Some Basic Concepts of Chemistry — NEET Notes

Nature of matter

Matter is anything that has mass and occupies space. NCERT classifies it on two parallel axes. At the physical level, matter exists in three states — solid, liquid, gas — distinguished by how tightly the constituent particles are held. Solids have a definite shape and volume, liquids have a definite volume but take the shape of the container, and gases occupy the entire space available. These three states are interconvertible by changing temperature or pressure: heating a solid melts it to a liquid and then vaporises it to a gas; cooling reverses the sequence.

At the chemical level, matter is either a pure substance or a mixture. A pure substance has constituent particles of fixed chemical composition — copper, silver, water, glucose. A mixture contains two or more pure substances in any ratio; its composition is variable. Mixtures are homogeneous (sugar solution, air) when the components are uniformly distributed, and heterogeneous (salt + sugar, grain + dirt) when they are visibly distinct. Components of a mixture can be separated by physical methods such as filtration, distillation, or crystallisation; components of a compound cannot.

Pure substances are further divided into elements (all atoms of one kind — sodium, oxygen, copper) and compounds (atoms of two or more elements combined in a fixed ratio — water, ammonia, carbon dioxide). The properties of a compound are completely different from those of its constituent elements. Hydrogen burns and oxygen supports combustion, but their compound — water — is used to extinguish fire. This independence of compound properties from elemental properties is one of the central observations that the atomic theory had to explain.

Properties of matter & measurement

Every substance has characteristic properties, sorted into two categories. Physical properties — colour, odour, melting point, boiling point, density — can be measured or observed without altering the identity of the substance. Chemical properties — combustibility, acidity, reactivity with acids and bases — require a chemical change to be observed. Quantitative measurement is the engine of chemistry, and every measurement is reported as a number followed by a unit.

The International System of Units (SI), established by the 11th CGPM in 1960, defines seven base units that cover every measurable quantity. Each unit was redefined in 2019 in terms of fixed numerical values of fundamental physical constants — making the SI a system grounded entirely in nature rather than in artefacts.

The chemistry laboratory uses derived units of everyday convenience — gram for mass (1 kg = 1000 g), cm³ or dm³ for volume, °C for temperature, g cm⁻³ for density. The litre (L) — a non-SI unit equal to 1 dm³ or 1000 cm³ — remains in widespread use for liquid volume. Mass is the amount of matter in a body and is invariant; weight is the gravitational force on that mass and varies from place to place. The two are routinely confused and NCERT is unusually firm about distinguishing them.

Two temperature conversions cover every NEET problem. The Fahrenheit–Celsius relation is °F = (9/5) × °C + 32. The Celsius–Kelvin relation is K = °C + 273.15. Density is mass per unit volume — the SI unit is kg m⁻³, but chemists use g cm⁻³. Higher density means more tightly packed particles.

Uncertainty & significant figures

Every experimental measurement carries some uncertainty — limited by the instrument and the experimenter. A platform balance might read 9.4 g where an analytical balance reads 9.4213 g; both are valid measurements, with different uncertainties. Significant figures are the meaningful digits in a measurement: all certain digits, plus one final uncertain digit. A reading of 11.2 mL means 11 is certain and 2 is the estimated tenth.

Five rules cover every significant-figure question NEET asks. (i) All non-zero digits count — 285 cm has three significant figures. (ii) Leading zeros do not count — 0.0052 has two significant figures. (iii) Zeros between non-zero digits count — 2.005 has four. (iv) Trailing zeros count only if there is a decimal point — 0.200 has three, 100 has one (unless written as 100. or 1.00 × 10²). (v) Counted exact numbers (2 balls, 20 eggs) have infinite significant figures.

Scientific notation makes both extremely large and extremely small numbers tractable. Any number can be written as N × 10ⁿ where 1 ≤ N < 10. Thus 232.508 becomes 2.32508 × 10² and 0.00016 becomes 1.6 × 10⁻⁴. In addition and subtraction, the result is rounded to the fewest decimal places among the inputs; in multiplication and division, the result is rounded to the fewest significant figures.

