Chemical Bonding Basics

Ionic compounds form when electrons transfer from metal to nonmetal, producing cations and anions held by electrostatic attraction. Sodium chloride NaCl is a 1:1 ratio of Na⁺ and Cl⁻; molar mass sums one Na and one Cl. Calcium carbonate CaCO₃ contains Ca²⁺ and CO₃²⁻ — the carbonate polyatomic ion stays intact in the formula. Molar mass calculations count every atom in the written formula, whether molecular or ionic.

Nonmetals share electrons to reach stable configurations. Water H₂O has two O–H covalent bonds; the molecule is discrete with molar mass 18.02 g/mol. Methane CH₄, ethanol C₂H₆O, and benzene C₆H₆ are likewise molecular — each formula represents one molecule whose mass you sum from atomic contributions. Compare water with hydrogen bonding versus methane CH₄, a nonpolar gas: same counting rules, different physical properties.

Sulfuric acid H₂SO₄ is covalent in the pure liquid but ionizes in water to H⁺ and SO₄²⁻. Ammonium chloride NH₄Cl contains the covalently bonded NH₄⁺ ion paired with Cl⁻. Polyatomic ions (SO₄²⁻, NO₃⁻, PO₄³⁻) must be treated as grouped units when parsing parentheses: aluminum sulfate Al₂(SO₄)₃ has three sulfate groups, not three isolated sulfurs and oxygens counted incorrectly.

Metals and network solids (diamond, quartz) do not have simple molecular formulas in the same sense as water. Classroom molar mass still uses the formula unit shown on the periodic table or in the problem (e.g., Mg for magnesium metal). Copper sulfate CuSO₄ for lab reagent calculations uses the written formula including hydrate water when applicable. Bonding type informs physical state and solubility; molar mass arithmetic stays the same once the correct formula is identified.

Knowing whether two elements are likely to form an ionic or covalent bond helps you predict a compound's formula before you even calculate its molar mass, which is a useful sanity check. Metals paired with nonmetals (like sodium and chlorine, or calcium and oxygen) almost always form ionic compounds, and the resulting formula follows a simple charge-balancing rule: the total positive charge from cations must equal the total negative charge from anions. Calcium (Ca²⁺) paired with chloride (Cl⁻) requires two chlorides per calcium to balance charge, giving CaCl₂ — not CaCl or CaCl₃.

Two nonmetals bonding together (like carbon and oxygen, or nitrogen and hydrogen) typically form covalent molecules, where the specific ratio of atoms depends on each element's number of available bonding electrons rather than a simple charge-balance rule. This is part of why carbon so often forms multiple different oxides (CO, CO₂) and multiple different compounds with hydrogen (CH₄, C₂H₆, C₃H₈, and so on) — covalent bonding allows more structural flexibility than the fixed charge-balancing rules of ionic bonding.

Aluminum forms Al³⁺ ions, and sulfate is the polyatomic ion SO₄²⁻. To balance charge, you need the smallest whole numbers of each ion whose total charges cancel: using the cross-multiplication trick (aluminum's charge becomes sulfate's subscript, and vice versa), you get Al₂(SO₄)₃ — 2 aluminum ions (total charge +6) balanced against 3 sulfate ions (total charge −6). This is exactly why aluminum sulfate's formula has those specific subscripts, and recognizing this charge-balance logic is a quick way to verify you've copied or remembered a polyatomic ion formula correctly before starting a molar mass calculation.

A frequent mistake is writing an ionic formula with the wrong subscripts because the charge-balancing step was skipped or done incorrectly — for instance, writing "CaCl" instead of the correct "CaCl₂" for calcium chloride, which would give an incorrect molar mass of about 75.5 g/mol instead of the correct 110.98 g/mol. Another mistake is treating a polyatomic ion's internal atoms as separable when applying an outer subscript, discussed at length in the "How to Calculate Molar Mass" and "Common Molar Mass Mistakes" guides — this error is really a bonding-conceptual mistake as much as an arithmetic one, since it comes from not recognizing that polyatomic ions act as a single bonded, indivisible unit within the larger formula.