Understanding Chemical Formulas

Formula Notation Introduction

• Formula notation is essential in chemistry for representing substance composition using symbols and numbers.
• Mastering chemical formulas is crucial for balancing chemical equations.
• Also helpful for predicting reactions and solving quantitative problems.

Types of Chemical Formulas

• There are multiple types of chemical formulas
• These include Empirical, Molecular, Structural, Condensed, and Ionic Formulas
• Each displays the arrangement of atoms in unique ways

Empirical Formula

• Represents the simplest whole-number ratio of atoms in a compound.
• Glucose (C₆H₁₂O₆) has an empirical formula of CH₂O.

Molecular Formula

• Shows the actual number of atoms of each element in a molecule.
• Hydrogen peroxide's molecular formula is H₂O₂.

Structural Formula

• Illustrates how atoms are arranged in a molecule.
• Ethanol's structural formula is CH₃CH₂OH.

Condensed Formula

• A compact way to write a structural formula without showing bonds.
• Propane's condensed formula is C₃H₈.

Ionic Formula

• Represents the ratio of ions in an ionic compound.
• Sodium chloride's ionic formula is NaCl, showing a 1:1 ratio of Na⁺ to Cl⁻.

Rules for Writing Chemical Formulas

• Elements are represented by symbols (H for hydrogen, O for oxygen).
• Subscripts indicate the number of atoms (H₂O means 2 hydrogen atoms and 1 oxygen atom).

Ionic Compounds

• The positive ion (cation) is written first, followed by the negative ion (anion).
• Charges must balance out to zero; e.g., Magnesium chloride is MgCl₂ (Mg²⁺ and 2 Cl⁻).

Covalent Compounds

• Prefixes may indicate the number of atoms (CO₂ = carbon dioxide).

Acids

• H is often written first (HCl for hydrochloric acid).

Valency and Oxidation States

• Valency is the combining power of an element.
• Oxidation state indicates the charge of an atom in a compound.
• Common valencies of elements include Sodium (1 in NaCl), Magnesium (2 in MgO), Aluminum (3 in Al₂O₃), Chlorine (1 in HCl), Oxygen (2 in H₂O), Nitrogen (3 in NH₃), and Sulfur (2, 4, 6 in H₂S, SO₂, SO₃).

Criss-Cross Method

• This can be used for Ionic Compounds
• For aluminum oxide (Al³⁺ and O²⁻), swap and drop to get Al₂O₃.

Polyatomic Ions

• Polyatomic ions are ions made up of multiple atoms.
• Ammonium (NH₄⁺) has a charge of +1.
• Hydroxide (OH⁻) has a charge of -1.
• Nitrate (NO₃⁻) has a charge of -1.
• Sulfate (SO₄²⁻) has a charge of -2.
• Carbonate (CO₃²⁻) has a charge of -2.
• Phosphate (PO₄³⁻) has a charge of -3.
• Calcium nitrate is Ca(NO₃)₂ (Ca²⁺ and NO₃⁻).

Formula Notation in Chemical Equations

• Atoms cannot be created or destroyed (Law of Conservation of Mass).
• Steps to balance equations: count atoms, adjust coefficients, and do not change subscripts.
• For the combustion of methane: CH₄ + O₂ → CO₂ + H₂O, the balanced equation is CH₄ + 2O₂ → CO₂ + 2H₂O.

Hydrated and Anhydrous Compounds

• Hydrated compounds contain water molecules in their structure.
• Copper(II) sulfate pentahydrate is CuSO₄·5H₂O.
• Anhydrous compounds have no water molecules.
• Anhydrous copper(II) sulfate is CuSO₄.

Molar Calculations & Formula Determination

• Molar Mass (g/mol) is the sum of atomic masses.
• Use this formula for Percentage Composition: % element = (mass of element in formula / total molar mass) × 100
• To calculate the % of oxygen in H₂O: Molar mass of H₂O = 2(1) + 16 = 18 g/mol, so % O = (16/18) × 100 = 88.89%.

Empirical Formula

• Convert mass % to moles, divide by the smallest number of moles, and write the formula in the simplest whole-number ratio.
• For a compound with 40% C, 6.7% H, 53.3% O, the empirical formula is CH₂O.

Empirical Formula Explained

• The empirical formula represents the simplest whole-number ratio of atoms in a compound.
• Some compounds share the same empirical and molecular formulas (e.g., water, H₂O).
• Glucose (C₆H₁₂O₆) has an empirical formula of CH₂O.
• Hydrogen Peroxide (H₂O₂) has an empirical formula of HO.
• Ethene (C₂H₄) has an empirical formula of CH₂.

Determining the Empirical Formula

• To determine the empirical formula, you should obtain the mass or percentage composition, convert mass to moles, find the simplest ratio, and round to whole numbers.
• Use the formula: Moles = Mass of element / Atomic mass of element (Ar)

Empirical Formula Example Calculation

• For a compound with 2.7 g of Al and 2.4 g of O, the empirical formula is Al₂O₃.
• For a compound with 40% carbon (C), 6.7% hydrogen (H), and 53.3% oxygen (O), the empirical formula is CH₂O.
• Assume a 100 g sample to convert percentages to masses.

Molecular Formula vs. Empirical Formula

• The molecular formula is a multiple of the empirical formula: Molecular Formula = (Empirical Formula)n, where n is a whole number.
• If a compound has an empirical formula of CH₂O and a molecular mass of 180 g/mol, the molecular formula is C₆H₁₂O₆ (Glucose).

Special Cases

• When calculating empirical fomulas, if a calculated value equates to 1.5 or 2.5, multiply by 2 to get whole numbers (e.g., C = 1, H = 1.5 becomes C₂H₃).
• If a calculated value equates to 1.33 or 1.66, multiply by 3 to get whole numbers (e.g., N = 1, O = 1.33 becomes N₃O₄).

Importance of Empirical Formula

• Can be used to determine unknown compounds in experiments.
• Used in stoichiometry (calculations in chemical reactions).
• Essential in industrial chemistry for calculating material composition.