AIEM — ICSE Exam Prep

Some Basic Concepts of Chemistry - Quick Revision

Law of Conservation of Mass (Lavoisier, 1789): in all physical and chemical changes there is no net change in mass; matter can neither be created nor destroyed.
Law of Definite Proportions (Proust): a given compound always contains exactly the same proportion of elements by mass, irrespective of source. Also called Law of Definite Composition.
Law of Multiple Proportions (Dalton, 1803): if two elements form more than one compound, the masses of one element combining with a fixed mass of the other are in the ratio of small whole numbers (e.g. O in water vs hydrogen peroxide = 16:32 = 1:2).
Gay Lussac's Law of Gaseous Volumes (1808): gases combine in simple volume ratios at the same T and P (H2:O2 = 100 mL:50 mL = 2:1).
Avogadro's Law (1811): equal volumes of all gases at the same T and P contain equal numbers of molecules.

Matter consists of indivisible atoms. 2. All atoms of an element have identical mass and properties; atoms of different elements differ. 3. Compounds form when atoms of different elements combine in a fixed ratio. 4. Chemical reactions are reorganisation of atoms; atoms are neither created nor destroyed. It explained the laws of chemical combination but could NOT explain the laws of gaseous volumes.

Standard since 1961: carbon-12 (12C) = exactly 12 amu; 1 amu = one-twelfth the mass of a 12C atom = 1.66056 x 10^-24 g. 'amu' is now written u (unified mass).
Average atomic mass weights isotope masses by relative abundance (carbon = 12.011 u). Periodic-table masses are average atomic masses.
Molecular mass = sum of atomic masses in a molecule (CH4 = 16.043 u; H2O = 18.02 u).
Formula mass used for ionic solids with no discrete molecules (NaCl = 23.0 + 35.5 = 58.5 u).

One mole = exactly 6.02214076 x 10^23 elementary entities = the Avogadro number (NA). An entity may be an atom, molecule, ion, electron, etc.
Molar mass = mass of one mole in grams; numerically equal to atomic/molecular/formula mass in u (water = 18.02 g mol^-1; NaCl = 58.5 g mol^-1).

Mass % of an element = (mass of element in compound / molar mass of compound) x 100.
Empirical formula = simplest whole-number ratio of atoms; molecular formula = exact number of each atom. Molecular formula = empirical formula x n, where n = molar mass / empirical formula mass.

Stoichiometry (Greek stoicheion = element, metron = measure) deals with masses/volumes of reactants and products from a balanced equation. Coefficients are stoichiometric coefficients (mole ratios); subscripts cannot be changed when balancing.
Limiting reagent: the reactant present in least amount, consumed first, limiting the product formed.
Concentration terms: mass per cent (w/w%), mole fraction, molarity (mol/L solution, temperature dependent), molality (mol/kg solvent, temperature independent).

SI base units (Table 1.1): length-metre, mass-kilogram, time-second, electric current-ampere, temperature-kelvin, amount of substance-mole, luminous intensity-candela.
Scientific notation: N x 10^n with 1 <= N < 10.
Significant figures convey uncertainty; rules govern zeros, rounding off, and operations.
Precision = closeness of repeated measurements; accuracy = agreement with the true value.