Some Basic Concepts of Chemistry

Chemistry is the branch of science that studies the composition, properties and transformation of matter. NCERT opens the chapter by pointing out how wide its reach is — chemistry gives us fertilisers, medicines, polymers, dyes, alloys and the fuels that run daily life, and it sits at the centre of biology, medicine, agriculture, geology and engineering.

You don't need to memorise a long list of applications for the exam, but you should be able to name two or three areas (health, energy, materials, agriculture) and give one concrete example in each, because 1-mark and 2-mark 'importance of chemistry' questions are common as openers.

Matter is anything that has mass and occupies space. It exists in three physical states — solid, liquid and gas — which differ in how closely their particles are packed and how strongly they attract one another. Solids have definite shape and volume, liquids have definite volume but take the shape of the container, and gases have neither definite shape nor definite volume; they expand to fill any container.

Matter is also classified chemically into pure substances and mixtures. A pure substance has a fixed composition throughout and cannot be separated into simpler substances by physical methods; it is further divided into elements (made of only one type of atom, e.g. Na, Cu, Ag) and compounds (two or more elements combined in a fixed ratio, e.g. water, carbon dioxide). A compound's properties are entirely different from its constituent elements. Mixtures contain two or more substances in variable proportion and can be separated by physical means; they are homogeneous (uniform composition throughout, e.g. sugar solution, air) or heterogeneous (non-uniform, e.g. sand in water, oil and water).

Every property of matter is either physical (measurable without changing chemical composition — colour, melting point, density) or chemical (describes how a substance transforms into a different substance — flammability, reactivity with acid). Physical properties are further extensive (depend on the amount of matter, e.g. mass, volume, length) or intensive (independent of amount, e.g. density, temperature, molar mass, refractive index).

Since chemistry is a quantitative science, measurement needs a common system of units. India, like most of the world, follows the International System of Units (SI), built on seven base units: metre (length), kilogram (mass), second (time), ampere (electric current), kelvin (temperature), mole (amount of substance) and candela (luminous intensity). Mass and weight are often used loosely in everyday speech but are different: mass is the amount of matter (constant everywhere, SI unit kg), while weight is the force of gravity on that mass (varies with location, unit newton).

No measurement is perfectly exact — every reading carries some uncertainty from the precision of the instrument. Scientists express this using significant figures: all the digits that are known with certainty plus one estimated (uncertain) digit. The rules examiners test most:

• All non-zero digits are significant (285 has 3 significant figures).
• Zeros between two non-zero digits are significant (2005 has 4).
• Leading zeros (before the first non-zero digit) are never significant — they only locate the decimal point (0.0025 has 2 significant figures).
• Trailing zeros in a number with a decimal point are significant (0.0500 has 3; 2.500 has 4).
• Trailing zeros in a number without a decimal point are ambiguous (100 could have 1, 2 or 3); this ambiguity is removed by writing the number in scientific notation, e.g. 1.00 × 10² clearly has 3.
• Exact numbers (from counting, or defined conversion factors like 1 dozen = 12) have infinite significant figures and never limit a calculation.

When adding or subtracting, the result is rounded to the least number of decimal places among the terms. When multiplying or dividing, the result is rounded to the least number of significant figures among the terms. Rounding follows the standard rule: if the digit to be dropped is more than 5, the preceding digit is raised by 1; if less than 5, it is left unchanged; if exactly 5, the preceding digit is raised by 1 if it is odd and left unchanged if it is even.

Quantitative experiments on chemical reactions led to five laws that Dalton later explained with his atomic theory.

• Law of conservation of mass (Lavoisier): matter can neither be created nor destroyed in a chemical reaction — the total mass of reactants equals the total mass of products in a closed system. Lavoisier verified this by carrying out reactions in sealed vessels and weighing before and after heating.
• Law of definite (constant) proportions (Proust): a given compound always contains the same elements combined in a fixed ratio by mass, regardless of its source or method of preparation. Pure water, whether from a river or made in the lab, is always 11.11% hydrogen and 88.89% oxygen by mass.
• Law of multiple proportions (Dalton): when two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other are in a ratio of small whole numbers. Example: carbon and oxygen form CO (12 g C : 16 g O) and CO₂ (12 g C : 32 g O); for a fixed 12 g of carbon, the oxygen masses 16 g and 32 g are in the ratio 1:2.
• Gay Lussac's law of gaseous volumes: when gases react, the volumes of reactants and gaseous products, measured at the same temperature and pressure, are in a ratio of small whole numbers. 2 volumes of hydrogen react with 1 volume of oxygen to give 2 volumes of water vapour.
• Avogadro's law: equal volumes of all gases, at the same temperature and pressure, contain an equal number of molecules. This law reconciled Gay Lussac's observation with Dalton's atomic theory by introducing the idea of a molecule as a group of atoms.

