Compounds, Solutions, and Their Composition

Converting Between Grams and Moles

Understanding the relationship between mass and amount of substance is fundamental in chemistry. The mole is the standard unit for amount of substance, and conversions require the compound's molecular mass (also called formula weight).

• Mole: The amount of substance containing Avogadro's number (6.022 × 1023) of particles (atoms, molecules, or ions).

• Molecular Mass (Formula Weight): The sum of the atomic masses of all atoms in a molecule or formula unit, expressed in grams per mole (g/mol).

• Conversion from grams to moles:

  • Use the formula:

• Conversion from moles to grams:

  • Use the formula:

• Example: To find moles in 18 g of H2O (molar mass = 18 g/mol):

Percent Composition of Elements in Compounds

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

• Percent Composition Formula:

• Example: In H2O, percent H = ; percent O =

Empirical and Molecular Formulas

The empirical formula gives the simplest whole-number ratio of elements in a compound, while the molecular formula shows the actual number of each atom in a molecule.

• Calculating Empirical Formula:

  1. Convert masses of elements to moles.

  2. Divide each by the smallest number of moles to get the simplest ratio.

  3. Write the empirical formula using these ratios.

• Converting Empirical Formula to Molecular Formula:

  • Find the empirical formula mass.

  • Divide the compound's molecular mass by the empirical formula mass to get a whole number (n).

  • Molecular formula = (Empirical formula)n

• Example: If empirical formula is CH2O (mass = 30 g/mol) and molecular mass is 180 g/mol, then n = 6, so molecular formula is C6H12O6.

Solutions: Definitions and Properties

Solutions are homogeneous mixtures of two or more substances. Understanding their components and properties is essential for chemical calculations.

• Solute: The substance dissolved in a solution (present in lesser amount).

• Solvent: The substance in which the solute is dissolved (present in greater amount; often water in aqueous solutions).

• Concentration: A measure of the amount of solute in a given quantity of solvent or solution.

• Saturated Solution: A solution containing the maximum amount of solute that can dissolve at a given temperature; additional solute will not dissolve.

Electrolytes and Nonelectrolytes

Substances can be classified based on their ability to conduct electricity when dissolved in water.

• Ionic Substances: Compounds composed of ions (e.g., NaCl) that dissociate into ions in solution and conduct electricity (electrolytes).

• Molecular Substances: Compounds composed of molecules (e.g., sugar, C12H22O11) that generally do not dissociate into ions and do not conduct electricity (nonelectrolytes).

• Electrolyte: A substance that produces ions in solution and conducts electricity.

• Nonelectrolyte: A substance that does not produce ions in solution and does not conduct electricity.

• Example: NaCl (ionic) is an electrolyte; glucose (molecular) is a nonelectrolyte.

Dissociation of Substances in Aqueous Solution

When substances dissolve in water, their behavior depends on their chemical nature.

• Ionic Compounds: Dissociate completely into cations and anions (e.g., NaCl → Na+ + Cl-).

• Strong Electrolytes: Dissociate completely (e.g., strong acids, strong bases, most salts).

• Weak Electrolytes: Partially dissociate (e.g., weak acids like acetic acid).

• Nonelectrolytes: Do not dissociate (e.g., most organic compounds like sugar).

Molarity: Definition and Use

Molarity is the most common unit of concentration in chemistry, expressing the amount of solute per liter of solution.

• Molarity (M):

• Why Use Molarity? It allows chemists to relate the volume of solution to the amount of solute, which is essential for stoichiometric calculations in reactions involving solutions.

• Example: To prepare 1.0 L of 0.5 M NaCl solution, dissolve 0.5 mol NaCl in enough water to make 1.0 L of solution.