Formulae

The empirical formula represents the simplest whole-number ratio of the elements present in a compound. It is the most basic form of a chemical formula and is often determined through experimental data such as mass or percentage composition.

The molecular formula provides the actual number of atoms of each element in a single molecule of a substance. It is always a whole-number multiple of the empirical formula (e.g., if the empirical formula is $CH_2$, the molecular formula could be $C_2H_4$ or $C_3H_6$).

To derive a molecular formula, the relative formula mass ($M_r$) of the actual compound is divided by the $M_r$ of the empirical formula to find a multiplier, which is then applied to all subscripts in the empirical formula.

A balanced equation ensures that the number of atoms for each element is identical on both the reactant and product sides, satisfying the Law of Conservation of Mass. Matter cannot be created or destroyed in a chemical reaction.

Coefficients are the large numbers placed in front of chemical formulae to adjust the quantity of molecules or formula units. Subscripts within a formula must never be changed when balancing, as this would change the identity of the substance.

State symbols provide essential context about the physical form of reactants and products: $(s)$ for solid, $(l)$ for liquid, $(g)$ for gas, and $(aq)$ for aqueous (dissolved in water).

Ionic equations focus only on the species that actually participate in a chemical change. They are derived by breaking down aqueous ionic compounds into their constituent ions.

Spectator ions are ions that appear on both sides of a chemical equation in the same state and form. Because they do not participate in the reaction, they are cancelled out to produce the net ionic equation.

This method is particularly useful for precipitation and neutralization reactions, highlighting the core chemical interaction, such as the formation of a solid from two aqueous solutions.

Step 1: Determine Mass: Identify the mass of each element in a sample. If given percentages, assume a 100g sample so that percentages translate directly to grams.

Step 2: Convert to Moles: Divide the mass of each element by its relative atomic mass ($A_r$). This converts the macroscopic mass into a count of atoms (moles).

Step 3: Find the Ratio: Divide each molar value by the smallest molar value obtained in Step 2. If the results are not whole numbers, multiply all values by a common factor to achieve a whole-number ratio.

Check Charge Balance: When writing formulae for ionic compounds, always verify that the sum of positive and negative charges equals zero. For transition metals, use the Roman numeral to identify the specific oxidation state.

The 'Brackets' Rule: Always use brackets for polyatomic ions if the subscript is greater than one. For example, Calcium Hydroxide is $Ca(OH)_2$, not $CaOH_2$.

Sanity Check Equations: After balancing an equation, perform a final count of every atom on both sides. A common mistake is forgetting to multiply the coefficient by the subscript (e.g., $2H_2O$ contains 4 Hydrogen atoms).

Simplifying Ratios: Ensure empirical formulae are truly in their simplest form. For example, $C_2H_4$ must be simplified to $CH_2$ if the question asks for the empirical formula.