Molecules and Compounds

Empirical and Molecular Formulas

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

• Empirical Formula: The simplest ratio of elements in a compound. For example, the empirical formula of hydrogen peroxide (H2O2) is HO.

• Molecular Formula: The actual number of atoms of each element in a molecule. For hydrogen peroxide, the molecular formula is H2O2.

• Relationship: The molecular formula is a whole-number multiple of the empirical formula.

• Calculation: To determine the molecular formula from the empirical formula, use the compound's molar mass:

  • Calculate the molar mass of the empirical formula.

  • Divide the compound's molar mass by the empirical formula mass to find the multiple (n):

  • Multiply the subscripts in the empirical formula by n to get the molecular formula.

• Example: If the empirical formula is CH2O (molar mass = 30 g/mol) and the compound's molar mass is 180 g/mol: Molecular formula = C6H12O6

Naming Compounds

Compounds are named according to systematic rules based on their composition and type (ionic or molecular).

• Ionic Compounds: Consist of a metal and a nonmetal. Name the cation (metal) first, then the anion (nonmetal with '-ide' ending). For transition metals, indicate the charge with Roman numerals. Example: NaCl is sodium chloride; FeCl2 is iron(II) chloride.

• Molecular (Covalent) Compounds: Consist of two nonmetals. Use prefixes to indicate the number of each atom (mono-, di-, tri-, etc.). The first element keeps its name; the second gets '-ide.' Example: CO2 is carbon dioxide; N2O4 is dinitrogen tetroxide.

• Acids: Binary acids use 'hydro-' and '-ic' (e.g., HCl is hydrochloric acid). Oxyacids use '-ic' or '-ous' depending on the polyatomic ion (e.g., H2SO4 is sulfuric acid).

Calculating Average Atomic Mass

The average atomic mass of an element is the weighted average of the masses of its naturally occurring isotopes.

• Formula:

• Example: Chlorine has two isotopes: Cl-35 (mass = 34.97 u, abundance = 75.77%) and Cl-37 (mass = 36.97 u, abundance = 24.23%).

Chemical Reactions and Chemical Quantities

Balancing Chemical Equations

Balancing chemical equations ensures the law of conservation of mass is obeyed: the number of atoms of each element is the same on both sides of the equation.

• Steps to Balance:

  1. Write the unbalanced equation with correct formulas for all reactants and products.

  2. Count the number of atoms of each element on both sides.

  3. Add coefficients to balance one element at a time, starting with the most complex molecule.

  4. Repeat until all elements are balanced.

  5. Check your work to ensure all atoms are balanced and coefficients are in the lowest whole-number ratio.

• Example: Balance the equation for the combustion of methane: Unbalanced: CH4 + O2 → CO2 + H2O Balanced: CH4 + 2 O2 → CO2 + 2 H2O