Chemistry is the branch of science that deals with the composition, structure, properties and transformation of matter. Matter is anything that occupies space and has mass. Chemistry is often called the central science because it connects physics with the life sciences and forms the foundation for understanding biological, geological and industrial processes. The study of chemistry is organised around the concept of atoms and molecules, and it explains why substances behave the way they do at the microscopic level.
Modern chemistry rests on the laws of chemical combination discovered through careful experiments by Antoine Lavoisier and Joseph Proust. These laws govern how elements combine to form compounds and led John Dalton to propose his atomic theory in 1808. This first chapter establishes the language of chemistry, including symbols, chemical formulae, atomic and molecular masses, and the mole concept, which together allow chemists to quantify every reaction and to move effortlessly between the macroscopic world of grams and the microscopic world of atoms.
Mastery of this chapter is essential because it provides the tools used in every subsequent chapter of Class 11 Chemistry. The mole concept, stoichiometry and concentration terms reappear in thermodynamics, equilibrium and practical chemistry. A student who understands limiting reagents and percentage composition can solve numerical problems across the syllabus with confidence.
Matter can be classified into mixtures and pure substances. A pure substance has a fixed composition, while a mixture has variable composition. Pure substances are further divided into elements and compounds. Elements cannot be broken down into simpler substances by chemical means, whereas compounds are made of two or more elements combined in a definite ratio.
Scientific measurements must be accurate and precise. Accuracy refers to how close a measured value is to the true value, while precision refers to how close a set of measurements are to one another. Significant figures indicate the number of meaningful digits in a measurement. In calculations, the final answer should have the same number of significant figures as the least precise measurement used.
The SI system is the internationally accepted system of units. Key SI base units include the metre (m) for length, kilogram (kg) for mass, second (s) for time, kelvin (K) for temperature, ampere (A) for electric current and mole (mol) for the amount of substance. Common derived quantities used in chemistry include density (g/mL), molarity (mol/L) and pressure (Pa). Concentration is often expressed as mass percentage, mole fraction, molarity or molality.
The five fundamental laws of chemical combination are the foundation of Dalton's atomic theory.
For example, hydrogen and oxygen always combine in a mass ratio of 1:8 to form water, and carbon and oxygen combine in ratios of 12:32 (CO2) and 12:16 (CO), which is a simple 2:1 ratio of oxygen masses, illustrating the law of multiple proportions.
Dalton's theory stated that all matter is made of indivisible atoms, all atoms of a given element are identical in mass and properties, atoms of different elements differ, compounds form when atoms of different elements combine in fixed small whole-number ratios, and atoms are neither created nor destroyed in chemical reactions.
The theory successfully explained the laws of chemical combination but had serious limitations. The discovery of subatomic particles, isotopes and isobars disproved the idea that atoms are indivisible and that all atoms of an element are identical. Isotopes (such as C-12 and C-14) are atoms of the same element with different mass numbers, while isobars are atoms of different elements with the same mass number. Compounds such as NO and N2O also show that atoms do not always combine in simple ratios.
Modern atomic theory corrects these points: atoms can be subdivided into protons, neutrons and electrons; atoms of the same element may differ in mass (isotopes); and atoms may combine in ratios other than simple whole numbers, as in sugars and proteins.
The atomic mass unit (amu or u) is defined as one-twelfth the mass of a carbon-12 atom.
$$1 \text{ u} = \frac{1}{12} \times \text{mass of one C-12 atom}$$
The atomic mass of an element is the average mass of its atoms, expressed in atomic mass units, weighted by the natural abundance of its isotopes. For example, chlorine has two isotopes, Cl-35 (75.77%) and Cl-37 (24.23%), giving an average atomic mass of about 35.5 u.
Molecular mass is the sum of the atomic masses of all atoms present in a molecule. For water:
$$M(H_2O) = 2 \times 1 + 1 \times 16 = 18 \text{ u}$$
Formula mass is used for ionic compounds such as NaCl, where individual molecules do not exist:
$$M(NaCl) = 23 + 35.5 = 58.5 \text{ u}$$
The mole is the amount of a substance that contains as many elementary entities as there are atoms in exactly 0.012 kg of carbon-12. One mole always contains Avogadro's number of entities:
$$N_A = 6.022 \times 10^{23} \text{ mol}^{-1}$$
The molar mass of a substance is its mass in grams per mole, numerically equal to its atomic or molecular mass. The number of moles is given by:
$$n = \frac{m}{M} = \frac{N}{N_A} = \frac{V}{V_m}$$
where m is mass, M is molar mass, N is number of particles and Vm is molar volume (22.4 L at STP for ideal gases).
