From electrons to shapes
Lewis structures track valence electrons as bonds and lone pairs. Formal charge helps compare plausible arrangements, and resonance describes a single electron distribution that cannot be captured by only one drawing. Electron groups repel and establish approximate geometry; lone pairs can compress bond angles. A molecule’s shape and bond polarities together determine whether it has a net dipole.
A Lewis structure begins with a total valence-electron count, adjusted for ionic charge. Bonds usually count as shared electron pairs; unshared pairs remain on individual atoms in the drawing. The octet rule is a useful guide for many second-period atoms, especially carbon, nitrogen, oxygen, and fluorine. It is not a universal law: hydrogen seeks a duet, electron-deficient boron compounds exist, and heavier central atoms require careful treatment.
Resonance structures are alternative drawings of the same nuclear arrangement, not molecules rapidly switching between different positions of atoms. In nitrate, equivalent N–O connections are better represented by a delocalized distribution than by one permanent double bond and two permanently different single bonds. Formal charge helps assess drawings, but the smallest charges do not automatically settle every case; electronegativity and experimentally observed structure matter too.
Several kinds of solid
Ionic solids contain extended arrays of oppositely charged ions; metals have mobile electrons; network-covalent solids contain continuous bonding; molecular solids consist of discrete molecules held together by intermolecular attractions. These structural differences explain why melting points, brittleness, and electrical conductivity differ. A formula like NaCl names a ratio in a crystal, not a separate NaCl molecule.
An ionic bond description emphasizes electron transfer and electrostatic attraction, while a covalent one emphasizes electron sharing. Real bonds can lie between those simplified extremes. Electronegativity differences help estimate polarity, but a cutoff value is a convention rather than a physical switch. A diatomic molecule's bond dipole points toward the atom drawing electron density more strongly; in a larger molecule the geometry determines whether those individual dipoles combine or cancel.
VSEPR works by arranging electron-density regions around a central atom to reduce repulsion. Count multiple bonds as one region for the basic shape prediction, and distinguish electron-domain geometry from the shape traced only by atoms. Water has a roughly tetrahedral electron-domain arrangement yet a bent molecular shape because two positions are occupied by lone pairs. The model predicts broad geometry; measured angles can depart from ideal values.
How bonds explain properties
Bond order and bond length reflect how electron density is shared. Greater bond order often means a shorter, stronger bond, but comparing different elements requires more care. A substance’s bulk behavior also depends on attractions between its particles and on its structure, not merely on the strength of one bond.
The energy required to separate a lattice into gaseous ions is tied to charge and distance. Ions with greater charge magnitude and smaller size often create stronger attractions, though crystal structure also matters. This helps explain why many ionic materials have high melting temperatures and why their solid forms do not conduct well: ions are fixed in the lattice. Melting or dissolving may allow ions to move and carry current.
For molecular compounds, the important question at melting can be completely different. Ice melts without decomposing every water molecule; the network of attractions between molecules reorganizes. Diamond is a covalent network, not a pile of small carbon molecules, and cutting all its bonds requires far more energy. Comparing these cases makes the “what particles are being separated?” question a reliable way to reason about properties.
A model has a job
Lewis drawings, VSEPR, and orbital models answer different questions. A Lewis structure counts valence electrons; VSEPR predicts a rough arrangement around a central atom; orbital ideas explain how electron density can support that arrangement. None is a photograph of a molecule. Expanded-octet drawings and resonance should be treated as representations with limits rather than literal alternating bonds.
A Lewis drawing is most useful for counting and connectivity, not for picturing the full three-dimensional electron cloud. VSEPR then uses those electron regions to predict shape, while bond polarity helps predict how the molecule interacts with an electric field or solvent. A model can be useful even when it simplifies the physics. The discipline is to ask what observation it is meant to explain and where it is likely to fail.
For example, a single Lewis structure for benzene would suggest alternating bond lengths. Experiment finds its ring bonds much more alike, which resonance and delocalized electron models represent better. That does not make the Lewis drawing worthless: it still communicates atoms, valence, and possible electron arrangements. It does mean that the drawing should not be mistaken for a frozen microscopic photograph.
From one bond to a whole material
An ionic lattice is stabilized by many attractions, which helps explain its high melting point. In a metal, mobile charge carriers account for conductivity even while atoms remain in a solid structure. A molecular substance may have strong covalent bonds within molecules yet melt easily because the attractions between molecules are weaker. Always name which interaction you are comparing.
In a metallic solid, atoms occupy a lattice and electrons can respond collectively to an applied field. This is why metals generally conduct as solids, unlike typical ionic crystals. Their layers can often slide without immediately destroying all bonding, helping explain malleability. A covalent network such as silicon has a different electronic structure and can become a semiconductor whose conductivity is altered by impurities and temperature.
For a molecular solid, size, shape, and intermolecular forces can dominate the melting behavior. Larger electron clouds often increase dispersion attraction, but packing and polarity also matter. The ability of a substance to dissolve likewise depends on the attractions formed with the solvent and those disrupted in both solute and solvent. No single label—“ionic” or “covalent”—predicts every property without structure and conditions.