Writing and Predicting Chemical Formulas, Nomenclature, and Composition Stoichiometry

Writing and Predicting Chemical Formulas

Formulas of Binary Ionic Compounds

Binary ionic compounds are composed of two elements: a metal (cation) and a nonmetal (anion). The formula is written by balancing the charges of the ions so that the compound is electrically neutral.

• Cation: Positively charged ion, usually a metal.

• Anion: Negatively charged ion, usually a nonmetal.

• Charges are determined by the group number in the periodic table.

For example, sodium (Na) in Group IA forms Na+, and chlorine (Cl) in Group VIIA forms Cl-. The formula for sodium chloride is NaCl.

Common Charges of Main Group Elements

Group	IA	IIA	IIIA	IVA	VA	VIA	VIIA	VIIIA
Cation/Anion	+1	+2	+3	±4	-3	-2	-1	0
Examples	Na+, K+	Mg2+, Ca2+	Al3+	C4+, C4-	N3-	O2-	F-, Cl-	He, Ne

Additional info: Transition metals may have variable charges, indicated by Roman numerals in their names (e.g., Fe(III) for Fe3+).

Steps for Writing Ionic Formulas

1. Write the symbol and charge for each ion.

2. Balance the charges so the total positive and negative charges are equal.

3. Write the formula with the cation first, followed by the anion. Use subscripts to indicate the number of each ion needed.

Example: Write the formula for magnesium chloride. Mg forms Mg2+, Cl forms Cl-. Two Cl- ions are needed to balance one Mg2+: $\text{Mg}^{2+} + 2\text{Cl}^- \rightarrow \text{MgCl}_2$

Nomenclature of Ionic Compounds

Rules for Naming Ionic Compounds

• Name the cation first, then the anion.

• The cation keeps its elemental name (e.g., sodium, magnesium).

• The anion name ends in "-ide" (e.g., chloride, oxide).

• For transition metals with variable charge, indicate the charge with Roman numerals (e.g., iron(III) chloride for FeCl3).

Example: NaCl is sodium chloride; FeCl2 is iron(II) chloride.

Polyatomic Ions and Compounds

Common Polyatomic Ions

Polyatomic ions are ions composed of more than one atom. They act as a single unit with a specific charge.

Ion	Formula	Name
NO3-	NO3-	Nitrate
SO42-	SO42-	Sulfate
CO32-	CO32-	Carbonate
PO43-	PO43-	Phosphate
NH4+	NH4+	Ammonium

When writing formulas with polyatomic ions, use parentheses if more than one polyatomic ion is needed (e.g., Ca(NO3)2).

Naming Compounds with Polyatomic Ions

• Name the cation first, then the polyatomic anion.

• Do not change the ending of the polyatomic ion's name.

Example: NaNO3 is sodium nitrate; CaSO4 is calcium sulfate.

Acids and Bases: Nomenclature

Naming Acids

• If the anion ends in "-ide": use the prefix "hydro-" and the suffix "-ic acid" (e.g., HCl is hydrochloric acid).

• If the anion ends in "-ate": use the suffix "-ic acid" (e.g., H2SO4 is sulfuric acid).

• If the anion ends in "-ite": use the suffix "-ous acid" (e.g., H2SO3 is sulfurous acid).

Calculations Involving Elements and Compounds (Composition Stoichiometry)

The Mole and Avogadro's Number

The mole is the SI unit for amount of substance. One mole contains Avogadro's number of particles (6.022 × 1023).

• Avogadro's Number: $6.022 \times 10^{23}$ particles/mol

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

Example: 1 mole of H2O contains $6.022 \times 10^{23}$ molecules and has a mass of 18.02 g.

Calculating Molar Mass

• Add the atomic masses of all atoms in the formula.

Example: Molar mass of CO2: $\text{MM} = 12.01\,\text{g/mol (C)} + 2 \times 16.00\,\text{g/mol (O)} = 44.01\,\text{g/mol}$

Conversions Between Mass, Moles, and Number of Particles

• To convert mass to moles: $\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• To convert moles to number of particles: $\text{particles} = \text{moles} \times 6.022 \times 10^{23}$

Example: How many molecules are in 36.0 g of H2O? $\text{moles} = \frac{36.0}{18.02} = 2.00$ mol $\text{molecules} = 2.00 \times 6.022 \times 10^{23} = 1.20 \times 10^{24}$ molecules

Percent Composition

Percent composition is the percent by mass of each element in a compound.

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

Example: Percent composition of H in H2O: $\%\,\text{H} = \frac{2.02}{18.02} \times 100\% = 11.2\%$

Empirical and Molecular Formulas

• Empirical formula: The simplest whole-number ratio of atoms in a compound.

• Molecular formula: The actual number of atoms of each element in a molecule.

Example: The empirical formula of glucose (C6H12O6) is CH2O.

Determining Empirical Formula from Percent Composition

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

2. Convert grams to moles for each element.

3. Divide by the smallest number of moles to get the simplest ratio.

Example: A compound contains 40% C, 6.7% H, and 53.3% O. Convert to moles: C: $\frac{40}{12.01} = 3.33$ mol H: $\frac{6.7}{1.01} = 6.64$ mol O: $\frac{53.3}{16.00} = 3.33$ mol Divide by 3.33: C 1, H 2, O 1. Empirical formula is CH2O.

Summary Table: Common Polyatomic Ions

Name	Formula	Charge
Nitrate	NO3-	-1
Sulfate	SO42-	-2
Carbonate	CO32-	-2
Phosphate	PO43-	-3
Hydroxide	OH-	-1
Ammonium	NH4+	+1

Practice Problems and Worked Examples

• Write the formula for calcium nitrate: Ca(NO3)2

• Name the compound Fe2(SO4)3: iron(III) sulfate

• Calculate the number of moles in 18.0 g of water: $\frac{18.0}{18.02} = 1.00$ mol

• Determine the empirical formula for a compound with 27.3% C, 72.7% O: C: $\frac{27.3}{12.01} = 2.27$ mol O: $\frac{72.7}{16.00} = 4.54$ mol Ratio: 1:2, so empirical formula is CO2

Additional info: The notes also include tables of variable charge metals, naming conventions for acids, and stepwise examples for all calculations.