Chemical bonding, structure and properties
Chemical bonding explains how atoms or ions are held together in substances. The type of bonding and the arrangement of particles help explain melting points, electrical conductivity and other properties. These ideas work best when the particles are named accurately. Ionic solids contain ions in extended lattices, rather than separate small molecules of an ionic compound moving independently through the crystal.
An ionic bond is an electrostatic attraction between oppositely charged ions. Sodium can lose an electron to form Na⁺, while chlorine can gain one to form Cl⁻. In solid sodium chloride, each is held within a giant ionic arrangement. The formula NaCl describes the simplest ratio of ions, not a single isolated pair that explains the entire crystal structure.
A covalent bond involves a shared pair of electrons between atoms. Hydrogen, water and methane contain small covalent molecules. Strong covalent bonds hold each together, but the attractions between separate molecules can be much weaker. Melting or boiling a simple molecular substance usually overcomes intermolecular attractions, rather than breaking all its covalent bonds. This distinction explains many comparatively low melting and boiling points.
Some covalent substances form giant networks instead of small molecules. Diamond contains a three-dimensional carbon network, and graphite has carbon layers with electrons that can move within them. Graphite therefore conducts electricity, while diamond normally does not. Describing both merely as ‘covalent’ misses the structural features that account for their different properties and common uses.
In metallic bonding, positive metal centres are attracted to delocalised electrons. Mobile electrons allow metals to conduct electrical . Layers can move while metallic bonding continues, helping explain malleability. Different metals and alloys have different strengths and melting points, so the model gives a general explanation rather than one fixed numerical property for all metals.
Chemical formulas must be electrically balanced when describing neutral ionic compounds. Magnesium forms Mg²⁺ and chloride is Cl⁻, so magnesium chloride is MgCl₂. Aluminium oxide combines Al³⁺ and O²⁻ in the ratio 2:3 to give Al₂O₃. Subscripts count the simplest whole-number ratio; they do not label the charges. In an exam explanation, connect a stated property to both bonding and particle mobility or structure.
From electron patterns to stable arrangements
Bond formation is associated with a lower- overall arrangement under the relevant conditions, not atoms consciously wanting a particular electron count. The octet rule is a useful guide for many main-group compounds but has exceptions, including hydrogen's two-electron shell and electron-deficient or expanded-valence examples. Draw electron-dot diagrams with a key showing which supplied each electron where required. Distinguish a bonding pair from a lone pair. Ionic attraction extends through a lattice, whereas a simple molecular covalent substance contains identifiable molecules. The formula alone is not enough to determine every property without knowing the arrangement of particles.
Shapes and lone-pair effects
A simple electron-pair repulsion model places electron regions around a central as far apart as possible. Four bonding regions give a tetrahedral arrangement, as in methane, with bond angles about 109.5°. Ammonia has three bonds and one lone pair, producing a trigonal-pyramidal molecular shape with an angle near 107°. Water has two bonds and two lone pairs and is bent, near 104.5°. Carbon dioxide has two electron regions around carbon and is linear, even though each region is a double bond. These are approximate model results. Molecular shape describes the positions, so do not name water tetrahedral merely because its four electron regions have an approximately tetrahedral arrangement.
Polar bonds and molecular polarity
A covalent bond can be polar when the bonded atoms attract shared electrons unequally. The more electronegative carries partial negative character and the other partial positive character. These partial charges are not necessarily complete ionic charges. Molecular polarity depends on both bond polarities and shape. Carbon dioxide has polar bonds but a symmetrical linear arrangement whose bond dipoles cancel. Water's bent shape prevents equivalent cancellation, so it has a net dipole. A containing polar bonds is therefore not automatically polar as a whole. Use a diagram to show vector cancellation rather than treating the word polar as an isolated label.
Intermolecular attractions and physical changes
London dispersion attractions occur between all atoms and molecules and generally grow with polarisability and contact opportunities. Permanent-dipole attractions occur between suitable polar molecules. Hydrogen bonding is a particularly important interaction when hydrogen is bonded to a highly electronegative such as nitrogen, oxygen or fluorine and interacts with a suitable lone pair. These attractions help explain water's boiling behaviour and differences among alcohols and hydrocarbons. Boiling a molecular liquid separates molecules against intermolecular attractions; it does not normally break all the covalent bonds inside each . Compare similar-sized molecules carefully before attributing every boiling-point difference to a single factor.
Coordinate bonding and familiar ions
In a coordinate or dative covalent bond, both electrons in the newly formed shared pair are supplied by one donor in the bonding description. Ammonia can donate its nitrogen lone pair to H⁺ to form NH₄⁺. Once formed, the ammonium 's N–H bonds are equivalent in the simple tetrahedral rather than one permanently weaker special bond. Coordinate descriptions explain how some particles form, but the resulting bond remains a covalent interaction. Keep the overall charge outside brackets in an electron diagram. A positive polyatomic does not that each has the same positive charge or has lost a proton.
Structure-to-property explanations
An ionic lattice usually has a high melting point because many strong electrostatic attractions must be overcome. A simple molecular solid may melt at a much lower temperature because weaker intermolecular attractions are overcome instead. Giant covalent structures can be very hard and high melting, but structure matters: graphite layers can slide and its delocalised electrons allow conduction along layers. Metals conduct through mobile electrons even when solid; ionic substances require mobile ions, as in a melt or suitable solution. An ionic compound does not automatically dissolve well in water. Solubility reflects competing interactions, and some ionic salts are sparingly soluble.
Chemical formulas, valency and charge balance
Use known charges to find the simplest neutral ionic ratio. Calcium nitrate needs one Ca²⁺ and two NO₃⁻ ions, giving Ca(NO₃)₂. The brackets show two whole nitrate groups, so there are two nitrogen and six oxygen atoms per formula unit. Aluminium sulfate is Al₂(SO₄)₃, combining total charges +6 and −6. Do not write charges as neutral-formula subscripts without simplifying a common ratio; Mg₂O₂ reduces to MgO for the formula-unit ratio. state is a bookkeeping concept and is not always identical to an 's actual charge in a covalent substance.
Self-check with reasoned answers
Explain why copper conducts when solid but solid sodium chloride does not. Copper has mobile delocalised electrons; sodium chloride's ions are held in lattice positions and cannot carry a bulk conduction through the crystal under the ordinary model. Explain why methane has a much lower boiling point than a giant covalent solid: separating methane molecules overcomes intermolecular attractions, while disrupting the giant network involves extensive covalent bonding. Finally, predict the formula of a compound from Al³⁺ and O²⁻ by finding equal total charges: two aluminium ions and three oxide ions give Al₂O₃, not Al₃O₂. Name the moving particle or the interaction whenever a question asks why.