Writing and Predicting Chemical Formulas, Nomenclature, and Stoichiometry

Writing and Predicting Chemical Formulas, Nomenclature, and Stoichiometry

Forming Ionic Compounds and Predicting Formulas

Ionic compounds are formed by the combination of cations (positively charged ions) and anions (negatively charged ions) in such a way that the overall charge of the compound is neutral. The chemical formula of an ionic compound reflects the simplest ratio of ions that achieves charge neutrality.

• Cations are typically formed by metals losing electrons.

• Anions are typically formed by nonmetals gaining electrons.

• The charges of ions can often be predicted based on their group in the periodic table.

Example: Sodium (Na) forms Na+ and chlorine (Cl) forms Cl-. The formula for sodium chloride is NaCl.

Common Ion Charges by Group

Group	1A	2A	3A	5A	6A	7A
Ion	+1	+2	+3	-3	-2	-1
Examples	Na+, K+	Mg2+, Ca2+	Al3+	N3-	O2-, S2-	F-, Cl-

Steps for Writing Ionic Formulas

1. Write the symbol and charge for the cation and anion.

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

3. Write the formula using subscripts to indicate the number of each ion needed.

Example: For magnesium chloride, Mg forms Mg2+ and Cl forms Cl-. Two Cl- ions are needed to balance one Mg2+: MgCl2.

Polyatomic Ions

Polyatomic ions are ions composed of more than one atom covalently bonded, carrying a net charge. They act as a single unit in compounds.

• Common polyatomic ions: NO3- (nitrate), SO42- (sulfate), CO32- (carbonate), NH4+ (ammonium).

• When more than one polyatomic ion is needed, use parentheses: e.g., Ca(NO3)2.

Table: Common Polyatomic Ions

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

Naming Ionic Compounds

The name of an ionic compound is constructed from the names of its ions.

• Name the cation first, then the anion.

• For monatomic anions, change the ending to "-ide" (e.g., chloride, oxide).

• For polyatomic ions, use the ion's name (e.g., sulfate, nitrate).

• For metals with variable charge (transition metals), indicate the charge with Roman numerals (e.g., FeCl2 is iron(II) chloride).

Writing Formulas for Compounds with Polyatomic Ions

When writing formulas for compounds containing polyatomic ions, the same charge balancing rules apply. Parentheses are used if more than one polyatomic ion is needed.

• Example: Aluminum sulfate: Al3+ and SO42-. The formula is Al2(SO4)3.

Acids and Their Nomenclature

Acids are compounds that release H+ ions in water. Their names depend on the anion present:

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

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

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

Table: Acid Nomenclature Examples

Acid Formula	Anion	Acid Name
HCl	Cl- (chloride)	Hydrochloric acid
HNO3	NO3- (nitrate)	Nitric acid
HNO2	NO2- (nitrite)	Nitrous acid
H2SO4	SO42- (sulfate)	Sulfuric acid
H2SO3	SO32- (sulfite)	Sulfurous acid

Stoichiometry: Calculations Involving Elements and Compounds

Stoichiometry involves quantitative relationships between reactants and products in chemical reactions. It is essential for determining the amounts of substances involved in reactions.

The Mole and Avogadro's Number

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

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

Example: 1 mol of H2O contains $6.022 \times 10^{23}$ molecules of water.

Molar Mass

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

• Calculated by summing the atomic masses of all atoms in the formula.

Example: Molar mass of H2O = 2(1.01) + 16.00 = 18.02 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: $\text{particles} = \text{moles} \times 6.022 \times 10^{23}$

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

Percent Composition

• Percent composition: The percent by mass of each element in a compound.

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

Example: For H2O, percent 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.

• To determine the empirical formula from percent composition, convert percentages to grams, then to moles, and find the simplest ratio.

Example: A compound with 40% C, 6.7% H, and 53.3% O has an empirical formula of CH2O.

Stoichiometric Calculations

• Use balanced chemical equations to relate moles of reactants and products.

• Convert given quantities to moles, use mole ratios, then convert to desired units.

Example: For the reaction $2H_2 + O_2 \rightarrow 2H_2O$, 4 mol H2 produces 4 mol H2O.

Summary Table: Key Conversions

From	To	Conversion Factor
Mass (g)	Moles	$\frac{1\,\text{mol}}{\text{molar mass (g)}}$
Moles	Particles	$6.022 \times 10^{23}$ particles/1 mol
Moles	Mass (g)	Molar mass (g)/1 mol

Practice Problems and Examples

• Numerous worked examples are provided throughout the notes, illustrating how to write formulas, name compounds, and perform stoichiometric calculations.

• Practice problems reinforce the application of concepts such as formula writing, naming, and quantitative calculations.

Additional info: These notes cover material relevant to the following Introduction to Chemistry topics: chemical composition, chemical reactions, quantities in chemical reactions, and chemical nomenclature. Tables have been reconstructed and summarized for clarity.