This lesson explains the "glue" that holds atoms together to form molecules and compounds. The goal for atoms is to achieve a stable electron configuration, like the noble gases (e.g., Neon: 2,8; Argon: 2,8,8), which usually means having a full outer shell of 8 electrons (the "octet rule").

1. Short Notes: Core Concepts

There are two main ways atoms achieve this stability: by transferring electrons or by sharing them.

A. Ionic Bonds (Transferring Electrons)

• Who: Happens between a metal and a non-metal.

• What happens: The metal atom loses one or more electrons to become a positively charged ion (cation). The non-metal atom gains those electrons to become a negatively charged ion (anion).

• The Bond: The bond is the strong electrostatic attraction between the positive cation and the negative anion.

• Example: Sodium Chloride (NaCl)

  • Sodium (Na: 2,8,1) wants to lose 1 electron to become stable (Na⁺: 2,8).

  • Chlorine (Cl: 2,8,7) wants to gain 1 electron to become stable (Cl⁻: 2,8,8).

  • Na gives its electron to Cl. The resulting Na⁺ and Cl⁻ ions are strongly attracted to each other.

• Structure: Ionic compounds form giant, regular, repeating structures called ionic lattices.

B. Covalent Bonds (Sharing Electrons)

• Who: Happens between non-metal atoms.

• What happens: Instead of transferring, the atoms share pairs of electrons so that each atom feels like it has a full outer shell.

• The Bond: The bond is the shared pair of electrons.

  • Single Bond: One pair of shared electrons (e.g., in H₂, H₂O, CH₄).

  • Double Bond: Two pairs of shared electrons (e.g., in O₂).

  • Triple Bond: Three pairs of shared electrons (e.g., in N₂).

• Structure: Covalent substances usually exist as individual, separate units called molecules. (Exceptions like diamond and graphite form giant atomic lattices).

C. Polarity and Intermolecular Bonds

• Electronegativity: The ability of an atom to attract the shared electrons in a covalent bond.

• Non-polar Covalent Bond: If two identical atoms share electrons (e.g., H-H or Cl-Cl), they share them equally. There is no charge difference across the molecule.

• Polar Covalent Bond: If two different atoms share electrons (e.g., H-Cl or H-O), the atom with higher electronegativity pulls the electrons closer, creating a slight negative charge (δ⁻) on that atom and a slight positive charge (δ⁺) on the other.

• Intermolecular Bonds: The slight positive end (δ⁺) of one polar molecule attracts the slight negative end (δ⁻) of a neighboring molecule. These are weak attractions between molecules (not to be confused with the strong covalent bonds within a molecule). These forces explain why water (H₂O) has a much higher boiling point than expected.

2. Tips & Tricks for the Exam

• Metal + Non-metal → Ionic Bond. (Look for an element from the left side of the periodic table with one from the right side).

• Non-metal + Non-metal → Covalent Bond. (Look for elements from the right side of the periodic table).

• Drawing Diagrams:

  • For ionic bonds, draw the atoms before and after the electron transfer. Show the electron moving with an arrow. Make sure to write the final charges on the ions (e.g., Na⁺, Cl⁻).

  • For covalent bonds, use "dot and cross" diagrams (Lewis structures). Draw the outer shells overlapping and show the shared pairs in the overlapping section.

3. Important Points & Common Exam Questions

• Properties of Ionic vs. Covalent Compounds: This is a very common comparison question.

• Why do ionic compounds conduct electricity only when molten or dissolved?

  • Answer: In the solid state, the ions are held in fixed positions in the lattice and cannot move. When molten or dissolved, the ions are free to move and carry an electrical charge.

• Explain the formation of a [given] compound. You should be able to draw a dot and cross diagram for simple compounds like MgO (ionic) or NH₃ (covalent).