General Chemistry: Bonding and Stoichiometry

Bonding and stoichiometry tie the whole Chem/Phys section together: periodic trends predict how atoms bond, bonding predicts molecular shape and intermolecular forces, and the mole concept converts all of it into measurable mass and moles. Master the logic chain below and a large block of discrete and passage questions becomes predictable.

Core Idea

• Periodic trends flow from effective nuclear charge and shell number. Moving right increases pull on electrons; moving down adds shells that shield and enlarge the atom.
• Bond type is a spectrum set by electronegativity difference. Large difference gives ionic, moderate gives polar covalent, near-zero gives nonpolar covalent, and shared delocalized electrons give metallic.
• The mole is a counting bridge. Avogadro's number links the atomic-scale count of particles to gram-scale mass, making balanced equations quantitative.

Atomic Structure and Periodic Trends

The periodic table organizes elements by increasing atomic number and recurring valence-shell configurations. Three trends dominate the MCAT:

• Atomic radius decreases left-to-right (rising effective nuclear charge pulls electrons in) and increases top-to-bottom (added shells).
• Ionization energy — the energy to remove an electron — increases left-to-right and decreases down. It is the inverse of atomic radius.
• Electronegativity — pull on shared bonding electrons — follows ionization energy: it rises toward fluorine (upper right, excluding noble gases).

Cations are smaller than their parent atoms; anions are larger.

Chemical Bonding

Bonding is classified by how valence electrons are handled:

• Ionic bonds form between a metal and nonmetal (large electronegativity difference, roughly greater than 1.7): electrons transfer, creating cations and anions held by electrostatic attraction.
• Covalent bonds share electrons between nonmetals. Nonpolar covalent bonds share electrons equally (difference near zero, as in diatomic elements); polar covalent bonds share unequally, creating partial charges and a dipole moment pointing toward the more electronegative atom.
• Metallic bonds feature a lattice of cations in a delocalized "sea" of electrons, explaining conductivity and malleability.

A molecule can have polar bonds yet be nonpolar overall if symmetry cancels the dipoles (CO2 is linear and nonpolar; H2O is bent and polar).

Lewis Structures, Formal Charge, and VSEPR

Draw Lewis structures by counting total valence electrons, connecting atoms, and completing octets (H wants 2). Formal charge = valence electrons − nonbonding electrons − half the bonding electrons; the best structure minimizes formal charges and places negative formal charge on the most electronegative atom.

VSEPR predicts geometry because electron pairs repel to maximize separation. Count electron domains (bonds + lone pairs) around the central atom:

• 2 domains: linear (180°)
• 3 domains: trigonal planar (120°); one lone pair gives bent
• 4 domains: tetrahedral (109.5°); one lone pair gives trigonal pyramidal, two give bent

Lone pairs repel more strongly, compressing bond angles slightly.

Intermolecular Forces

IMFs govern boiling point, solubility, and viscosity, ranked strongest to weakest:

• Hydrogen bonding: H bonded to N, O, or F attracted to a lone pair on another N, O, or F — the strongest IMF and the reason water has a high boiling point.
• Dipole-dipole: attraction between permanent dipoles of polar molecules.
• London dispersion forces: temporary induced dipoles present in all molecules; they strengthen with molar mass and surface area, and dominate in large nonpolar molecules.

The Mole Concept and Stoichiometry

Avogadro's number (6.022 × 10^23) defines one mole of particles. Molar mass (g/mol) numerically equals the atomic or formula mass in amu, converting grams to moles. To do stoichiometry: balance the equation, convert given mass to moles, apply the mole ratio from coefficients, then convert to the target quantity. The limiting reagent is fully consumed first and caps product formed; the other reactant is in excess. Percent yield = (actual yield ÷ theoretical yield) × 100.

High-Yield Exam Patterns

• Rank atomic radius, ionization energy, or electronegativity by position — expect a two-element or two-ion comparison.
• Given a molecule, predict polarity by combining bond dipoles with molecular geometry (symmetry cancellation is the classic trap).
• Identify molecular shape and bond angle from the number of electron domains, adjusting for lone pairs.
• Rank boiling points using IMF strength; hydrogen bonding almost always wins.
• Compute a limiting reagent by converting both reactants to moles of product and taking the smaller value.
• Use percent yield to back-calculate actual mass from a theoretical amount.

Common Traps to Avoid

• Assuming a molecule with polar bonds is automatically polar — symmetric geometry (CO2, CCl4) cancels the net dipole.
• Confusing ionization energy and electronegativity direction with atomic radius; radius runs opposite to the other two.
• Picking the reactant present in the smaller mass as limiting instead of converting to moles first.
• Forgetting to balance the equation before applying mole ratios.
• Treating London dispersion as negligible; in large nonpolar molecules it can exceed a small molecule's dipole forces.