Some Basic Concepts of Chemistry

2.1 Core concepts

Chemistry is the branch of science that studies the preparation, properties, structure and reactions of material substances. It underpins the manufacture of fertilisers, alkali, acids, salts, dyes, polymers, drugs (cisplatin and taxol in cancer therapy, AZT in HIV control), soaps, detergents and metals, and it is central to weather prediction, ozone-layer protection, and pollution control (NCERT §1.1, p. 4).

Nature and states of matter. Matter is anything that has mass and occupies space. At the macroscopic level it occurs in three physical states — solid (definite volume, definite shape, particles close-packed in fixed positions), liquid (definite volume, no definite shape, particles close but mobile), and gas (no definite volume or shape, particles far apart and moving freely) — and these states are interconvertible by changing temperature or pressure (NCERT §1.2.1, p. 4–5; Fig. 1.1).

Classification of matter. Matter is first classified at the bulk level into pure substances (fixed composition, cannot be separated by physical methods) and mixtures (variable composition, separable by physical methods). Mixtures are homogeneous (sugar in water, air, brass — uniform throughout) or heterogeneous (sugar + salt, grains + pulses + dirt — non-uniform). Pure substances split further into elements (one type of atom — Na, Cu, H₂, O₂) and compounds (two or more elements combined in a fixed ratio — H₂O, CO₂, sugar); a compound's properties differ from those of the elements that make it (NCERT §1.2.2, p. 5–6; Figs. 1.2–1.4).

Properties of matter and SI units. Properties are physical (colour, melting point, boiling point, density — measured without changing chemical identity) or chemical (combustibility, acidity, reactivity — require a chemical change) (NCERT §1.3.1, p. 6). The International System of Units (SI), adopted by the 11th CGPM in 1960 and redefined in 2019, has seven base units — metre (length), kilogram (mass), second (time), ampere (electric current), kelvin (thermodynamic temperature), mole (amount of substance), and candela (luminous intensity) — from which all derived units are built (NCERT §1.3.3, Table 1.1, p. 7). Mass is the amount of matter present and is invariant; weight is the gravitational force on a body and changes with location. Volume has units of (length)³; 1 L = 1000 mL = 1000 cm³ = 1 dm³ (NCERT §1.3.4–1.3.5, p. 9). Density (mass per unit volume) has SI unit kg m⁻³, although chemists routinely use g cm⁻³. The Celsius scale is set by °C = (°F − 32) × 5/9, and the Kelvin scale by K = °C + 273.15; negative readings exist on °C but not on K (NCERT §1.3.6–1.3.7, p. 10; Fig. 1.8).

Uncertainty in measurement. Scientific notation expresses every quantity as N × 10ⁿ with 1 ≤ N < 10, so that very large (6.022 × 10²³) and very small (1.66 × 10⁻²⁴ g) numbers fit comfortably on the page and in calculators (NCERT §1.4.1, p. 11). The rules for significant figures are: all non-zero digits are significant; captive zeros between non-zero digits are significant; leading zeros are not; trailing zeros are significant only if a decimal point is present (so 100 has 1 sig fig but 100. has 3 and 100.0 has 4); counted exact numbers have infinite sig figs (NCERT §1.4.2, p. 12). In addition/subtraction, the answer is limited by the fewest decimal places; in multiplication/division, by the fewest significant figures. Precision is the closeness of repeated readings to one another; accuracy is the closeness of a reading to the true value (Table 1.4, p. 13). Dimensional analysis (factor-label method) uses unit-conversion factors equal to unity (e.g. 2.54 cm/1 in) to interconvert units cleanly (NCERT §1.4.3, p. 13–14).

Laws of chemical combination. (i) Conservation of Mass (Lavoisier, 1789) — matter is neither created nor destroyed in any physical or chemical change. (ii) Definite Proportions (Proust) — a given compound always contains the same elements in the same proportion by mass irrespective of source or method of preparation. (iii) Multiple Proportions (Dalton, 1803) — when two elements form more than one compound, the masses of one combining with a fixed mass of the other bear a small whole-number ratio (H in H₂O vs H₂O₂ → O masses 16 : 32 = 1 : 2). (iv) Gay-Lussac's Law of Gaseous Volumes (1808) — gases combine or are produced in simple volume ratios at the same T and P (H₂ + ½O₂ → H₂O is volumetrically 2 : 1 : 2). (v) Avogadro's Law (1811) — equal volumes of all gases at the same T and P contain equal numbers of molecules (NCERT §1.5, p. 14–16; Fig. 1.9).

