Chemical Formulas

Why This Matters

Chemical formulas are the language of chemistry—they tell you not just what atoms are present, but how many, in what ratio, and how they're arranged. In Honors Chemistry, you're being tested on your ability to move fluidly between different formula types, calculate quantities from formulas, and use formulas to predict compound behavior. These skills form the foundation for stoichiometry, reaction prediction, and understanding molecular properties.

The concepts here connect to nearly every major unit you'll encounter: mole calculations, bonding theory, reaction balancing, and redox chemistry. When you see a formula on an exam, you should immediately recognize what type it is, what information it provides, and what calculations you can perform with it. Don't just memorize definitions—know what each formula type reveals about a compound and when to use each one.

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Types of Chemical Formulas

Different formula types communicate different information about a compound. The key is knowing what level of detail each provides and when each is most useful.

Empirical Formula

• Simplest whole-number ratio of atoms in a compound—derived from experimental data like combustion analysis
• Does not show actual atom counts; glucose ($C_6H_{12}O_6$) has the empirical formula $CH_2O$
• Starting point for determining molecular formulas when combined with molar mass data

Molecular Formula

• Shows the actual number of each atom in one molecule of a covalent compound
• Always a whole-number multiple of the empirical formula; the multiplier equals molecular mass ÷ empirical mass
• Required for calculating molecular weight and performing stoichiometric conversions

Structural Formula

• Depicts atom arrangement and bonding—shows which atoms connect to which
• Reveals molecular geometry through Lewis structures, condensed formulas, or line-angle drawings
• Essential for distinguishing isomers—compounds with identical molecular formulas but different structures

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Formulas for Different Compound Types

The type of bonding in a compound determines how its formula is written and interpreted. Ionic compounds use formula units; covalent compounds use molecular formulas.

Ionic Compound Formulas

• Cation written first, anion second—formula reflects the ratio needed to balance charges to zero
• Represent formula units, not molecules; $NaCl$ means a 1:1 ratio in a crystal lattice
• Charges must cancel; for $Al_2O_3$, two $Al^{3+}$ ions balance three $O^{2-}$ ions

Covalent Compound Formulas

• Nonmetals sharing electrons to form discrete molecules with specific atom counts
• Prefixes indicate quantity—mono-, di-, tri-, tetra-, penta-, hexa-, etc.
• Order matters: less electronegative element first (except for certain common compounds like water)

Hydrate Formulas

• Ionic compounds with water molecules trapped in the crystal structure
• Written with a centered dot: $CuSO_4 \cdot 5H_2O$ (copper(II) sulfate pentahydrate)
• Water can be driven off by heating—mass difference reveals water content

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Formula Units and Structural Representations

Understanding what a formula physically represents helps you interpret it correctly. Formula units describe ionic lattices; Lewis structures show electron distribution.

Formula Units

• Smallest repeating unit in an ionic crystal—not a discrete molecule
• Represents the ion ratio that produces electrical neutrality
• Used for all ionic compound calculations—treat the formula unit like you would a molecule for molar mass

Polyatomic Ions

• Charged groups of covalently bonded atoms that act as a single unit in ionic compounds
• Must be memorized—common ones include $NO_3^-$, $SO_4^{2-}$, $PO_4^{3-}$, $NH_4^+$
• Parentheses preserve the unit: $Ca(NO_3)_2$ contains two complete nitrate ions

Lewis Structures

• Diagrams showing bonding pairs and lone pairs of valence electrons
• Follow the octet rule (with exceptions for expanded octets and incomplete octets)
• Predict molecular geometry and reactivity—foundation for VSEPR theory

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Calculations from Formulas

Formulas aren't just labels—they're tools for quantitative analysis. Molar mass connects formulas to measurable quantities.

Molar Mass Calculation

• Sum of all atomic masses in the formula, expressed in $g/mol$
• Use coefficients and subscripts: for $Ca(OH)_2$, calculate $40.08 + 2(16.00 + 1.01) = 74.10 \text{ g/mol}$
• Gateway to stoichiometry—converts between grams and moles for any substance

Percent Composition

• Mass percentage of each element: $\frac{\text{mass of element in 1 mol}}{\text{molar mass of compound}} \times 100\%$
• Used to determine empirical formulas from experimental mass data
• Indicates purity when compared to theoretical values

Stoichiometry Calculations

• Mole ratios from balanced equations connect reactants to products
• Conversion pathway: grams → moles → mole ratio → moles → grams
• Limiting reactant determines actual yield—excess reactant is left over

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Balancing and Oxidation States

Chemical formulas must obey conservation laws. Balancing ensures mass conservation; oxidation numbers track electron distribution.

Balancing Chemical Equations

• Adjust coefficients only—never change subscripts, which would change the compound's identity
• Conserves mass and atoms—same number of each element on both sides
• Systematic approach: balance metals first, then nonmetals, then hydrogen, then oxygen

Oxidation Numbers

• Assigned charges that track electron distribution in compounds and ions
• Key rules: elements = 0; monatomic ions = charge; oxygen = $-2$ (usually); hydrogen = $+1$ (usually)
• Must sum to overall charge—zero for neutral compounds, ion charge for polyatomic ions

Nomenclature Rules

• Systematic naming based on compound type—ionic, covalent, and acids each have distinct rules
• Roman numerals indicate oxidation state for transition metals: $FeCl_3$ = iron(III) chloride
• Acids follow patterns: binary acids use hydro-___-ic; oxyacids use -ic (more O) or -ous (less O)

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Quick Reference Table

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Self-Check Questions

1. A compound has an empirical formula of $CH_2O$ and a molar mass of 180 g/mol. What is its molecular formula, and what calculation connects these two formula types?

2. Which two formula types would you need to distinguish between two isomers that have identical molecular formulas? Why is an empirical formula insufficient?

3. Compare how you would write the formula for magnesium chloride versus dinitrogen tetroxide. What determines whether you use prefixes or charge-balancing?

4. If an FRQ asks you to determine the formula of an unknown compound from combustion analysis data, what sequence of calculations would you perform, and which formula type would you find first?

5. Explain why changing a subscript in a chemical equation is fundamentally different from changing a coefficient. What does each number represent, and what law governs equation balancing?