Learning Objectives
- Understand the importance of chemistry and its scope
- Learn about the laws of chemical combination
- Understand Dalton's atomic theory and the mole concept
- Perform stoichiometric calculations and determine empirical/molecular formulas
- Understand concentration terms and limiting reagent concept
Key Concepts
Laws of Chemical Combination
Law of Conservation of Mass (Lavoisier): Matter can neither be created nor destroyed in a chemical reaction. Total mass of reactants = Total mass of products.
Law of Definite Proportions (Proust): A given compound always contains the same elements in the same fixed proportion by mass.
Law of Multiple Proportions (Dalton): When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in simple whole number ratios.
Gay-Lussac's Law of Gaseous Volumes: Gases combine in simple whole number ratios by volume at same T and P.
Avogadro's Law: Equal volumes of all gases at same T and P contain the same number of molecules.
Dalton's Atomic Theory
Matter consists of indivisible atoms. Atoms of the same element are identical. Atoms combine in simple whole number ratios. Atoms cannot be created or destroyed.
Atomic and Molecular Masses
Atomic mass unit (u): 1 u = 1/12 × mass of ¹²C = 1.66054 × 10⁻²⁴ g.
Molecular mass: Sum of atomic masses of all atoms in a molecule.
Formula mass: Used for ionic compounds (e.g., NaCl: 23 + 35.5 = 58.5 u).
Mole Concept
1 mole = 6.022 × 10²³ particles (Avogadro's number, Nₐ).
Molar mass: Mass of 1 mole of substance in grams = numerical value of atomic/molecular mass in u.
Number of moles: n = given mass / molar mass = number of particles / Nₐ = volume at STP / 22.4 L
At STP (0°C, 1 atm): 1 mole of any gas occupies 22.4 L.
Percentage Composition and Empirical/Molecular Formula
% of element = (mass of element in 1 mole / molar mass) × 100
Empirical formula: Simplest whole number ratio of atoms. Steps: % → mass → moles → simplest ratio.
Molecular formula: n × empirical formula, where n = molecular mass / empirical formula mass.
Stoichiometry and Limiting Reagent
Balanced equations show molar ratios. The limiting reagent is the reactant that is completely consumed first and determines the amount of product formed.
Percentage yield = (actual yield / theoretical yield) × 100%
Concentration Terms
Molarity (M): moles of solute / litres of solution
Molality (m): moles of solute / kg of solvent
Mole fraction: xₐ = nₐ / (nₐ + n_b); xₐ + x_b = 1
Mass percentage: (mass of solute / mass of solution) × 100
Parts per million (ppm): (mass of solute / mass of solution) × 10⁶
Summary
Chemistry relies on fundamental laws of chemical combination. The mole concept connects the microscopic world of atoms to measurable quantities. Stoichiometric calculations use balanced equations and molar ratios. The limiting reagent determines the maximum product yield. Various concentration terms express solution composition for different applications.
Important Terms
- Mole: Amount of substance containing Nₐ = 6.022 × 10²³ particles
- Molar Mass: Mass of one mole in grams (g/mol)
- Avogadro's Number: 6.022 × 10²³ mol⁻¹
- Empirical Formula: Simplest ratio of atoms in a compound
- Limiting Reagent: Reactant completely consumed first
- Molarity: Moles of solute per litre of solution
Quick Revision
- 1 u = 1.66054 × 10⁻²⁴ g; 1 mole = 6.022 × 10²³ particles
- n = mass/molar mass = particles/Nₐ = volume(STP)/22.4 L
- Empirical formula: % → mass → moles → ratio
- Molecular formula = n × empirical formula
- Molarity = mol/L; Molality = mol/kg solvent
- Limiting reagent gives least moles of product