Dalton's atomic theory (1808) crystallises these laws: matter is made of indivisible atoms; atoms of an element are identical in mass and properties; compounds form when atoms combine in fixed ratios; chemical reactions only rearrange atoms — never create or destroy them (NCERT §1.6, p. 16).

Atomic, molecular and formula mass. Atomic mass is given relative to ¹²C, which is assigned exactly 12 u. One unified atomic mass unit (u or amu) = 1.66056 × 10⁻²⁴ g. The average atomic mass weights isotopic masses by natural abundance (C = 12.011 u, blending ¹²C, ¹³C and trace ¹⁴C). Molecular mass is the sum of atomic masses (H₂O = 18.02 u, CH₄ = 16.043 u). For ionic substances without discrete molecules (NaCl), the analogous quantity is called formula mass (NaCl = 58.5 u) (NCERT §1.7, p. 16–18; Fig. 1.10).

The mole. One mole is the SI amount of substance containing exactly 6.02214076 × 10²³ elementary entities — atoms, molecules, ions, electrons, or specified groups — and this fixed value defines the Avogadro constant N_A. The molar mass in g mol⁻¹ is numerically equal to the atomic/molecular/formula mass in u (NCERT §1.8, p. 18; Fig. 1.11). Mole↔mass↔particle interconversion uses molar mass and N_A.

Percentage composition, empirical and molecular formulas. Mass per cent of an element = (mass of that element in 1 mol / molar mass) × 100. For ethanol C₂H₅OH: C = 52.14 %, H = 13.13 %, O = 34.73 %. From percentage composition: convert to moles → divide by smallest mole value → multiply to whole numbers → that ratio is the empirical formula. The molecular formula = (empirical formula) × n, where n = (molar mass) / (empirical formula mass) (NCERT §1.9–1.9.1, p. 18–20).

Stoichiometry and limiting reagent. A balanced equation gives molar (and hence mass and volume) ratios. For CH₄ + 2 O₂ → CO₂ + 2 H₂O, combustion of 16 g CH₄ yields 36 g H₂O and 44 g CO₂ (Problem 1.3, p. 20–22). The limiting reagent is the reactant present in the least stoichiometric quantity — it is consumed first and decides the yield. Problem 1.5 (p. 22): mixing 50.0 kg N₂ with 10.0 kg H₂ to make NH₃ — moles needed of H₂ (5358) exceed moles available (4960), so H₂ is limiting; NH₃ formed = 3.30 × 10³ mol = 56.1 kg.

Concentration of solutions. Four common expressions — mass per cent (mass of solute × 100/mass of solution), mole fraction (xₐ = nₐ/(nₐ+n_b)), molarity M = moles of solute / volume of solution in L, and molality m = moles of solute / mass of solvent in kg. The dilution shortcut M₁V₁ = M₂V₂ follows from conservation of moles. Molarity changes with temperature because the solution volume expands; molality does not, because the solvent mass is invariant (NCERT §1.10.2, p. 21–24).

2.2 Definitions to memorise

2.3 Diagrams / processes to remember

Make sure you can sketch and label each diagram below. Fig. 1.1 (p. 5) shows the three states with their characteristic particle arrangement: solids close-packed and ordered, liquids close but disordered, gases far apart and rapidly moving. Fig. 1.2 (p. 5) is the master classification flowchart — Matter → Pure substance (Elements/Compounds) or Mixture (Homogeneous/Heterogeneous); CUET frequently tests where a given example (air, milk, brass, sodium chloride solution) lands on this tree. Fig. 1.3 (p. 6) contrasts atoms (Na, Cu) with diatomic molecules (H₂, N₂, O₂), and Fig. 1.4 depicts H₂O and CO₂ molecules — useful for distinguishing elemental versus compound molecules.

