Naming Compounds, Acids, and Chemical Formulas: A General Chemistry Study Guide

Naming Binary Molecular Compounds

Rules and Prefixes

Binary molecular compounds are composed of two nonmetal elements. Their naming follows specific rules and uses prefixes to indicate the number of atoms present.

• Rule 1: The first element in the formula is named using its full element name.

• Rule 2: The second element is named as if it were an anion (ending with -ide).

• Rule 3: Prefixes are used to denote the number of atoms of each element. The prefix mono- is never used for the first element.

Number	Prefix
1	mono
2	di
3	tri
4	tetra
5	penta
6	hexa
7	hepta
8	octa
9	nona
10	deca

• Note: Prefixes are only used when there is no metal present.

Examples

Compound	Systematic Name	Common Name
N2O	Dinitrogen monoxide	Nitrous oxide
NO	Nitrogen monoxide	Nitric oxide
NO2	Nitrogen dioxide
N2O3	Dinitrogen trioxide
N2O4	Dinitrogen tetroxide
N2O5	Dinitrogen pentoxide

Example: CO is called carbon monoxide, not monocarbon monoxide. SiO2 is called silicon dioxide, not monosilicon dioxide.

Formulas from Names and Vice Versa

Practice and Application

It is important to be able to write chemical formulas from systematic names and to name compounds from their formulas.

• Example: Vanadium (V) fluoride → VF5

• Example: Dioxygen difluoride → O2F2

• Example: Rubidium peroxide → Rb2O2

• Example: Gallium oxide → Ga2O3

Additional info: Practice converting between names and formulas is essential for mastering chemical nomenclature.

Naming Acids

Types and Rules

Acids are compounds that release H+ ions in solution. Their names depend on the presence or absence of oxygen in the anion.

• Type 1: Anion does not contain oxygen

  • The acid is named with the prefix hydro- and the suffix -ic.

  • Examples: HCl = hydrochloric acid, HBr = hydrobromic acid, HI = hydroiodic acid, H2S = hydrosulfuric acid, HCN = hydrocyanic acid

• Type 2: Anion contains oxygen

  • The acid name is formed from the root of the anion name with the suffix -ic or -ous depending on the ending of the anion:

    • If the anion ends in -ate, the acid name ends in -ic.

    • If the anion ends in -ite, the acid name ends in -ous.

  • Examples:

    • H2SO4 (sulfate, SO42-) → sulfuric acid

    • H2SO3 (sulfite, SO32-) → sulfurous acid

Acid	Anion	Name
HClO4	perchlorate	perchloric acid
HClO3	chlorate	chloric acid
HClO2	chlorite	chlorous acid
HClO	hypochlorite	hypochlorous acid

Hydrated Ionic Compounds

Definition and Examples

Hydrates are ionic compounds that include a specific number of water molecules in their structure. The number of water molecules is indicated after a dot in the formula.

• Example: CuSO4·5H2O is called copper(II) sulfate pentahydrate.

Molar Mass

Definition and Calculation

Molar mass is the mass per mole of any material (element, molecule, etc.), often called molecular weight. It is calculated by summing the atomic masses of all elements in the molecule.

• Formula: $\text{Molar mass} = \sum (\text{number of atoms of element}) \times (\text{atomic mass of element})$

Examples:

• Water (H2O): $2 \times 1.01 + 16.00 = 18.02$ g/mol

• Table salt (NaCl): $22.99 + 35.45 = 58.44$ g/mol

• Calcium carbonate (CaCO3): $40.08 + 12.01 + 3 \times 16.00 = 100.09$ g/mol

• Acetone (C3H6O): $3 \times 12.01 + 6 \times 1.01 + 16.00 = 58.08$ g/mol

• Caffeine (C8H10N4O2): $8 \times 12.01 + 10 \times 1.01 + 4 \times 14.01 + 2 \times 16.00 = 194.19$ g/mol

• Acetaminophen (C8H9NO2): $8 \times 12.01 + 9 \times 1.01 + 14.01 + 2 \times 16.00 = 151.16$ g/mol

Mass Percentage Composition

Definition and Calculation

The mass percentage of an element in a compound is the mass of the element divided by the total mass of the compound, multiplied by 100%.

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

Examples:

• Water (H2O): Hydrogen = $\frac{2.02}{18.02} \times 100\% = 11.2\%$, Oxygen = $\frac{16.00}{18.02} \times 100\% = 88.8\%$

• Barium sulfate (BaSO4): Calculate mass % for Ba, S, and O.

• Methanol (CH4O): Calculate mass % for C, H, and O.

• Caffeine (C8H10N4O2): Calculate mass % for C, H, N, and O.

Note: The sum of all mass percentages in a compound must equal 100%.

Empirical and Molecular Formulas

Definitions

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

Examples:

• Ethane (C2H6): Empirical formula = CH3

• Benzene (C6H6): Empirical formula = CH

• Caffeine (C8H10N4O2): Empirical formula = C4H5N2O

Determining Empirical Formula from Combustion Analysis

Combustion analysis is used to determine the empirical formula of a compound by measuring the masses of CO2 and H2O produced from a known mass of sample.

1. Base the calculation on 100 grams of compound (each percent = mass in grams).

2. Determine the number of moles of each element using atomic masses.

3. Divide each value by the smallest number of moles calculated.

4. If necessary, multiply by an integer to obtain whole numbers.

Acceptable rounding: 9.92 rounds to 10, 1.08 rounds to 1. Do not round common fractions (e.g., 1.25, 1.33, 1.5).

Determining Molecular Formula

1. Obtain the empirical formula.

2. Compute the mass of the empirical formula.

3. Calculate the ratio: $\frac{\text{molar mass (molecular weight)}}{\text{empirical formula mass}}$

4. The integer from step 3 is the number of empirical formula units in the molecular formula.

Example: If the empirical formula mass is 30 g/mol and the molar mass is 60 g/mol, the molecular formula is twice the empirical formula.

Additional info: These procedures are fundamental for analyzing unknown compounds and are widely used in laboratory settings.