L20. Structure and Properties: Why Salt, Sugar, Water, and Metals Behave Differently

Chemical Names and Formulas

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L20. Structure and Properties: Why Salt, Sugar, Water, and Metals Behave Differently

Why does salt conduct electricity when dissolved but sugar does not, even though both dissolve in water?

From particle pictures to real behavior

You have learned how to draw Lewis models and write formulas. Now use those particle pictures to explain visible behavior. Look closely at common observations: table salt (NaCl) dissolves and its solution conducts electricity; table sugar (C12H22O11) dissolves but its solution does not conduct. The key is what the formula and structure tell us about what moves when the substance changes phase. Ionic compounds like NaCl are made of positive and negative ions held in a crystal lattice. When salt dissolves, the lattice breaks into separate ions that float freely in water. Represent this with a clear equation that conserves atoms and charge: NaCl(s) → Na+(aq) + Cl−(aq) This equation shows one Na atom and one Cl atom on each side, and a total charge of zero on both sides (+1 and −1 sum to 0). Those free ions carry electric current. By contrast, many molecular compounds (like sugar) dissolve as whole neutral molecules; no charged particles are released, so the solution does not conduct electricity.

• Ionic solid (NaCl): rigid lattice of Na+ and Cl− ions → when dissolved, mobile ions → conducts electricity. Molecular solid (sugar): neutral molecules held by intermolecular forces → dissolves as whole molecules → no charge carriers → no conductivity.

Shape and polarity: how direction matters

Not all molecules are the same even when they contain the same atoms. Polarity depends on both how strongly atoms pull electrons (electronegativity) and the molecule’s shape. Two clear examples are carbon dioxide (CO2) and water (H2O). Both contain oxygen and carbon or hydrogen, but their shapes make their overall polarity different. CO2 is linear: O=C=O. Each C–O bond is polar, but the two bond dipoles point in opposite directions and cancel. The molecule has no net dipole moment, so it is nonpolar and doesn’t mix well with water. H2O is bent (about 104.5°). The two O–H bond dipoles do not cancel; they add to give a net dipole. Water is polar and interacts strongly with other polar substances and ions. Quick comparison:

• CO2 — linear, bond dipoles cancel → nonpolar → poor solubility in water. H2O — bent, bond dipoles add → polar → strong solubility with ions and polar molecules.

Metals and networks: other structures, other properties

Some solids are neither simple ionic crystals nor discrete molecular solids. Metals form a lattice of positive metal ions in a “sea” of delocalized electrons. Those electrons move freely in the solid, so metals conduct electricity as solids. Metals are also malleable because layers of ions can slide while the electron sea still holds the metal together. Network covalent solids are another special case. Diamond is a network of carbon atoms covalently bonded in a rigid 3-D framework. This network gives diamond an extremely high melting point and makes it very hard, because breaking it requires breaking many strong covalent bonds. Graphite, also made of carbon, has layers with strong bonds within layers and weak interactions between layers; it conducts along the layers and is slippery. Takeaway: the microscopic arrangement of atoms and electrons — lattice vs molecule vs network — explains macroscopic properties like melting point, conductivity, solubility, and hardness.

Putting structure and properties together

The arrangement of atoms and electrons — whether as ions in a lattice, discrete neutral molecules, a sea of electrons, or a covalent network — directly controls how a substance dissolves, whether it conducts electricity, and how hard or soft it is. Use your Lewis models and formulas to predict these behaviors: ask whether charged particles will form and whether bond directions add or cancel.

Later lessons on the mole and molar mass will let you calculate how many particles are present in samples and how concentration (mol/L) affects properties like conductivity and freezing point. For now, practice linking particle pictures to what you can test in the lab: dissolve, measure conductivity, and compare melting points.