Structure of molecules: Class 9 Chemistry

Key points

Atoms join to become more stable. Most atoms do this by getting a full outer shell like the nearest noble gas. The octet rule says that atoms tend to have 8 electrons in the outer shell. Hydrogen, lithium and beryllium follow the duplet rule and try to reach 2 electrons like helium.

A chemical bond is the force that holds atoms together in a compound. It forms by transfer of electrons, by sharing of electrons, or by the attraction between a sea of electrons and metal ions.

An ionic bond forms by the transfer of electrons from a metal to a non-metal. The metal becomes a positive ion (cation) and the non-metal becomes a negative ion (anion). The strong electrostatic attraction between them is the bond. In sodium chloride, Na gives one electron to Cl to form Na⁺ and Cl⁻. In magnesium chloride, Mg gives two electrons, one to each of two Cl atoms, so the formula is MgCl₂.

Properties of ionic compounds:

• Hard, brittle crystalline solids with high melting and boiling points, because of strong forces in the lattice.
• Soluble in water but usually not in organic solvents.
• Conduct electricity when molten or dissolved in water, because the ions are free to move, but not as solids.

A covalent bond forms when two atoms share one or more pairs of electrons, usually between non-metal atoms. Sharing one pair gives a single bond (H-H in H₂ or Cl-Cl), two pairs a double bond (O=O in O₂, and C=C in ethene) and three pairs a triple bond (N≡N in N₂). In methane, carbon shares four pairs with four hydrogen atoms. In ammonia, nitrogen has three bond pairs and one lone pair. In water, oxygen has two bond pairs and two lone pairs.

Properties of covalent compounds:

• Low melting and boiling points, since the forces between molecules are weak. Many are gases or liquids.
• Usually insoluble in water but soluble in organic solvents (polar ones such as sugar and ethanol dissolve in water).
• Do not conduct electricity, because they have no free ions.

Electronegativity difference decides the nature of a bond. If the atoms are identical, or the difference is zero, the bond is non-polar covalent, as in Cl₂. If the difference is moderate, the shared pair is pulled towards the more electronegative atom, which gets a partial negative charge (δ−) while the other gets δ+. This is a polar covalent bond, as in HCl and H₂O. A very large difference, roughly above 1.7 to 2.0, gives an ionic bond.

A coordinate covalent (dative) bond is a covalent bond in which both shared electrons come from one atom. The ammonium ion NH₄⁺ forms when the lone pair on NH₃ is donated to an H⁺ ion. Once formed, all four N-H bonds are the same.

Metallic bonding is the attraction between positive metal ions and a sea of free (delocalised) valence electrons. It explains why metals conduct electricity and heat, are malleable and ductile, and have high melting points and a shiny surface.

Forces between molecules are much weaker than bonds inside molecules:

• Dipole-dipole forces: attraction between the δ+ end of one polar molecule and the δ− end of another.
• London dispersion forces: weak attractions in all molecules, caused by temporary shifts of electrons. They are the only forces in non-polar substances such as noble gases and Cl₂.
• Hydrogen bonding: a strong dipole attraction between a hydrogen atom bonded to F, O or N and a lone pair on F, O or N of a nearby molecule.

Hydrogen bonding explains why water has a high boiling point (100 °C) compared with hydrogen sulphide, and why ice floats. In ice the molecules are held in an open hexagonal structure that takes up more volume than the liquid, so ice is less dense than water.

Study tip: Practise drawing dot and cross diagrams for NaCl, MgCl₂, H₂, O₂, N₂, H₂O, NH₃ and CH₄ with all outer electrons shown. For comparison questions, use a short table of properties and give a reason (ions or molecules, strong or weak forces) for every difference.

Definitions and flashcards

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Short questions with model answers

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Common mistakes

• Writing that ionic compounds are made of molecules. They are lattices of ions.
• Showing a double bond with only one pair of dots or forgetting lone pairs.
• Confusing breaking a covalent bond with melting or boiling a covalent substance, which only overcomes intermolecular forces.
• Saying hydrogen bonding occurs with any hydrogen atom. It needs F, O or N.
• Forgetting that solid ionic compounds do not conduct.

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