Introduction
Chemistry studies matter and its transformations. Every substance has physical & chemical properties.
Matter and Its Classification
Matter is anything that has mass and occupies space. It can exist in different states: solid, liquid, gas, and plasma. Matter can be broadly classified into pure substances (elements and compounds) and mixtures (homogeneous and heterogeneous).
| Law | Statement | Example |
|---|---|---|
| Conservation of Mass (Lavoisier) | Matter can neither be created nor destroyed. | C + O₂ → CO₂ (12g + 32g = 44g) |
| Definite Proportions (Proust) | A given compound always contains exactly the same proportion of elements by weight. | H₂O always has H:O in 1:8 mass ratio. |
| Multiple Proportions (Dalton) | If two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in the ratio of small whole numbers. | CO and CO₂. |
| Gay-Lussac's Law of Combining Volumes | When gases combine in a chemical reaction, they do so in a simple ratio by volume provided all gases are at same T and P. | H₂ + Cl₂ → 2HCl |
| Avogadro's Law | Equal volumes of all gases at the same temperature and pressure should contain equal number of molecules. | 1L H₂ and 1L O₂ have same molecules. |
Atomic and Molecular Theory
According to Dalton's Atomic Theory, matter consists of indivisible atoms. All atoms of a given element have identical properties. Compounds are formed when atoms of different elements combine in a fixed ratio. Atoms can neither be created nor destroyed in a chemical reaction.
We measure Atomic mass in amu (atomic mass unit) or u (unified mass). Molecular mass is the sum of atomic masses of the elements present in a molecule. Formula mass is used for ionic compounds.
Mole Concept
The mole concept provides a bridge between the macroscopic and microscopic worlds. A mole is the amount of substance that contains as many particles as there are atoms in exactly 12g of the carbon-12 isotope.
Molar mass ≠ Molecular mass. Molar mass is in g/mol (macroscopic), molecular mass is in amu (atomic scale). For H₂O: mol. mass = 18 amu, molar mass = 18 g/mol.
Avogadro's Number
1 mole = 6.022 × 10²³ particles. This is known as Avogadro's Number (Nₐ). Molar mass is the mass of 1 mole of a substance in grams.
| Element/Compound | Formula | Molar Mass (g/mol) |
|---|---|---|
| Hydrogen | H | 1 |
| Oxygen | O | 16 |
| Water | H₂O | 18 |
| Sodium Chloride | NaCl | 58.5 |
| Carbon Dioxide | CO₂ | 44 |
| Sulfuric Acid | H₂SO₄ | 98 |
Concentration Terms
The concentration of a solution expresses the amount of solute present in a given amount of solvent or solution.
Mass percentage is the mass of solute per 100g of the solution.
Solved Examples
(a) Find moles in 36g of H₂O.
Moles = Given mass / Molar mass = 36g / 18 g/mol = 2 moles.
(b) Find number of molecules in 0.5 mol CO₂.
N = n × Nₐ = 0.5 × 6.022 × 10²³ = 3.011 × 10²³ molecules.
(c) Find molarity of 4g NaOH in 500mL solution.
Molar mass of NaOH = 40 g/mol.
Moles of NaOH = 4 / 40 = 0.1 mol.
Volume = 500 mL = 0.5 L.
Molarity = 0.1 / 0.5 = 0.2 M.