Percentage composition gives the mass percentage of each element in a compound. For a compound with a known formula:
$$\text{Mass }\% \text{ of element} = \frac{\text{mass of element in one mole of compound}}{\text{molar mass of compound}} \times 100$$
For example, water has 88.89% oxygen and 11.11% hydrogen by mass. An empirical formula gives the simplest whole-number ratio of atoms in a compound, while a molecular formula gives the actual number of atoms in one molecule. If the empirical formula is known and the molecular mass is given, the molecular formula is:
$$\text{Molecular formula} = n \times \text{Empirical formula}, \quad n = \frac{\text{molecular mass}}{\text{empirical formula mass}}$$
For example, the empirical formula of hydrogen peroxide is HO (mass 17 u); since its molecular mass is 34 u, n = 2 and the molecular formula is H2O2.
Stoichiometry deals with the quantitative relationships between reactants and products in a balanced chemical equation. The coefficients in a balanced equation give the mole ratios, which can be converted to mass, volume or number of particles. Consider the combustion of methane:
$$\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$$
One mole of methane reacts with two moles of oxygen to give one mole of carbon dioxide and two moles of water. From any balanced equation we can convert moles of a known substance to moles of any other substance using the stoichiometric ratio.
The limiting reagent is the reactant that is completely consumed first in a reaction, and it determines the maximum amount of product formed. The other reactants are said to be in excess. The theoretical yield is the amount of product calculated from the limiting reagent, while the actual yield is what is really obtained. The percentage yield is:
$$\text{Percentage yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100$$
Concentration of solutions is expressed using several terms. Molarity (M) is moles of solute per litre of solution, molality (m) is moles of solute per kilogram of solvent, mole fraction is the ratio of moles of one component to total moles, and mass percentage is grams of solute per 100 g of solution. During dilution, M1V1 = M2V2, since the moles of solute remain constant.
| Quantity | Formula | Unit / Value |
|---|---|---|
| Number of moles | n = m / M | mol |
| Number of particles | N = n x N_A | N_A = 6.022 x 10^23 |
| Molar mass | M = m / n | g mol^-1 |
| Molarity | M = moles of solute / volume in L | mol L^-1 |
| Molality | m = moles of solute / mass of solvent (kg) | mol kg^-1 |
| Mole fraction | x_A = n_A / (n_A + n_B) | unitless |
| Molar volume (STP) | V_m = 22.4 L | L mol^-1 |
| Molecular formula factor | n = molecular mass / empirical formula mass | unitless |
| Density | d = m / V | g mL^-1 |
| Law | Scientist | Statement |
|---|---|---|
| Conservation of mass | Lavoisier | Mass is neither created nor destroyed in a reaction |
| Definite proportions | Proust | A compound always has the same composition by mass |
| Multiple proportions | Dalton | Ratios of combining masses are small whole numbers |
| Reciprocal proportions | Richter | Combining ratios with a fixed third element are related |
| Gaseous volumes | Gay Lussac | Reacting gas volumes are in simple whole-number ratios |
| Category | Example | Separation Method |
|---|---|---|
| Element | Oxygen, Sodium | Cannot be broken by chemical means |
| Compound | Water, NaCl | Chemical change |
| Homogeneous mixture | Salt in water | Evaporation, distillation |
| Heterogeneous mixture | Sand in water | Filtration, sedimentation |
Some Basic Concepts of Chemistry introduces the quantitative language of the subject. The laws of chemical combination and Dalton's atomic theory established that matter is composed of atoms combining in fixed ratios, and the mole concept provides the bridge between the microscopic world of particles and the macroscopic world of measurable quantities. Stoichiometry, percentage composition, empirical formulas and concentration terms equip students to solve realistic chemical problems. A thorough command of this chapter is indispensable for mastering every later topic in the Class 11 syllabus, and careful attention to units, significant figures and the limiting reagent will serve students well in examinations and beyond.