Regents Chemistry chemical bonding and molecular structure: a complete skills guide to ionic, covalent and metallic bonding, polarity and properties

Why bonding ties the whole course together

Chemical bonding is where atomic structure turns into the properties of real substances. The valence electrons you counted in the first unit decide which bonds form, electronegativity decides how polar those bonds are, shape decides whether a molecule is polar, and intermolecular forces decide the melting and boiling points. This guide pulls together the matching dot-point pages, each with Regents-format practice: types of chemical bonds, electronegativity and bond polarity, Lewis structures and molecular polarity, intermolecular forces, and properties of ionic and molecular substances.

The three types of bond

Atoms bond to reach a lower-energy, more stable arrangement, usually a full outer level. An ionic bond is the electrostatic attraction between ions formed when a metal transfers electrons to a nonmetal. A covalent bond forms when two nonmetals share electrons (single, double or triple). A metallic bond is the attraction between positive metal ions and a sea of mobile valence electrons. Forming a bond releases energy; breaking one absorbs energy.

Electronegativity and bond polarity

Subtract the two atoms' electronegativities (Table S) to classify the bond. A difference of about zero is nonpolar covalent; a small to moderate difference is polar covalent, with a partial negative charge on the more electronegative atom; a large difference is ionic. For ranking, the largest difference is the most polar bond. Identical atoms always give a nonpolar covalent bond.

Lewis diagrams and molecular polarity

A Lewis electron-dot diagram shows valence electrons as shared (bonding) pairs and lone (nonbonding) pairs, with most atoms reaching an octet and hydrogen reaching two. A molecule's overall polarity needs two things: polar bonds and an asymmetrical shape. Symmetrical molecules ($\text{CO}_2$, linear) cancel their bond dipoles and are nonpolar; bent ($\text{H}_2\text{O}$) and pyramidal ($\text{NH}_3$) molecules do not cancel and are polar.

Intermolecular forces and boiling point

Forces between molecules are weaker than the bonds inside them, but they set melting and boiling points. From strongest to weakest: hydrogen bonding (H bonded to N, O or F), dipole-dipole (polar molecules), and dispersion (all molecules, and the only force for nonpolar molecules). Stronger forces mean higher boiling points. Hydrogen bonding explains water's high boiling point, its surface tension, and why ice floats.

Properties from bonding

The recurring Regents move is to read off a substance's properties and name the bonding: a high-melting solid that conducts only when molten or dissolved is ionic; a low-melting nonconductor is molecular; a malleable conductor is metallic.

Check your knowledge

Attempt these under timed conditions, then check the solutions.

1. State the type of bonding in (a) $\text{CaCl}_2$, (b) $\text{O}_2$, (c) iron metal. (3 marks)
2. Using electronegativities (Na $0.9$, Cl $3.2$, H $2.2$), classify the bonds in $\text{NaCl}$ and $\text{HCl}$. (2 marks)
3. Describe the Lewis electron-dot diagram of a water molecule, including lone pairs. (2 marks)
4. Explain why $\text{CO}_2$ is nonpolar even though it has polar bonds. (1 mark)
5. Rank dispersion forces, hydrogen bonding and dipole-dipole forces from weakest to strongest. (1 mark)
6. Explain why solid sodium chloride does not conduct electricity but its solution does. (2 marks)