Dimensional analysis — also called the factor-label method — converts a quantity from one unit to another by multiplying by unit factors equal to 1. To convert 3 in to centimetres, multiply by (2.54 cm / 1 in): 3 × 2.54 = 7.62 cm. The unit factor is always chosen so that the unwanted unit cancels. This single technique handles every NEET unit-conversion problem, from converting litres to cubic metres to converting days to seconds.

Laws of chemical combination

Five laws were stitched together over roughly forty years (1789–1811) to convert chemistry from an alchemical craft into a quantitative science. Each was derived from careful mass measurements, each was empirical, and together they forced the conclusion that matter is atomic. NEET tests them in two formats: identifying the law from a worked example, and applying a mass-ratio calculation to a multi-compound case.

The most fertile of these laws was Avogadro's. By distinguishing between atoms and molecules — a distinction Dalton resisted — Avogadro could explain Gay-Lussac's volume ratios. Hydrogen + oxygen combining in 2:1 by volume to give 2 volumes of water vapour only works if hydrogen and oxygen are diatomic (H₂ and O₂) and water is H₂O. Avogadro's idea was ignored for almost fifty years until Cannizzaro revived it at the 1860 Karlsruhe Congress.

Dalton's atomic theory

In 1808 John Dalton published A New System of Chemical Philosophy, in which he proposed that all matter is composed of indivisible atoms. His four postulates absorbed every empirical law that came before. The theory's power was that it explained — at a single stroke — conservation of mass, definite proportions, and multiple proportions, while predicting new chemistry that had not yet been observed.

The theory had two weaknesses. Dalton insisted that atoms of the same kind cannot combine, so he could not explain Gay-Lussac's gaseous volumes — which require diatomic molecules of elements (H₂, O₂, N₂). And he could not say why atoms combine — the answer came only with the electron and chemical bonding decades later. NCERT presents Dalton's theory as the first quantitatively useful model of matter, not as a complete one.

All matter — whether element, compound or mixture — is composed of small particles called atoms.— John Dalton, 1808

Atomic & molecular masses

Atoms are far too small to weigh individually — even the heaviest atoms come in at around 10⁻²² grams. Nineteenth-century chemists could only determine the mass of one atom relative to another. Hydrogen was initially taken as the standard with mass 1; in 1961 the international community switched to carbon-12, with one ¹²C atom defined as exactly 12 atomic mass units (amu, now written u).

Most elements occur naturally as mixtures of isotopes, so the value listed in the periodic table is an average atomic mass weighted by isotopic abundance. Carbon, for instance, exists as ¹²C (98.892%), ¹³C (1.108%), and trace ¹⁴C (≈10⁻¹⁰%). Plugging in:

(0.98892 × 12 u) + (0.01108 × 13.00335 u) + (2 × 10⁻¹² × 14.00317 u) = 12.011 uAverage atomic mass of natural carbon

Molecular mass is the sum of atomic masses of every atom in one molecule. Methane (CH₄) has molecular mass 12.011 + 4(1.008) = 16.043 u. Water has 2(1.008) + 16.00 = 18.02 u. Formula mass is used in place of molecular mass for ionic compounds — those that do not exist as discrete molecules. Sodium chloride is a three-dimensional lattice of Na⁺ and Cl⁻ ions in which one Na⁺ is surrounded by six Cl⁻ and vice versa; we therefore quote its formula mass as 23.0 + 35.5 = 58.5 u.

Mole concept & molar mass

The mole is chemistry's bridge between the microscopic world of atoms and the macroscopic world of grams. A dozen counts twelve, a score counts twenty, a gross counts 144; a mole counts 6.02214076 × 10²³ — Avogadro's number. The mole was added in 1971 as the seventh SI base unit, and in 2019 was redefined so that this number itself is fixed by convention rather than derived from a mass measurement of carbon-12.

