Chemical Bonds

• Coulomb's Law: The strength of the bond is governed by the equation $F = k \frac{q_1 q_2}{r^2}$. This implies that bonds are stronger when the charges ($q$) are higher and the distance ($r$) between the nuclei is smaller.

• Electronegativity ($\chi$): This is a measure of an atom's ability to attract shared electrons. The difference in electronegativity ($\Delta\chi$) between two atoms determines the bond's character, ranging from purely covalent to ionic.

• Potential Energy Wells: As two atoms approach, attractive forces lower the potential energy until a minimum is reached at the equilibrium bond length. If they move closer, nuclear repulsion causes the potential energy to spike sharply.

Predicting Bond Type

• Step 1: Identify the elements: Generally, metal + non-metal forms ionic bonds, while non-metal + non-metal forms covalent bonds.
• Step 2: Calculate $\Delta\chi$: Use the Pauling scale. If $\Delta\chi > 1.7$, the bond is considered mostly ionic. If $0.4 < \Delta\chi < 1.7$, it is polar covalent. If $\Delta\chi < 0.4$, it is non-polar covalent.

Drawing Lewis Structures

• Sum Valence Electrons: Count the total valence electrons for all atoms in the molecule, adjusting for charge if it is an ion.
• Arrange Atoms: Place the least electronegative atom in the center (excluding Hydrogen) and connect others with single bonds.
• Distribute Remaining Electrons: Complete the octets of outer atoms first, then place any remaining electrons on the central atom. If the central atom lacks an octet, form multiple bonds.

• Check for Exceptions: Always remember that Beryllium and Boron often form compounds with fewer than eight electrons (incomplete octets), while elements in Period 3 and below can have expanded octets (more than eight electrons).

• Formal Charge Calculation: Use the formula $FC = V - L - \frac{1}{2}S$ (where $V$ is valence electrons, $L$ is lone pair electrons, and $S$ is shared electrons) to determine the most plausible Lewis structure. The structure with formal charges closest to zero is usually the most stable.

• Bond Length vs. Strength: Remember the inverse relationship. Triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds. In exams, use this to rank molecules by stability or reactivity.

• The 'Metal-Nonmetal' Trap: Do not assume a bond is ionic just because it involves a metal and a non-metal. For example, $AlCl_3$ has significant covalent character; always check the electronegativity difference if data is provided.

• Confusing Polarity: A molecule can have polar bonds but be a non-polar molecule overall if the geometry is symmetrical (e.g., $CO_2$). Always consider the VSEPR shape before concluding molecular polarity.

• Bond Breaking Energy: Students often mistakenly think bond breaking releases energy. Breaking a bond always requires an input of energy (endothermic), while forming a bond releases energy (exothermic).