General Chemistry Study Guide: Nomenclature, Conversions, and Solution Concentrations

Inorganic Nomenclature

Systematic Naming of Inorganic Salts and Acids

Inorganic nomenclature is the system used to name chemical compounds, particularly salts and acids, based on their composition and structure. Understanding these rules is essential for clear communication in chemistry.

• Inorganic Salts: Named by stating the cation (positive ion) first, followed by the anion (negative ion).

• Acids: Named based on the anion present. For example, anions ending in -ide become hydro-...-ic acid (e.g., HCl: hydrochloric acid), while -ate becomes -ic acid (e.g., H2SO4: sulfuric acid), and -ite becomes -ous acid (e.g., H2SO3: sulfurous acid).

• Polyatomic Ions: Many compounds contain polyatomic ions, which are ions composed of more than one atom. Examples include sulfate (SO42−), nitrate (NO3−), and ammonium (NH4+).

Example: Na2SO4 is named sodium sulfate.

Table: Common Polyatomic Ions

Polyatomic Ion Name	Polyatomic Ion Formula
Ammonium	NH4+
Carbonate	CO32−
Oxalate	C2O42−
Hydroxide	OH−
Nitrate	NO3−
Sulfate	SO42−
Phosphate	PO43−
Acetate	CH3COO−
Chlorate	ClO3−
Permanganate	MnO4−

Unit Conversions and Molar Mass

Unit Conversions

Unit conversions are essential for solving quantitative problems in chemistry. They involve changing a measurement from one unit to another using conversion factors.

• Conversion Factor: A ratio that expresses how many of one unit are equal to another unit (e.g., 1 inch = 2.54 cm).

• Dimensional Analysis: A method that uses conversion factors to move from one unit to another.

Example: To convert 5.0 grams to milligrams: $5.0\ \text{g} \times \frac{1000\ \text{mg}}{1\ \text{g}} = 5000\ \text{mg}$

Molecular, Formula, and Molar Mass

The mass of a chemical compound can be described in several ways:

• Molecular Mass: The sum of the atomic masses of all atoms in a molecule (for covalent compounds).

• Formula Mass: The sum of the atomic masses of all atoms in a formula unit (for ionic compounds).

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

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

Example: Molar mass of H2O: $2 \times 1.01\ \text{g/mol} + 16.00\ \text{g/mol} = 18.02\ \text{g/mol}$

Avogadro's Number and Counting Particles

Avogadro's number ($6.022 \times 10^{23}$) is the number of particles (atoms, molecules, ions) in one mole of a substance.

• To find the number of particles: $\text{Number of particles} = \text{moles} \times 6.022 \times 10^{23}$

• To find moles from particles: $\text{Moles} = \frac{\text{Number of particles}}{6.022 \times 10^{23}}$

Example: Number of molecules in 0.50 mol of CO2: $0.50 \times 6.022 \times 10^{23} = 3.01 \times 10^{23}$ molecules

Calculating Concentrations

Molarity (M)

Molarity is the most common unit of concentration in chemistry, defined as moles of solute per liter of solution.

• Formula: $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

Example: 0.25 mol NaCl dissolved in 0.500 L solution: $M = \frac{0.25}{0.500} = 0.50\ \text{M}$

Mole Fraction (X)

The mole fraction is the ratio of the moles of one component to the total moles in the mixture.

• Formula: $X_A = \frac{n_A}{n_{\text{total}}}$

Example: In a solution with 1 mol NaCl and 9 mol H2O: $X_{\text{NaCl}} = \frac{1}{1+9} = 0.10$

Mass Percent (%)

Mass percent expresses the mass of solute as a percentage of the total mass of the solution.

• Formula: $\text{Mass Percent} = \frac{\text{mass of solute}}{\text{total mass of solution}} \times 100\%$

Example: 5 g NaCl in 100 g solution: $\frac{5}{100} \times 100\% = 5\%$

Density and Solution Calculations

Density is often used to relate mass and volume in solution calculations.

• Formula: $\text{Density} = \frac{\text{mass}}{\text{volume}}$

Example: If a solution has a density of 1.10 g/mL and a volume of 250 mL, its mass is $1.10 \times 250 = 275$ g.

Dilution Calculations

Preparing Dilute Solutions

Dilution involves adding solvent to a solution to decrease its concentration. The relationship between the concentrations and volumes before and after dilution is given by:

• Formula: $M_1V_1 = M_2V_2$

• $M_1$ = initial molarity, $V_1$ = initial volume

• $M_2$ = final molarity, $V_2$ = final volume

Example: To prepare 500 mL of 0.10 M NaCl from a 1.0 M stock solution: $V_1 = \frac{M_2V_2}{M_1} = \frac{0.10 \times 0.500}{1.0} = 0.050$ L = 50 mL

Calculating Volume of Water to Add

To achieve a desired concentration, calculate the volume of solvent to add using the dilution formula.

• Formula: $V_{\text{added}} = V_2 - V_1$

Example: If you have 100 mL of 0.50 M solution and want 0.10 M, $V_2 = \frac{M_1V_1}{M_2} = \frac{0.50 \times 100}{0.10} = 500$ mL. Add $500 - 100 = 400$ mL water.

Summary Table: Concentration Units

Unit	Definition	Formula
Molarity (M)	Moles of solute per liter of solution	$M = \frac{\text{mol solute}}{\text{L solution}}$
Mole Fraction (X)	Ratio of moles of component to total moles	$X_A = \frac{n_A}{n_{\text{total}}}$
Mass Percent (%)	Mass of solute per total mass, as a percentage	$\frac{\text{mass solute}}{\text{mass solution}} \times 100\%$

Key Takeaways

• Systematic naming and formula writing are foundational skills in chemistry.

• Unit conversions and molar mass calculations are essential for quantitative work.

• Concentration units (molarity, mole fraction, mass percent) are used to describe solutions and are interconvertible.

• Dilution calculations use the $M_1V_1 = M_2V_2$ relationship.

Additional info: For a complete list of polyatomic ions and more practice problems, refer to your course resources or data sheets provided by your instructor.