An "elementary entity" can be anything you choose to specify — an atom, a molecule, an ion, an electron, a formula unit, or any other particle. One mole of hydrogen atoms is 6.022 × 10²³ H atoms; one mole of hydrogen molecules is 6.022 × 10²³ H₂ molecules, which contains 2 × 6.022 × 10²³ atoms. Confusing these is the single most reliable NEET trap from this chapter.

The molar mass of a substance is the mass of one mole, expressed in grams per mole. Numerically, molar mass in g mol⁻¹ equals atomic, molecular, or formula mass in u — only the units differ. Molar mass of water is 18.02 g mol⁻¹; of sodium chloride 58.5 g mol⁻¹; of methane 16.043 g mol⁻¹.

Percentage composition & empirical formula

If you have an unknown compound, the first question chemistry asks is — what is in it, and in what ratio? Mass per cent of an element in a compound is the mass of that element divided by the molar mass of the compound, multiplied by 100. For water, mass per cent of hydrogen = (2 × 1.008 / 18.02) × 100 = 11.18%; mass per cent of oxygen = (16.00 / 18.02) × 100 = 88.79%. Mass per cent is independent of the sample size — a direct corollary of the law of definite proportions.

From per cent composition, two formulae can be derived. The empirical formula gives the simplest whole-number ratio of atoms. The molecular formula gives the actual number of atoms in one molecule. For glucose, the empirical formula is CH₂O and the molecular formula is C₆H₁₂O₆ — the molecular formula is 6 × the empirical formula.

NEET 2021 tested this in compact form (Q.82): an organic compound is 78% C and 22% H by mass; find the empirical formula. Moles: C = 78/12 = 6.5; H = 22/1 = 22. Ratio: H/C = 22/6.5 = 3.38 ≈ 3. Empirical formula = CH₃. The arithmetic takes thirty seconds when the steps are memorised.

Stoichiometry & stoichiometric calculations

The word stoichiometry comes from Greek stoicheion (element) and metron (measure). It is the quantitative bookkeeping of a chemical reaction — calculating how much of each reactant is consumed and how much of each product is formed. The currency is the balanced chemical equation, which respects the law of conservation of mass by having the same number of atoms of each element on both sides.

Consider the combustion of methane:

CH₄ (g) + 2 O₂ (g) → CO₂ (g) + 2 H₂O (g)Stoichiometric coefficients: 1 : 2 : 1 : 2

The coefficients (1, 2, 1, 2) are called stoichiometric coefficients. They simultaneously represent four parallel ratios that NCERT spells out exactly:

The standard stoichiometric calculation has three steps. First, write and balance the equation. Second, convert the given mass to moles using molar mass. Third, use the mole ratio in the balanced equation to find moles of the unknown, then convert back to mass, volume, or particles as the question demands.

NEET 2023 Q.57 is a textbook application. Heat 20 g of 20% pure CaCO₃; calculate the mass of CO₂ produced. CaCO₃ → CaO + CO₂. Pure CaCO₃ in the sample = 20 × 0.20 = 4 g. Mole ratio is 1:1: 100 g CaCO₃ gives 44 g CO₂, so 4 g gives (4 × 44 / 100) = 1.76 g of CO₂. Three short lines.

Limiting reagent

Real reactions are rarely run with exactly stoichiometric amounts. One reactant is usually in excess; the other is in short supply. The reactant in short supply runs out first and stops the reaction — it limits how much product can be formed and is called the limiting reagent. The excess reactant has leftover unconsumed material at the end. Every yield calculation must be done with reference to the limiting reagent — never the excess one.

Identifying the limiting reagent takes two arithmetic steps. Convert each reactant to moles using molar mass. Then, for each reactant, calculate how many moles of product it would form if fully consumed. The reactant that gives the smaller yield is limiting. NCERT's worked example uses the Haber synthesis of ammonia.

Reactions in solution introduce one more concentration measure beyond mass-per-cent and mole fraction: molarity (mol of solute per litre of solution) and molality (mol of solute per kg of solvent). Molarity is the more commonly used in NEET stoichiometry problems; molality, being independent of temperature, appears in colligative-property calculations.