Introduction to Chemistry: Nomenclature, Formula Writing, and Mole Calculations Study Guide

Chapter 5: Nomenclature and Formula Writing

Naming Ionic and Molecular Compounds

Understanding how to name chemical compounds is a foundational skill in chemistry. The process differs for ionic and molecular compounds, and it is important to follow systematic rules to avoid common mistakes.

• Ionic Compounds: Composed of metals and nonmetals. The metal (cation) is named first, followed by the nonmetal (anion) with its ending changed to -ide. If the metal can have more than one charge (transition metals), a Roman numeral is used to indicate its charge.

• Molecular Compounds: Composed of nonmetals only. Greek prefixes are used to indicate the number of each atom present (e.g., di- for two, tri- for three).

• Polyatomic Ions: Compounds containing polyatomic ions retain the name of the ion (e.g., sulfate, phosphate).

Example: SO2 is named sulfur dioxide (molecular compound), while FeCl3 is iron(III) chloride (ionic compound with transition metal).

Table: Sample Compounds and Their Names

Compound	Name
CoBr2	Cobalt(II) bromide
BeBr2	Beryllium bromide
VPO4	Vanadium(III) phosphate
Cr2(SO4)3	Chromium(III) sulfate
Al2(SO4)3	Aluminum sulfate
TiS	Titanium(II) sulfide
PCl3	Phosphorus trichloride

Common Mistakes:

• Using Roman numerals for molecular compounds (not needed).

• Forgetting to use Greek prefixes for molecular compounds.

• Not using the correct charge for transition metals.

Writing Formulas from Compound Names

To write the chemical formula from a compound's name, first determine if the compound is ionic or molecular. Use the appropriate rules for each type.

• Ionic Compounds: Use the charges to determine subscripts so the total charge is zero.

• Molecular Compounds: Use Greek prefixes to determine the number of each atom.

Example: Potassium sulfide is K2S (ionic, K+ and S2−), while selenium tetrabromide is SeBr4 (molecular, prefix tetra- means four Br atoms).

Table: Compound Names and Formulas

Name	Formula
Barium nitrite	Ba(NO2)2
Selenium tetrabromide	SeBr4
Potassium sulfide	K2S
Aluminum chloride	AlCl3
Dinitrogen pentoxide	N2O5

Chapter 6: Mole Calculations and Empirical/Molecular Formulas

Calculating Mass, Moles, and Number of Molecules

Converting between mass, moles, and number of molecules is a key skill in chemistry. The relationships are based on the molar mass and Avogadro's number.

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

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

Key Equations:

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

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

Example: To find the number of molecules in 0.0153 moles of CO2:

• $0.0153 \text{ mol} \times 6.022 \times 10^{23} = 9.21 \times 10^{21}$ molecules$

Stepwise Conversion Process

1. Write units of answer and starting unit.

2. Add units to conversion factor to cancel out starting units and put in answer units.

3. Relate top to bottom in conversion factor.

Example: To convert 6.33 kg N2O to moles:

• Convert kg to g: $6.33 \text{ kg} \times \frac{1000 \text{ g}}{1 \text{ kg}} = 6330 \text{ g}$

• Convert g to moles: $\frac{6330 \text{ g}}{44.0 \text{ g/mol}} = 144 \text{ moles}$

Empirical and Molecular Formulas

The empirical formula is the simplest whole-number ratio of atoms in a compound. The molecular formula is the actual number of atoms of each element in a molecule.

• To determine the empirical formula, convert the mass of each element to moles, then divide by the smallest number of moles to get the ratio.

• To determine the molecular formula, divide the molar mass of the compound by the molar mass of the empirical formula, then multiply the subscripts in the empirical formula by this integer.

Key Equations:

• Empirical formula calculation: $\text{Empirical formula} = \text{smallest whole-number ratio of moles of each element}$

• Molecular formula calculation: $\text{Molecular formula} = (\text{Empirical formula})_n$, where $n = \frac{\text{Molar mass of compound}}{\text{Molar mass of empirical formula}}$

Example: A compound with 62.04 g C, 10.44 g H, and 27.52 g O has an empirical formula of C5H10O2. If the molar mass is 290 g/mol, the molecular formula is C10H20O4.

Empirical Formula from Experimental Data

Given the mass of each element in a compound, the empirical formula can be determined by converting each mass to moles and finding the simplest ratio.

Example: A compound contains 2.231 g Se and 3.221 g F. The empirical formula is SeF6.

Common Mistakes

• Using grams instead of moles for subscripts in formulas.

• Stopping at the empirical formula and forgetting to calculate the molecular formula.

• Not using scientific notation correctly in calculations.

Additional info: These notes are based on review questions and answers for an introductory chemistry course, focusing on nomenclature, formula writing, and mole calculations, which are core topics in general chemistry.