Ionic & Covalent Bonding

Why Atoms Bond

Atoms form chemical bonds to lower their energy and reach a stable electron arrangement, usually a full valence shell of eight electrons - the octet rule (two for hydrogen and helium). They achieve this in two main ways: by transferring electrons or by sharing them. The kind of atoms involved decides which path is taken, and that decision drives every physical property the NY Regents exam asks you to predict.

Ionic Bonding: Electron Transfer

An ionic bond forms when a metal transfers one or more electrons to a nonmetal.

• The metal loses electrons and becomes a positively charged cation (for example, Na becomes Na+).
• The nonmetal gains those electrons and becomes a negatively charged anion (Cl becomes Cl-).
• The opposite charges attract through electrostatic force, which is the ionic bond.

In sodium chloride (NaCl), sodium gives its one valence electron to chlorine. Both ions now have full octets. Ionic compounds do not exist as single molecules; they form a repeating three-dimensional crystal lattice, so a formula like NaCl is a ratio, not a discrete molecule. The energy released when gaseous ions assemble into this lattice (lattice energy) is large, which is why ionic solids are so stable and have such high melting points.

The charges of the ions come straight from the periodic table: Group 1 forms +1, Group 2 forms +2, Group 13 forms +3, Group 15 forms -3, Group 16 forms -2, and Group 17 forms -1. Knowing these charges lets you predict ionic formulas without memorizing each compound.

Properties of Ionic Compounds

• High melting and boiling points - strong lattice attractions take a lot of energy to break.
• Crystalline solids at room temperature.
• Conduct electricity when molten or dissolved (the ions are then free to move) but not as a solid.
• Many are soluble in water.

Covalent Bonding: Electron Sharing

A covalent bond forms when two nonmetals share one or more pairs of electrons so each atom reaches a full shell. The shared atoms form a discrete molecule.

• Single bond: one shared pair (H-H in H2).
• Double bond: two shared pairs (O=O in O2).
• Triple bond: three shared pairs (N#N in N2, the strongest and shortest).

A diatomic molecule is two atoms bonded together. Memorize the seven elements that occur naturally as diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2 (a useful phrase is "Have No Fear Of Ice Cold Beer"). The bond order also affects bond length and strength: a triple bond is shorter and stronger than a double bond, which is shorter and stronger than a single bond. This is why nitrogen gas (N2), with its triple bond, is so unreactive and stable.

Network solids such as diamond and silicon dioxide (SiO2) are a special covalent case - one continuous network of covalent bonds, giving extremely high melting points and hardness. They are the exception to the "covalent means low melting point" rule.

Properties of Covalent (Molecular) Substances

• Lower melting and boiling points than ionic compounds.
• Often gases or liquids, or soft solids, at room temperature.
• Poor conductors of electricity (no free-moving charges).
• Many are insoluble in water unless they are polar.

Using Electronegativity to Predict Bond Type

The electronegativity difference between two bonded atoms predicts the bond type. Look up each value on the Properties of Selected Elements table (legacy Table S) and subtract.

Worked example: Is the bond in potassium chloride (KCl) ionic or covalent?

1. From the Properties of Selected Elements table (legacy Table S): K is about 0.8, Cl is about 3.2.
2. Difference = 3.2 - 0.8 = 2.4.
3. 2.4 is greater than 1.7, so the bond is ionic.

This also matches the metal-plus-nonmetal rule: K is a metal, Cl a nonmetal.

Metallic Bonding (Briefly)

In a metal, valence electrons are not held by any single atom; they move freely through a "sea of mobile electrons" surrounding positive metal ions. This metallic bond explains why metals conduct electricity and heat, are malleable (bendable), and are ductile (drawn into wires).

Common Exam Traps

• Calling NaCl a molecule. It is an ionic compound forming a lattice; its formula gives a ratio of ions.
• Predicting conductivity wrong. Solid ionic compounds do not conduct; only molten or aqueous (dissolved) forms do.
• Forgetting the diatomic seven. Writing oxygen as O instead of O2 loses credit.
• Skipping the Properties of Selected Elements table (legacy Table S). Borderline bonds (difference near 1.7) must be checked with actual values, not guessed.
• Mixing up cation and anion. Cations are positive (metals); anions are negative (nonmetals). A memory aid: the t in cation looks like a plus sign.
• Assuming all covalent substances are insoluble. Polar molecules like sugar dissolve well in water.

Knowing the metal-versus-nonmetal pattern, the electron-transfer-versus-sharing mechanism, and the electronegativity cutoff from the Properties of Selected Elements table (legacy Table S) lets you classify any bond and then predict the substance's melting point, state, and conductivity with confidence.