Chemical compounds come in two main flavors: ionic and molecular. Ionic compounds are like salt, with charged particles held together by strong electrical forces. Molecular compounds, like water, have atoms sharing electrons in covalent bonds.

Knowing the periodic table helps predict compound types. Metals on the left tend to form ionic bonds with nonmetals on the right. Nonmetals usually form molecular compounds with each other. Understanding these patterns is key to predicting chemical behavior.

Ionic and Molecular Compounds

Ionic vs molecular compounds

• Ionic compounds consist of positively charged cations (metals) and negatively charged anions (nonmetals) held together by strong electrostatic forces called ionic bonds (NaCl)
  • Form when metals transfer electrons to nonmetals resulting in high melting and boiling points
  • Conduct electricity when dissolved in water or molten due to the presence of mobile ions (KCl solution)
• Molecular compounds composed of atoms held together by covalent bonds formed by the sharing of electrons between nonmetals (H₂O)
  • Held together by relatively weaker intermolecular forces such as hydrogen bonding or van der Waals forces
  • Generally have lower melting and boiling points compared to ionic compounds (CO₂ sublimes at -78.5 ℃)
  • Usually do not conduct electricity in any state because they do not contain mobile ions (CH₄)

Compound types from periodic trends

• Ionic compounds typically form between metals on the left side of the periodic table and nonmetals on the right side (NaBr)
  • Elements with large differences in electronegativity are more likely to form ionic bonds (LiF)
• Molecular compounds typically form between nonmetals on the right side of the periodic table (Cl₂)
  • Elements with similar or slightly different electronegativity values tend to share electrons and form covalent bonds (HCl)
• Electronegativity trends in the periodic table play a key role in predicting bond types
  • Increases from left to right across a period as the atomic radius decreases and effective nuclear charge increases (C < N < O < F)
  • Decreases from top to bottom within a group as the atomic radius increases and effective nuclear charge decreases (F > Cl > Br > I)

Formulas of ionic compounds

1. Determine the charges of the cation and anion

   • Metal cations usually have positive charges equal to their group number (K⁺ from Group 1, Ca²⁺ from Group 2)
   • Nonmetal anions usually have negative charges equal to (8 - group number) (S²⁻ from Group 16, P³⁻ from Group 15)

2. Balance the charges of the cation and anion to achieve a neutral compound

   • The sum of the charges in the formula must equal zero
   • Use subscripts to indicate the number of each ion needed for charge balance (Al³⁺ and O²⁻ → Al₂O₃)

• Examples showcase the application of these rules
  • Potassium bromide: K⁺ and Br⁻ → KBr
  • Calcium fluoride: Ca²⁺ and F⁻ → CaF₂
  • Aluminum sulfide: Al³⁺ and S²⁻ → Al₂S₃

Electronic Structure and Bonding

• The octet rule guides the formation of stable compounds, where atoms tend to gain, lose, or share electrons to achieve a full outer shell of 8 electrons
• Valence electrons, the outermost electrons of an atom, play a crucial role in chemical bonding and reactivity
• Lewis structures provide a visual representation of valence electrons and bonding in molecules
• Polarity in molecules arises from uneven distribution of electron density
  • Dipole moments quantify the degree of charge separation in polar molecules