Chemical Formulas & Equations
Learn to read and write chemical formulas, understand what a balanced chemical equation means, balance simple reactions step by step, and classify the four main types of chemical reactions.
Reading Chemical Formulas
A chemical formula uses element symbols and subscript numbers to show the type and number of atoms in a molecule or compound. The subscript (small number after the symbol) tells you how many atoms of that element are present.
Anatomy of a Chemical Formula
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Common Chemical Formulas to Know
| Substance | Formula | Atoms Present |
|---|---|---|
| Water | H₂O | 2 hydrogen, 1 oxygen |
| Carbon dioxide | CO₂ | 1 carbon, 2 oxygen |
| Oxygen gas | O₂ | 2 oxygen (diatomic molecule) |
| Sodium chloride (table salt) | NaCl | 1 sodium, 1 chlorine |
| Glucose (sugar) | C₆H₁₂O₆ | 6 carbon, 12 hydrogen, 6 oxygen |
| Ammonia | NH₃ | 1 nitrogen, 3 hydrogen |
| Calcium carbonate | CaCO₃ | 1 calcium, 1 carbon, 3 oxygen |
Diatomic elements: Some elements exist as pairs of atoms in their natural state. Remember: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂. A useful mnemonic: "Have No Fear Of Ice Cold Beer" (H, N, F, O, I, Cl, Br). These diatomic forms must be used in balanced equations.
Balancing Chemical Equations
The Law of Conservation of Mass states that matter cannot be created or destroyed in a chemical reaction. The number of atoms of each element must be the same on both sides of a balanced equation. We balance equations by adding coefficients (numbers in front of formulas) — we never change subscripts.
General Format
Substances before the arrow react to form new substances after the arrow.
Step-by-Step Method for Balancing
Write the unbalanced equation with correct formulas for all reactants and products.
Count atoms of each element on each side of the arrow.
Add coefficients (whole numbers in front of formulas) to balance elements that are unequal. Balance one element at a time. Balance metals and non-metals before hydrogen and oxygen.
Check: Recount all atoms on both sides. Confirm the equation is fully balanced.
Types of Chemical Reactions
Chemical reactions can be grouped into four main types based on how the atoms rearrange.
1. Synthesis (Combination)
Two or more substances combine to form a single product.
Example: 2Mg + O₂ → 2MgO (magnesium burns in oxygen to form magnesium oxide)
2. Decomposition
A single compound breaks down into two or more simpler substances (often when heated).
Example: 2H₂O → 2H₂ + O₂ (electrolysis of water) or CaCO₃ → CaO + CO₂ (heating limestone)
3. Single Displacement (Replacement)
A more reactive element displaces a less reactive one from a compound.
Example: Zn + 2HCl → ZnCl₂ + H₂ (zinc displaces hydrogen from hydrochloric acid)
4. Double Displacement
Two compounds swap ions to form two new compounds. Often produces a precipitate, gas, or water.
Example: NaCl + AgNO₃ → AgCl↓ + NaNO₃ (silver chloride precipitate forms)
Combustion is a special type often studied separately: a fuel reacts rapidly with oxygen to produce CO₂ and H₂O, releasing heat and light. Example: CH₄ + 2O₂ → CO₂ + 2H₂O (methane burning in oxygen). Combustion reactions are a type of synthesis/oxidation.
Key Vocabulary
Coefficient
A number placed in front of a chemical formula in an equation to balance it. A coefficient multiplies all atoms in that formula. Changing coefficients does NOT change the identity of the substance.
Reactant
A substance that is present at the start of a chemical reaction and is consumed (used up) during the reaction. Reactants are written on the left side of the arrow in a chemical equation.
Product
A substance formed as a result of a chemical reaction. Products are written on the right side of the arrow. New substances with different properties to the reactants are produced.
Conservation of Mass
The law stating that mass is neither created nor destroyed in a chemical reaction. The total mass of reactants equals the total mass of products. This is why equations must be balanced.
Worked Examples
Balance the equation: H₂ + O₂ → H₂O
Step 1: Count atoms (unbalanced)
| Element | Reactants | Products | Balanced? |
|---|---|---|---|
| H | 2 | 2 | Yes |
| O | 2 | 1 | No |
Step 2: Put 2 in front of H₂O: H₂ + O₂ → 2H₂O. Now O is balanced (2=2) but H isn't (2≠4).
Step 3: Put 2 in front of H₂: 2H₂ + O₂ → 2H₂O
Check: H: 4=4 ✓, O: 2=2 ✓. Balanced equation: 2H₂ + O₂ → 2H₂O
Balance: Fe + O₂ → Fe₂O₃ (iron rusting)
Unbalanced: Fe: 1 vs 2 (need 2 Fe); O: 2 vs 3 (harder — find LCM)
Balance Fe: Put 4 in front of Fe and 2 in front of Fe₂O₃: 4Fe + O₂ → 2Fe₂O₃
Count O: Products now have 6 oxygen. Put 3 in front of O₂: 4Fe + 3O₂ → 2Fe₂O₃
Check: Fe: 4=4 ✓, O: 6=6 ✓. Balanced: 4Fe + 3O₂ → 2Fe₂O₃
Identify the reaction type: 2KClO₃ → 2KCl + 3O₂
Observation: One substance (potassium chlorate, KClO₃) breaks down into two substances (KCl and O₂).
Pattern matches: AB → A + B
Reaction type: Decomposition. This is the classic lab method for producing oxygen gas — heating potassium chlorate with a manganese dioxide catalyst.
Knowledge Check
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Key Concepts Summary
- ✓ Chemical formulas show the type and number of atoms in a substance. Subscripts show number of atoms of each element.
- ✓ The Law of Conservation of Mass requires equations to be balanced — same atoms on both sides.
- ✓ Balance equations by adjusting coefficients only — never change subscripts in formulas.
- ✓ Synthesis: A + B → AB. Decomposition: AB → A + B. Single displacement: A + BC → AC + B. Double displacement: AB + CD → AD + CB.
- ✓ Seven elements exist as diatomic molecules in their natural form: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂.
- ✓ Combustion reactions involve a fuel reacting with oxygen to produce CO₂ and H₂O, releasing energy.