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.

For example, sodium (Na) forms Na+ and chlorine (Cl) forms Cl-, combining to form NaCl.

Table: Common Charges of Main Group Elements

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

Naming Ionic Compounds

The name of an ionic compound is constructed from the names of its constituent ions. The cation is named first, followed by the anion. For monatomic anions, the ending is changed to '-ide'.

• Example: NaCl is named sodium chloride.

• For compounds containing polyatomic ions, use the name of the polyatomic ion as is (e.g., Na2SO4 is sodium sulfate).

Writing Formulas for Ionic Compounds

To write the formula for an ionic compound:

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 using subscripts to indicate the number of each ion needed for neutrality.

Example: To write the formula for magnesium chloride:

• Mg forms Mg2+, Cl forms Cl-.

• Two Cl- ions are needed to balance one Mg2+ ion.

• Formula: MgCl2

Polyatomic Ions and Compounds Containing Polyatomic Ions

Polyatomic ions are ions composed of more than one atom. They act as a single unit with a specific charge. Common polyatomic ions include sulfate (SO42-), nitrate (NO3-), and ammonium (NH4+).

Table: Common Polyatomic Ions

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

When writing formulas with polyatomic ions, use parentheses if more than one polyatomic ion is needed.

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

Acids and Bases: Naming and Formulas

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

• If the anion ends in '-ide', the acid name is 'hydro-' + root + '-ic acid' (e.g., HCl: hydrochloric acid).

• If the anion ends in '-ate', the acid name is root + '-ic acid' (e.g., H2SO4: sulfuric acid).

• If the anion ends in '-ite', the acid name is root + '-ous acid' (e.g., H2SO3: sulfurous acid).

Table: Naming Acids Based on Anion

Anion Ending	Acid Name	Example
-ide	hydro-...-ic acid	HCl: hydrochloric acid
-ate	...-ic acid	HNO3: nitric acid
-ite	...-ous acid	HNO2: nitrous acid

Stoichiometry: Calculations Involving Elements and Compounds

Stoichiometry involves the calculation of quantities in chemical reactions, including the determination of the number of atoms, molecules, or formula units in a given sample, and the relationships between mass, moles, and number of particles.

The Mole and Avogadro's Number

• Mole (mol): The amount of substance containing as many entities (atoms, molecules, ions) as there are atoms in 12 g of carbon-12.

• Avogadro's Number: $6.022 \times 10^{23}$ entities per mole.

Key Equations

• Number of particles = moles $\times$ Avogadro's number

• Mass (g) = moles $\times$ molar mass (g/mol)

• Moles = $\dfrac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

Calculating Formula Mass and Molecular Mass

The formula mass (for ionic compounds) or molecular mass (for covalent compounds) is the sum of the atomic masses of all atoms in the formula.

• Example: The formula mass of NaCl is $22.99 + 35.45 = 58.44$ g/mol.

Percent Composition

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

• Percent by mass = $\dfrac{\text{mass of element in 1 mol of compound}}{\text{molar mass of compound}} \times 100\%$

Empirical and Molecular Formulas

The empirical formula gives the simplest whole-number ratio of atoms in a compound, while the molecular formula gives 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.

• The molecular formula is a whole-number multiple of the empirical formula.

Summary Table: Steps for Writing and Naming Compounds

Type of Compound	How to Write Formula	How to Name
Ionic (metal + nonmetal)	Balance charges, write symbols, use subscripts	Cation name + anion name (-ide)
Covalent (nonmetal + nonmetal)	Use prefixes to indicate number of atoms	Prefix + element name + prefix + element name (-ide)
Acid	H + anion	See acid naming rules above

Worked Examples

• Example 1: Write the formula for aluminum oxide.

  • Al3+ and O2- combine to form Al2O3.

• Example 2: Calculate the number of molecules in 2.0 mol of CO2.

  • Number of molecules = $2.0 \times 6.022 \times 10^{23} = 1.204 \times 10^{24}$ molecules

• Example 3: Find the percent composition of H in H2O.

  • H: $2 \times 1.01 = 2.02$ g; O: $16.00$ g; total = $18.02$ g

  • Percent H = $\dfrac{2.02}{18.02} \times 100\% = 11.2\%$

Additional info: These notes include expanded explanations, tables, and examples to clarify the process of writing chemical formulas, naming compounds, and performing basic stoichiometric calculations, as covered in introductory college chemistry.