Topic F: Microscopic World II

Microscopic World II deepens the bonding picture by explaining why bonds and molecules can be polar, and how weak forces between molecules control melting points, boiling points and solubility. It then contrasts giant and simple molecular structures, drawing together the structure-property links that run through the whole syllabus.

Electronegativity and bond polarity

Electronegativity is a measure of how strongly an atom attracts the shared pair of electrons in a covalent bond. It increases across a period and decreases down a group, peaking at fluorine. When two bonded atoms have different electronegativities, the shared electrons are pulled toward the more electronegative atom, giving it a small negative charge and leaving the other atom slightly positive. This separation of charge is a polar bond, or dipole. If both atoms are identical, as in Cl2, the bond is non-polar. The greater the difference in electronegativity, the more polar the bond, and a very large difference shifts the bonding toward ionic character. Bond polarity underlies molecular polarity, solubility and intermolecular forces.

Polar and non-polar molecules

Whether a whole molecule is polar depends on both its bond polarities and its shape. If the individual bond dipoles point in directions that cancel by symmetry, the molecule has no overall dipole and is non-polar. Carbon dioxide has two polar C=O bonds but is linear and symmetric, so the dipoles cancel and the molecule is non-polar. Tetrachloromethane, CCl4, is tetrahedral and symmetric and is also non-polar. Water has polar O-H bonds and a bent shape, so the dipoles add together and the molecule is strongly polar. Polar molecules dissolve readily in polar solvents such as water, while non-polar molecules dissolve in non-polar solvents, the rule that like dissolves like.

Van der Waals (dispersion) forces

All molecules attract one another through weak van der Waals forces, also called London dispersion forces. They arise because the constant motion of electrons creates instantaneous, temporary dipoles that induce opposite dipoles in neighbouring molecules, producing a fleeting attraction. These forces are weak but become stronger as molecules get larger and contain more electrons, because bigger electron clouds are more easily distorted. This explains why boiling points rise down a homologous series and down a group of similar molecules: the larger the molecule, the stronger the dispersion forces and the more energy needed to separate the molecules. Van der Waals forces are the only attraction between non-polar molecules such as the halogens and noble gases.

Dipole-dipole forces and hydrogen bonding

Polar molecules attract each other through permanent dipole-dipole forces, where the slightly positive end of one molecule is attracted to the slightly negative end of another. These are stronger than van der Waals forces between molecules of similar size. The strongest intermolecular force is hydrogen bonding, a special case that occurs when hydrogen is bonded directly to nitrogen, oxygen or fluorine, and is attracted to a lone pair on N, O or F in a neighbouring molecule. Hydrogen bonding explains the unusually high boiling point of water, the lower volatility of alcohols compared with alkanes of similar mass, and the fact that ice is less dense than liquid water because of its open hydrogen-bonded structure.

Giant versus simple molecular structures

Covalent substances fall into two structural types with very different properties. Simple molecular substances, such as carbon dioxide, iodine and water, consist of small molecules held to each other by weak intermolecular forces; melting and boiling break only these weak forces, so they have low melting and boiling points and do not conduct electricity. Giant covalent structures, such as diamond, graphite and silicon dioxide, have atoms joined by strong covalent bonds throughout the whole structure, so melting requires breaking many strong bonds and the melting points are very high. The key examination skill is to identify the structure type, then explain the property by naming exactly which forces or bonds must be overcome.

Structure, properties and uses

Linking structure to properties explains why materials are chosen for particular uses. Diamond, with every carbon bonded to four others, is extremely hard and is used in cutting and drilling tools. Graphite, with layers that slide and delocalised electrons, conducts electricity and acts as a lubricant and as electrodes. Simple molecular substances are soft and low-melting, suited to use as gases, volatile liquids or soft solids. Metals, with delocalised electrons and sliding layers, are chosen for wiring and structural use. By reasoning from bonding to structure to property, you can predict whether a substance will be hard or soft, high or low melting, a conductor or an insulator, and soluble or insoluble in water.