Some basic concepts of chemistry
Class 11 · Updated 2026-10-03. Published for practice.
Laws that back up stoichiometry
Conservation of mass: in a closed system, total mass before and after a reaction is the same. Your balanced equation is a statement of that idea in atom counts.
Definite proportions: a compound always has the same element mass ratio (water is always about 11% H by mass). Multiple proportions: when two elements form more than one compound, the mass ratios of one element that combine with a fixed mass of the other are small whole numbers (CO vs CO2 with fixed C).
Gay-Lussac’s law of combining volumes and Avogadro’s hypothesis tie gas volumes at the same temperature and pressure to mole ratios. Those ideas are why 2 H2 + O2 → 2 H2O reads as a 2 : 1 : 2 volume story for gases under the same conditions.
From equation to moles
Coefficients in a balanced equation are mole ratios. Convert given data (mass, volume of gas, concentration × volume) to moles first, then use the ratio to find what you need.
Mass percent in a compound: (mass of element in one formula unit / molar mass of compound) × 100. Empirical formula from percent composition: assume 100 g, convert each element to moles, divide by the smallest mole count, then scale to whole numbers.
If two reactants are both given, find the limiting reagent before you calculate product yield. The excess reactant is whatever is left after the limiting one is used up.
Where this meets the lab
Preparing a solution of known molarity needs molar mass, a balance reading, and volume in litres: moles = M × V(L), mass = moles × molar mass. Rounding and significant figures should match what your teacher expects in the file.
Titration arithmetic in Class 12 still rests on mole ratios from the balanced equation. Getting Unit 1 stoichiometry solid now saves pain later when you combine M1V1/n1 = M2V2/n2 with redox half-reactions.
This note is a map, not a substitute for your NCERT numericals. Work the end-of-chapter problems for unit conversion and limiting reagent until the steps feel automatic.
Common exam slips
Using molecular mass when the question gave empirical data, or forgetting to multiply empirical formula by n when molecular mass is supplied.
Treating dm3 and L as different in mole calculations — they are the same unit. Mixing mL into M × V without converting to litres.
Quoting more significant figures in the answer than the weakest measurement in the question. Match the data, not the calculator display.