Writing and Predicting Chemical Formulas, Stoichiometry, and Nomenclature: Study Notes for Introductory Chemistry

Chemical Formulas and Nomenclature

Writing and Predicting Chemical Formulas

Chemical formulas represent the types and numbers of atoms in a compound. Predicting and writing formulas requires understanding the charges of ions and the rules of combination for ionic and covalent compounds.

• Ionic Compounds: Formed from metals and nonmetals. The total positive charge must balance the total negative charge.

• Covalent Compounds: Formed from nonmetals. Atoms share electrons; formulas reflect the actual number of atoms.

• Polyatomic Ions: Groups of atoms that carry a charge and act as a single unit in compounds (e.g., SO42-, NH4+).

Common Ion Charges by Group

The charge of ions can often be predicted by their group in the periodic table:

Group	Common Ion Charge
IA (Alkali metals)	+1
IIA (Alkaline earth metals)	+2
IIIA	+3
IVA	+4 or -4
VA	-3
VIA	-2
VIIA (Halogens)	-1
VIIIA (Noble gases)	0

Example: Sodium (Na, group IA) forms Na+; Chlorine (Cl, group VIIA) forms Cl-; the formula for sodium chloride is NaCl.

Formulas for Compounds with Polyatomic Ions

When writing formulas for compounds containing polyatomic ions, use parentheses if more than one of the ion is needed.

• Example: Calcium nitrate: Ca2+ and NO3- combine to form Ca(NO3)2.

Naming Chemical Compounds

Nomenclature is the system of naming chemical compounds. The rules differ for ionic and covalent compounds.

• Ionic Compounds: Name the cation (metal) first, then the anion (nonmetal or polyatomic ion). For transition metals, indicate the charge with Roman numerals.

• Covalent Compounds: Use prefixes to indicate the number of each atom (mono-, di-, tri-, etc.).

• Acids: If the anion ends in -ide, the acid name begins with 'hydro-' and ends with '-ic acid'. If the anion ends in -ate or -ite, use '-ic acid' or '-ous acid' respectively.

Common Polyatomic Ions

Name	Formula	Charge
Ammonium	NH4	+1
Nitrate	NO3	-1
Sulfate	SO4	-2
Phosphate	PO4	-3
Hydroxide	OH	-1
Carbonate	CO3	-2

Chemical Composition and Stoichiometry

Stoichiometric Calculations

Stoichiometry involves quantitative relationships between reactants and products in chemical reactions. Calculations often use the mole concept and Avogadro's number.

• Mole: The amount of substance containing $6.022 \times 10^{23}$ particles (Avogadro's number).

• Molar Mass (MM): The mass of one mole of a substance, expressed in grams per mole (g/mol).

• Formula Mass: The sum of atomic masses of all atoms in a chemical formula.

Basic Equations

• Number of moles: $n = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• Number of particles: $\text{Number of particles} = n \times 6.022 \times 10^{23}$

• Mass from moles: $\text{mass (g)} = n \times \text{molar mass (g/mol)}$

Example Calculation

• How many moles are in 18 g of water (H2O)? Molar mass of H2O = 18 g/mol $n = \frac{18}{18} = 1$ mole

Percent Composition

Percent composition expresses the mass percentage of each element in a compound.

• Formula: $\%\,\text{element} = \frac{\text{mass of element in 1 mole}}{\text{molar mass of compound}} \times 100$

• Example: In H2O, mass of H = 2 g, mass of O = 16 g, molar mass = 18 g. $\%\,\text{H} = \frac{2}{18} \times 100 = 11.1\%$ $\%\,\text{O} = \frac{16}{18} \times 100 = 88.9\%$

Empirical and Molecular Formulas

The empirical formula shows the simplest whole-number ratio of atoms in a compound. The molecular formula shows the actual number of atoms.

• Example: Glucose has a molecular formula C6H12O6 and an empirical formula CH2O.

• To determine empirical formula from percent composition:

  1. Convert percentages to grams (assume 100 g sample).

  2. Convert grams to moles for each element.

  3. Divide by the smallest number of moles to get ratios.

  4. Round to nearest whole number.

Tables and Reference Data

Selected Tables from Notes

Table: Common Polyatomic Ions

Ion	Formula	Charge
Ammonium	NH4	+1
Nitrate	NO3	-1
Sulfate	SO4	-2
Phosphate	PO4	-3
Hydroxide	OH	-1
Carbonate	CO3	-2

Table: Prefixes for Covalent Compounds

Number	Prefix
1	Mono-
2	Di-
3	Tri-
4	Tetra-
5	Penta-
6	Hexa-
7	Hepta-
8	Octa-
9	Nona-
10	Deca-

Additional info:

• These notes cover topics from chapters 5 (Molecules and Compounds) and 6 (Chemical Composition), and also include stoichiometry (chapter 8).

• Worked examples and practice problems are included throughout to reinforce concepts.

• Tables for common ions, prefixes, and percent composition calculations are provided for reference.