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

Writing and Predicting Chemical Formulas

Formulas of Binary Ionic Compounds

Binary ionic compounds are formed from the combination of metal cations and nonmetal anions. The chemical formula represents the simplest ratio of ions that results in a neutral compound.

• Cation: A positively charged ion, usually a metal.

• Anion: A negatively charged ion, usually a nonmetal.

• Formula Unit: The simplest whole-number ratio of ions in an ionic compound.

To write the formula, balance the charges so the total positive and negative charges are equal.

Common Ion Charges by Group

Group	1A	2A	3A	5A	6A	7A
Cation	Na+, K+	Mg2+, Ca2+	Al3+
Anion				N3−, P3−	O2−, S2−	F−, Cl−, Br−, I−

Additional info: Charges are inferred from periodic table group numbers.

Steps to Write Ionic Compound Formulas

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

2. Balance the charges by adjusting subscripts so the total charge is zero.

3. Write the formula with the cation first, followed by the anion.

Example: Magnesium chloride: Mg2+ and Cl− combine to form MgCl2.

Naming Ionic Compounds

Rules for Naming

• Name the cation first, then the anion.

• For monatomic anions, use the root of the element name plus the suffix “-ide”.

• For polyatomic ions, use the name of the ion as is.

Example: NaCl is named sodium chloride.

Polyatomic Ions

Polyatomic ions are groups of atoms that carry a charge and act as a single unit in compounds.

Ion	Formula	Charge
Ammonium	NH4+	+1
Sulfate	SO42−	−2
Nitrate	NO3−	−1
Phosphate	PO43−	−3

Compounds Containing Polyatomic Ions

Writing Formulas

• Use parentheses around polyatomic ions when more than one is needed.

• Balance the charges as with binary ionic compounds.

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

Acids and Acid Nomenclature

Binary Acids

• Composed of hydrogen and one other nonmetal.

• Named as “hydro-” + root of nonmetal + “-ic acid”.

Example: HCl is hydrochloric acid.

Oxyacids

• Contain hydrogen, oxygen, and another element (usually a nonmetal).

• If the polyatomic ion ends in “-ate”, the acid name ends in “-ic acid”.

• If the polyatomic ion ends in “-ite”, the acid name ends in “-ous acid”.

Example: H2SO4 (sulfate) is sulfuric acid; H2SO3 (sulfite) is sulfurous acid.

Stoichiometry and Chemical Calculations

Basic Repeating Units

• Atom: The smallest unit of an element.

• Molecule: The smallest unit of a covalent compound.

• Formula Unit: The smallest unit of an ionic compound.

The Mole and Avogadro’s Number

The mole is a counting unit in chemistry, representing $6.022 \times 10^{23}$ particles (atoms, molecules, or formula units).

• Avogadro’s Number: $6.022 \times 10^{23}$

• Mole (mol): The amount of substance containing Avogadro’s number of entities.

Molar Mass

Molar mass is 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 g/mol) + 16.00 g/mol = 18.02 g/mol.

Stoichiometric Calculations

• Use molar mass to convert between mass and moles.

• Use Avogadro’s number to convert between moles and number of particles.

Key Equations:

• $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass (g/mol)}}$

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

Percent Composition

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

• $\%\,\text{Element} = \frac{\text{Mass of element in 1 mol of compound}}{\text{Molar mass of compound}} \times 100$

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.

Steps to Determine Empirical Formula

1. Convert mass 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 the simplest ratio.

Summary Tables

Common Polyatomic Ions

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

Common Acid Names

Acid	Formula	Name
HCl	Hydrochloric acid	Binary acid
H2SO4	Sulfuric acid	Oxyacid
HNO3	Nitric acid	Oxyacid
H2CO3	Carbonic acid	Oxyacid

Key Concepts and Applications

• Writing chemical formulas requires understanding ion charges and balancing them for neutrality.

• Naming compounds follows systematic rules based on the type of compound (ionic, covalent, acid).

• Stoichiometry allows conversion between mass, moles, and number of particles using molar mass and Avogadro’s number.

• Percent composition and empirical formulas are essential for analyzing chemical compounds.

Example Application: Calculate the number of moles in 18 g of water: $\text{Moles} = \frac{18\,\text{g}}{18.02\,\text{g/mol}} = 1.00\,\text{mol}$

Additional info: These notes cover topics from chapters on chemical composition, chemical reactions, and quantities in chemical reactions, as well as nomenclature and formula writing.