A molecule's formula does not determine its shape by itself. Electron pairs repel into a three-dimensional arrangement, and that shape determines whether individual bond dipoles cancel. The result controls intermolecular attractions, volatility, solubility and many real-world uses.

Around a central atom, bonding pairs and lone pairs repel. A double or triple bond counts as one electron domain for shape prediction because its electron density lies in one direction. First identify the total number of domains, arrange them as far apart as possible, then describe the shape made by the atoms only.

Two domains are linear, three trigonal planar, four tetrahedral, five trigonal bipyramidal and six octahedral. Electron-pair repulsion predicts ideal angles, while real angles reflect differences in pair-pair repulsion.

Electron-pair repulsion sets an approximate molecular geometry, while bond electronegativities create individual bond dipoles. Overall molecular polarity depends on the vector sum of those dipoles, so polar bonds can cancel in a symmetric molecule. Intermolecular forces then help explain boiling points, solubility and viscosity, but comparisons should control for molecular size and shape. Hydrogen bonding requires an appropriate H–N, H–O or H–F donor and an acceptor, not merely the presence of hydrogen.

Common molecular shapes must be recalled accurately: $\mathrm{BeC}{\mathrm{l}}_2$ and $\mathrm{C}{\mathrm{O}}_2$ are linear, $\mathrm{BC}{\mathrm{l}}_3$ is trigonal planar, $\mathrm{C}{\mathrm{H}}_4$ is tetrahedral, $\mathrm{PC}{\mathrm{l}}_5$ is trigonal bipyramidal and $\mathrm{S}{\mathrm{F}}_6$ is octahedral. State the angle only when it belongs to the named position in the shape.

For five domains, equatorial positions are $120^{\circ }$ apart and axial-equatorial angles are $90^{\circ }$. For six domains, adjacent bonds are $90^{\circ }$. These spatial distinctions matter when explaining repulsion or polarity.

Repulsion strength follows lone pair–lone $pair>lone$ pair–bond $pair>bond$ pair–bond pair. Lone pairs are held by one nucleus rather than shared between two, so their electron density is more concentrated near the central atom and occupies more space.

Thus methane has the ideal tetrahedral angle of $109.5^{\circ }$, ammonia has about $107^{\circ }$ and water about $104.5^{\circ }$. Do not say that lone pairs attract bonds closer together; they repel bonding pairs more strongly and force the bonding pairs nearer to one another.

A bond is polar when atoms of different electronegativity share electrons unequally. A molecule is polar only if its bond dipoles do not cancel as vectors. Symmetrical $\mathrm{C}{\mathrm{O}}_2$ has two polar $C=O$ bonds but no overall dipole; bent water has polar O–H bonds that reinforce to give a polar molecule.

Judge polarity after finding the shape. Consider identical outer atoms, different outer atoms, lone pairs and any geometry that makes dipoles cancel. This is more reliable than deciding from the formula or from the presence of one electronegative atom.

London forces arise from temporary fluctuations in electron density that induce dipoles in neighbouring particles. They act between all atoms and molecules and grow with electron number, polarizability and surface contact. Branching reduces surface contact, so more-branched isomers usually have lower boiling temperatures.

Permanent dipole–permanent dipole attractions occur between polar molecules. Hydrogen bonding is a particularly strong case when H is bonded to N, O or F and can interact with a lone pair on a neighbouring molecule. Stronger intermolecular forces raise boiling temperature and reduce volatility because more energy is needed to separate molecules.

Molecular polarity is a vector question. A polar bond has an unequal sharing of electrons because the bonded atoms have different electronegativities, but the molecule is polar only if the individual bond dipoles do not cancel. Symmetrical linear, trigonal-planar or tetrahedral molecules can be non-polar even when every bond is polar.

Lone pairs matter twice: they occupy electron domains in VSEPR reasoning and they often produce a net dipole because the molecule becomes asymmetric. This is why CO₂ is non-polar whereas H₂O is polar, and why NH₃ has a net dipole whereas BF₃ does not. The resulting intermolecular forces then influence boiling point, solubility and viscosity.

Comparing safe liquid samples or supplied boiling-point data lets you test predictions based on molecular shape and intermolecular forces. Use a controlled comparison rather than attributing every difference to one variable.

Never decide molecular polarity before you have considered the three-dimensional shape.