce{H2S}$, and $\ce{HCl}$ don't exhibit hydrogen bonding because they lack sufficiently electronegative atoms (carbon, silicon, sulfur, and chlorine aren't electronegative enough, despite chlorine being quite electronegative).

Other anomalous properties of water

The extensive hydrogen bonding network in water is responsible for several additional unusual properties:

Surface tension:

Water has a relatively high surface tension compared to other liquids of similar molecular mass. The hydrogen bonds between water molecules at the surface create a "film" effect. This property allows:

• Some insects to walk on water surfaces without sinking
• Water droplets to form beads on surfaces rather than spreading out completely

Detergents reduce surface tension by interfering with hydrogen bonding, which is why they make water "wetter" and better at cleaning.

Viscosity:

Water has higher viscosity than expected for its molecular size. Viscosity is the resistance to flow—honey has high viscosity while petrol has low viscosity. The network of hydrogen bonds between water molecules creates resistance to flow, though water still flows quite easily compared to many other liquids.

These dozens of unusual properties make water a remarkable solvent and an essential substance for life, though for A-Level chemistry you primarily need to focus on density, melting point, and boiling point.

Hydrogen bonding in biological systems

Hydrogen bonding plays crucial roles beyond water chemistry. One particularly important example is in DNA structure.

The famous double helix structure of DNA is held together by hydrogen bonds between complementary base pairs:

• Adenine (A) pairs with thymine (T) forming two hydrogen bonds
• Cytosine (C) pairs with guanine (G) forming three hydrogen bonds

These hydrogen bonds are strong enough to hold the DNA strands together, but weak enough to be separated when DNA needs to replicate. This balance is essential for genetic information storage and transfer.

You'll encounter many more examples of hydrogen bonding when you study organic chemistry and biochemistry later in your course, particularly in alcohols, carboxylic acids, and amino acids.