Some Basic Concepts of Chemistry

• Importance & Nature of Matter

  • Matter: Anything with mass and occupies space
  • Physical Classification
    • Solid: Definite shape and volume
    • Liquid: Definite volume, flexible shape
    • Gas: Neither definite shape nor volume
  • Chemical Classification
    • Pure Substances
      • Elements: Single type of atom (e.g., $\text{Fe}$, $\text{O}_2$)
      • Compounds: Fixed ratio of different atoms (e.g., $\text{H}_2\text{O}$)
    • Mixtures
      • Homogeneous: Uniform composition throughout (e.g., air, solutions)
      • Heterogeneous: Non-uniform composition (e.g., oil in water)

• Properties & Measurement

  • SI Base Units: 7 standard units (kg, m, s, K, mol, A, cd)
  • Uncertainty in Measurement
    • Scientific Notation: Exponential representation ($N \times 10^n$)
    • Precision: Closeness of repeated measurements
    • Accuracy: Agreement of measurement with the true value
    • Significant Figures: Digits known with certainty plus one estimated
    • Dimensional Analysis: Conversion of units using unit factors

• Laws of Chemical Combinations

  • Law of Conservation of Mass: Mass is neither created nor destroyed (Lavoisier)
  • Law of Definite Proportions: Fixed element ratio by mass in compounds (Proust)
  • Law of Multiple Proportions: Simple integer ratios for elements forming multiple compounds (Dalton)
  • Gay Lussac’s Law: Gases react in simple volume ratios at constant T and P
  • Avogadro’s Law: Equal gas volumes contain equal number of molecules at same T and P

• Dalton’s Atomic Theory

  • Matter is made of indivisible atoms
  • Identical atoms for a given element
  • Compounds formed by fixed combination of different atoms
  • Reactions involve rearrangement, not destruction of atoms

• Atomic & Molecular Mass

  • Atomic Mass Unit (amu / u): 1/12th mass of a Carbon-12 atom
  • Average Atomic Mass: Calculated using natural abundance of isotopes
  • Molecular Mass: Sum of atomic masses of all atoms in a molecule
  • Formula Mass: Mass of formula unit for ionic compounds (e.g., $\text{NaCl}$)

• Mole Concept & Molar Mass

  • Mole: $6.022 \times 10^{23}$ entities (Avogadro Constant - $N_A$)
  • Molar Mass: Mass of one mole of a substance in grams
  • Molar Volume: $22.4 \text{ L}$ for any ideal gas at STP

• Empirical & Molecular Formulae

  • Mass Percentage: Mass of element per $100\text{ g}$ of compound
  • Empirical Formula: Simplest whole-number ratio of atoms present
  • Molecular Formula: Actual number of atoms present in a molecule
  • Formula Relation: $\text{Molecular Formula} = n \times \text{Empirical Formula}$

• Stoichiometry & Reaction Calculations

  • Stoichiometric Coefficients: Numbers indicating mole/molecule ratios in balanced equations
  • Limiting Reagent: Reactant consumed completely, limiting product formation
  • Reactions in Solutions
    • Mass Percent: $(\text{Mass of Solute} / \text{Total Mass}) \times 100$
    • Mole Fraction ($x$): Ratio of moles of a component to total moles
    • Molarity ($M$): Moles of solute per Liter of solution ($\text{mol/L}$)
    • Molality ($m$): Moles of solute per Kilogram of solvent ($\text{mol/kg}$)