Fig. 1.6 (p. 9) is a cube illustrating volume relationships: 1 L = 1000 mL = 1000 cm³ = 1 dm³. Fig. 1.7 (p. 10) shows the four volume-measuring devices a CUET candidate must be able to identify by name and by precision class — graduated cylinder, burette, pipette, volumetric flask. Fig. 1.8 (p. 10) puts °C, °F and K side by side so that you can compare the freezing point of water (0 °C / 32 °F / 273.15 K) and human body temperature (37 °C / 98.6 °F / 310.15 K).

Fig. 1.9 (p. 16) is the central visual of the laws of chemical combination: two volumes of H₂ + one volume of O₂ → two volumes of water vapour, with little boxes showing equal numbers of molecules per equal volumes (Avogadro's bridge from Gay-Lussac). Fig. 1.10 (p. 17) depicts cubic packing of Na⁺ and Cl⁻ in NaCl — each ion surrounded by six counter-ions — explaining why NaCl has a formula mass rather than a molecular mass. Fig. 1.11 (p. 18) lays out one mole of various substances (12 g C, 18 g H₂O, 23 g Na, 32 g S, 56 g Fe, 63.5 g Cu, 200 g Hg) side by side to drive home that "one mole" denotes a count of particles, not a fixed mass.

Three workflow processes are worth memorising. The mole-conversion tree (p. 18) maps mass ↔ moles ↔ number of particles via molar mass and N_A, and mass ↔ volume via density — this is the spine of every CUET numerical. The empirical → molecular formula process (p. 19–20) takes percentage composition → divide by atomic mass to get moles → divide by the smallest mole to get the empirical ratio → multiply by n = M_molar/M_empirical to get the molecular formula. The stoichiometry-with-limiting-reagent process (p. 20–22) is: convert all masses to moles → compare available moles to stoichiometric moles required → identify the limiting reagent → use it to compute product moles → convert back to mass or volume. Practising these three procedural pipelines once each will solve >70 % of CUET kech101 numericals.

2.4 Common confusions / NTA trap points

• Mass vs weight — students treat them as synonyms; NCERT explicitly states (p. 9) that mass is constant while weight varies with gravitational acceleration. Distractor: "weight is the amount of matter".
• Significant figures with terminal zeros — 100 has only 1 sig fig, 100. has 3, 100.0 has 4 (p. 12). NTA loves this trap; the decimal point is decisive.
• Definite vs Multiple Proportions — "Definite" is about ONE compound (always the same composition); "Multiple" is about TWO or more compounds of the same two elements (whole-number mass ratios). Don't confuse the two laws.
• Gay-Lussac vs Avogadro — Gay-Lussac talks about combining volumes; Avogadro extends it to equal molecules per equal volume. They are different statements, and CUET MCQs deliberately swap the names.
• Molarity vs molality — molarity depends on solution volume and so varies with temperature; molality uses solvent mass and so is temperature-independent (p. 24).
• Limiting-reagent shortcut — compare moles available to moles required by stoichiometry; the reactant that yields the fewer moles of product is limiting. Smaller mass alone is NOT the criterion (Problem 1.5).
• amu definition — exactly 1/12 the mass of one ¹²C atom (p. 16); not "the mass of a hydrogen atom" (that is the older 1803 system).
• Empirical vs molecular formula — glucose has empirical formula CH₂O but molecular formula C₆H₁₂O₆. CUET asks you to find n = molar mass / empirical formula mass; don't forget to multiply through.
• Mole fraction has no units — it is a ratio of moles, so its sum over all components must equal 1. Trap: choosing % composition as a "mole fraction".
• Avogadro number is now exact — since 2019 it is fixed at 6.02214076 × 10²³ mol⁻¹ by definition, not measured (p. 18). NTA may phrase this as "approximate" in a distractor.
• Counting particles in diatomic molecules — 1 mole of Cl₂ contains 6.022 × 10²³ molecules but 2 × 6.022 × 10²³ Cl atoms. Watch the wording.
• Density-percentage-to-molarity — M = (10 × density × % w/w) / molar mass is a derived shortcut for problems where concentrated solutions are described by weight percent and density (Exercise 1.38, p. 29).

2.5 Key reactions & formulas