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Schooling · Class 11 · Chemistry · Some Basic Concepts of Chemistry · Chapter on a page

CBSE Class 11 · Chemistry · NCERT Chemistry Part-I

Some Basic Concepts of Chemistry: chapter on a page

A cut of Shishya's notes for this chapter: the first paragraph of the introduction and the key ideas, formulas and facts, common mistakes and quick revision lists, word for word. Nothing is added or reworded; the worked examples are left out (they are on the chapter page).

About the notes: Written by Shishya's AI on 27 Sept 2026 from the chapter's title and class level, in Shishya's own words — not a copy or summary of the textbook. Read the official chapter for the book's own text, activities and exercises.

Fact-checked on 21 Oct 2026: an AI-assisted check of this note against cited sources; the errors it found were corrected that day. Some errors may remain.

Mind map

Drawn from the same notes by a fixed rule (no AI): each label is the item's own words — its bold term, its label before a colon, a short formula as written, or its first words.

Mind map of Some Basic Concepts of Chemistry: Key ideas, Formulas and facts, Watch out forKey ideasLaw of conservation of massLaw of definite proportionsLaw of multiple proportionsMole conceptMolar mass and molecular massStoichiometryLimiting reagentFormulas and factsNumber of molesNumber of particlesMolar volume of an ideal gas at STPPercentage compositionMolarityMolalityMass percentWatch out forAdding atomic masses of only the…Using mass ratio directly as mole…Reporting an answer with more…Confusing molar mass (g mol⁻¹) with…Forgetting to identify the limiting…Some Basic Conceptsof Chemistry
The mind map as a list
  • Key ideas
    • Law of conservation of mass
    • Law of definite proportions
    • Law of multiple proportions
    • Mole concept
    • Molar mass and molecular mass
    • Stoichiometry
    • Limiting reagent
  • Formulas and facts
    • Number of moles
    • Number of particles
    • Molar volume of an ideal gas at STP
    • Percentage composition
    • Molarity
    • Molality
    • Mass percent
  • Watch out for
    • Adding atomic masses of only the…
    • Using mass ratio directly as mole…
    • Reporting an answer with more…
    • Confusing molar mass (g mol⁻¹) with…
    • Forgetting to identify the limiting…

What this chapter is about

This chapter lays the foundation for all the chemistry you will study in senior secondary school. It begins with the nature of matter—how substances exist as elements, compounds and mixtures—and moves to the laws that govern chemical combinations. You learn to express amount of substance through the mole, connect laboratory masses to the number of atoms or molecules, and use balanced equations to predict how much product forms from given reactants.

Key ideas

  • Law of conservation of mass: In a chemical reaction, total mass of reactants equals total mass of products; matter is neither created nor destroyed.
  • Law of definite proportions: A pure compound always contains its constituent elements in a fixed ratio by mass, regardless of source.
  • Law of multiple proportions: When two elements form more than one compound, the different masses of one element that combine with a fixed mass of the other are in a ratio of small whole numbers.
  • Mole concept: One mole of any substance contains exactly 6.02214076 × 10²³ (≈ 6.022 × 10²³) entities (Avogadro constant, Nₐ). The molar mass (g mol⁻¹) equals the relative atomic or molecular mass expressed in grams.
  • Molar mass and molecular mass: Molecular mass is the sum of atomic masses of all atoms in a formula; molar mass is the mass of one mole of that substance.
  • Stoichiometry: A balanced equation gives the mole ratios in which reactants combine and products form, enabling mass and volume calculations.
  • Limiting reagent: The reactant that is completely consumed first limits the amount of product; any other reactant present in excess remains unreacted.
  • Empirical and molecular formulas: The empirical formula shows the simplest whole-number ratio of atoms; the molecular formula shows the actual number of atoms in one molecule.

Formulas and facts to remember

  • Number of moles: n = m / M, where m is mass in grams, M is molar mass in g mol⁻¹.
  • Number of particles: N = n × Nₐ, with Nₐ = 6.022 × 10²³ mol⁻¹.
  • Molar volume of an ideal gas at STP (0 °C, 1 bar): 22.7 L mol⁻¹ (or 22.4 L mol⁻¹ at older STP of 1 atm).
  • Percentage composition: (mass of element in 1 mol / molar mass of compound) × 100.
  • Molarity (M): moles of solute per litre of solution, unit mol L⁻¹.
  • Molality (m): moles of solute per kilogram of solvent, unit mol kg⁻¹.
  • Mass percent: (mass of solute / mass of solution) × 100.
  • Significant figures rule for multiplication/division: Result keeps as many significant figures as the quantity with the fewest.

Common mistakes

  • Adding atomic masses of only the unique elements and ignoring subscripts → count every atom, multiplied by its subscript.
  • Using mass ratio directly as mole ratio in stoichiometry → convert masses to moles first, then apply the balanced equation's coefficients.
  • Reporting an answer with more significant figures than the data justify → use the fewest significant figures among the measured quantities.
  • Confusing molar mass (g mol⁻¹) with molecular mass (u) → numerically equal, but units differ; use g mol⁻¹ for lab calculations.
  • Forgetting to identify the limiting reagent → always check which reactant runs out first before finding product mass.

Quick revision

  • One mole = 6.022 × 10²³ particles; molar mass in grams equals relative atomic/molecular mass.
  • n = m / M links mass to moles; N = n × Nₐ links moles to particle count.
  • Balance the equation first; use mole ratios, not mass ratios, for stoichiometry.
  • Limiting reagent decides maximum product; excess reagent remains.
  • Empirical formula: simplest whole-number atom ratio; molecular formula may be a multiple of it.
  • Keep significant figures consistent with the precision of your measurements.