In 1808 John Dalton proposed an atomic theory to explain the laws of chemical combination. Its main postulates:

• Matter consists of indivisible particles called atoms.
• Atoms of a given element are identical in mass and properties; atoms of different elements differ in mass and properties.
• Atoms cannot be created or destroyed in a chemical reaction (this explains the law of conservation of mass).
• Compounds are formed when atoms of different elements combine in fixed, simple whole-number ratios (this explains the law of definite proportions and the law of multiple proportions).
• The relative number and kind of atoms in a given compound are constant.

Dalton's theory explained the first three laws of chemical combination successfully but could not explain Gay Lussac's law of gaseous volumes — it had no concept of a molecule as a group of atoms that could itself be divided during a reaction. Avogadro's hypothesis (introducing the molecule) later filled this gap.

Atoms are far too light to weigh in grams conveniently, so atomic masses are expressed relative to a standard. The modern standard is the carbon-12 isotope: one atomic mass unit (amu, symbol u) is defined as exactly 1/12th the mass of one atom of carbon-12 in its ground state. One amu equals 1.66 × 10⁻²⁴ g (1.66 × 10⁻²⁷ kg).

The atomic mass of an element is its mass relative to this standard. Most elements exist as a mixture of isotopes, so the atomic mass quoted on the periodic table is the average atomic mass — the weighted average of the masses of all naturally occurring isotopes, weighted by their percentage abundance. For example, carbon occurs as 98.892% ¹²C (mass 12.000 u) and 1.108% ¹³C (mass 13.003 u); the weighted average works out to 12.011 u, the value on the periodic table.

Molecular mass is the sum of the atomic masses of all atoms in a molecule, e.g. the molecular mass of water (H₂O) is 2(1.008) + 16.00 = 18.02 u. For ionic compounds, which do not exist as discrete molecules, the term formula mass is used instead — the sum of atomic masses of all atoms in one formula unit, e.g. NaCl has formula mass 23.0 + 35.5 = 58.5 u.

Because atoms and molecules are far too small to count or handle individually, chemists use a counting unit called the mole, the SI unit for amount of substance. One mole is defined as the amount of a substance that contains as many elementary entities (atoms, molecules, ions or other particles) as there are atoms in exactly 12 g of carbon-12. This number is the Avogadro constant, N_A = 6.022 × 10²³ mol⁻¹.

The molar mass of a substance is the mass of one mole of it, expressed in g mol⁻¹; numerically it equals the atomic mass, molecular mass or formula mass expressed in grams (e.g. molar mass of water = 18.02 g mol⁻¹). This gives three interconvertible routes to the same mole value: mass, number of particles, and (for a gas) volume. For a gas, one mole occupies a fixed volume at a given temperature and pressure regardless of its identity — this is the molar volume. At STP (standard temperature and pressure, now defined as 273.15 K and 1 bar), one mole of any ideal gas occupies 22.7 L; the older convention (1 atm, 273.15 K, sometimes called NTP) gives 22.4 L, which many competitive-exam books still use — always check which convention a question uses.

The percentage composition of a compound tells you what fraction of its total mass is contributed by each element. It is found by dividing the total mass of that element in one mole of the compound by the molar mass of the compound, and multiplying by 100.

Percentage composition data (usually from combustion or elemental analysis) is used to work backward to the empirical formula — the simplest whole-number ratio of atoms of each element in a compound. The standard method: assume 100 g of the compound, convert each element's percentage directly to grams, divide each mass by that element's atomic mass to get moles, then divide every mole value by the smallest one to get the simplest ratio (multiplying through by a small integer if the ratios aren't whole numbers). Once you have the empirical formula, the molecular formula is obtained using the compound's experimentally known molar mass: molecular formula = n × (empirical formula), where n = molar mass ÷ empirical formula mass. For ionic compounds, which exist as extended lattices rather than discrete molecules, the empirical formula is itself the standard way of writing the compound (NaCl, MgO).

Stoichiometry is the calculation of the quantitative relationships between reactants and products in a balanced chemical equation. A balanced equation tells you the mole ratio in which substances react and are formed; from that ratio, and the moles of any one substance, you can calculate the moles (and hence mass or volume) of every other substance in the reaction.

In practice, reactants are rarely mixed in the exact stoichiometric ratio. The reactant that gets used up first, stopping the reaction and limiting the amount of product formed, is the limiting reagent; the other reactant, left over at the end, is said to be in excess. To find the limiting reagent, convert every given reactant mass to moles, divide by its coefficient in the balanced equation, and the reactant with the smallest resulting value is limiting — all product calculations must then be based on it, never on the reactant in excess.

Chemists also express the concentration of a solution in several standard ways: mass percentage (mass of solute per 100 g of solution), mole fraction (moles of one component divided by total moles of all components), molarity (moles of solute per litre of solution — the most commonly used unit in the lab, though it is temperature-dependent since volume expands with heat), molality (moles of solute per kilogram of solvent — temperature-independent, since mass doesn't change with temperature), and parts per million (ppm) for very dilute solutions.