2.6 Ionic and Molecular Compounds

Ionic and Molecular Compounds

Ionic vs molecular compounds

Chemical compounds fall into two main categories based on how their atoms are held together.

Ionic compounds consist of positively charged cations (metals) and negatively charged anions (nonmetals) held together by strong electrostatic forces called ionic bonds. Table salt (NaCl) is a classic example.

• They form when metals transfer electrons to nonmetals, resulting in high melting and boiling points
• They conduct electricity when dissolved in water or melted, because the ions become free to move (for example, a KCl solution conducts current)

Molecular compounds are composed of atoms held together by covalent bonds, where nonmetals share electrons rather than transferring them. Water ($H_2O$) is a familiar example.

• They're held together by relatively weaker intermolecular forces like hydrogen bonding or van der Waals forces
• They generally have lower melting and boiling points compared to ionic compounds ($CO_2$ sublimes at $-78.5°C$)
• They usually don't conduct electricity in any state because they contain no mobile ions ($CH_4$, for instance, is a non-conductor)

Compound types from periodic trends

The periodic table gives you a reliable way to predict what type of compound two elements will form. The key factor is electronegativity, which measures how strongly an atom attracts electrons in a bond.

• Ionic compounds typically form between metals (left side of the table) and nonmetals (right side). A large difference in electronegativity between two elements makes ionic bonding more likely. LiF is a good example: lithium has very low electronegativity and fluorine has the highest.
• Molecular compounds typically form between two nonmetals. When elements have similar electronegativity values, they tend to share electrons rather than transfer them. $Cl_2$ (two identical atoms) and HCl (two nonmetals with a modest electronegativity difference) are both molecular.

Electronegativity follows predictable trends on the periodic table:

• Increases left to right across a period, because atomic radius decreases and effective nuclear charge increases ($C < N < O < F$)
• Decreases top to bottom within a group, because atomic radius increases and electrons are farther from the nucleus ($F > Cl > Br > I$)

Formulas of ionic compounds

Writing the formula for an ionic compound comes down to balancing charges so the compound is electrically neutral. Here's the process:

1. Determine the charge of each ion.

   • Metal cations usually carry a positive charge equal to their group number. For example, $K^+$ from Group 1, $Ca^{2+}$ from Group 2.
   • Nonmetal anions usually carry a negative charge equal to $(8 - \text{group number})$. For example, $S^{2-}$ from Group 16, $P^{3-}$ from Group 15.

2. Balance the charges so they sum to zero. Use subscripts to show how many of each ion you need.

A few examples to see this in action:

• Potassium bromide: $K^+$ and $Br^-$ already balance one-to-one → KBr
• Calcium fluoride: $Ca^{2+}$ needs two $F^-$ ions to balance → $CaF_2$
• Aluminum oxide: $Al^{3+}$ and $O^{2-}$ need a 2:3 ratio (since $2 \times 3 = 3 \times 2 = 6$) → $Al_2O_3$
• Aluminum sulfide: $Al^{3+}$ and $S^{2-}$ follow the same cross-over logic → $Al_2S_3$

Electronic Structure and Bonding

The octet rule is the guiding principle behind most bonding: atoms tend to gain, lose, or share electrons until they have 8 electrons in their outer shell (a full valence shell). This drive toward a stable electron configuration is why ionic and covalent bonds form in the first place.

• Valence electrons are the outermost electrons of an atom and are the ones directly involved in bonding. The number of valence electrons an element has determines how it bonds and how many bonds it can form.
• Lewis structures give you a visual way to represent valence electrons and bonding in molecules. Dots represent lone (unbonded) electrons, and lines represent shared pairs.
• Polarity arises when electron density is unevenly distributed in a molecule. If one atom pulls on the shared electrons more strongly than the other (due to higher electronegativity), the bond is polar. Dipole moments quantify this charge separation: the larger the dipole moment, the more polar the molecule.