Atoms and Molecules

When substances react chemically, do their masses follow any fixed pattern? Two important laws answer this question and form the foundation of chemistry.

Law of Conservation of Mass: This law states that mass can neither be created nor destroyed in a chemical reaction. The total mass of the reactants before a reaction is always equal to the total mass of the products after the reaction. This law was established by Antoine L. Lavoisier in 1789 through careful experiments carried out in closed (sealed) vessels, so that no matter could escape or enter during the reaction.

For example, when sodium sulphate solution reacts with barium chloride solution inside a sealed flask, a white precipitate of barium sulphate is formed along with sodium chloride in solution. Weighing the flask before and after the reaction shows that the mass remains exactly the same, confirming that no mass is lost or gained.

Law of Constant Proportions (Law of Definite Proportions): Given by the chemist Joseph Proust, this law states that in a chemical compound the elements are always present in a definite proportion by mass, no matter how or from where the compound is obtained. For example, water obtained from a river, from the sea, or prepared in a laboratory always contains hydrogen and oxygen combined in the same ratio of 1:8 by mass.

Together, these two laws show that chemical combination is not random but follows fixed, predictable rules, which later led John Dalton to propose his atomic theory.

The laws of chemical combination raised an important question: why do elements always combine in fixed proportions? John Dalton answered this in 1808 with his atomic theory, which proposed that matter is made up of extremely small, indivisible particles called atoms.

Atoms are so small that they cannot be seen even with a powerful optical microscope; the radius of a single atom is typically around 10^-10 m. To make communication simple, every element is given a unique symbol, usually the first letter (capitalised) or the first two letters of its name. For example, hydrogen is denoted by H, oxygen by O, and carbon by C. Some symbols are derived from Latin names, such as Na for sodium (natrium), K for potassium (kalium), and Fe for iron (ferrum).

Since a single atom is far too light to be weighed on any laboratory balance, chemists express atomic mass as a relative value. One atomic mass unit (u) is defined as exactly 1/12th of the mass of one atom of carbon-12. On this scale, a hydrogen atom has a mass of about 1 u, an oxygen atom about 16 u, and a sodium atom about 23 u. These relative atomic masses are used throughout chemical calculations to compare atoms of different elements.

Atoms rarely exist alone; they usually combine to form molecules. A molecule is the smallest particle of an element or a compound that can exist independently and still show all the chemical properties of that substance.

Molecules of elements are made up of atoms of the same element joined together. The number of atoms present in one molecule of an element is called its atomicity. Noble gases like helium (He) exist as single atoms and are called monatomic. Gases such as hydrogen (H₂), oxygen (O₂), nitrogen (N₂) and chlorine (Cl₂) are diatomic, made of two atoms each. Some elements form more complex molecules: ozone (O₃) is triatomic, phosphorus (P₄) is tetra-atomic, and sulphur (S₈) has eight atoms in each molecule.

Molecules of compounds are formed when atoms of two or more different elements combine in a fixed ratio. For example, a molecule of water (H₂O) contains two atoms of hydrogen and one atom of oxygen; a molecule of carbon dioxide (CO₂) contains one carbon atom and two oxygen atoms; and a molecule of ammonia (NH₃) contains one nitrogen atom and three hydrogen atoms.

Many substances, especially those formed between metals and non-metals, do not form neutral molecules but instead form charged particles called ions. An ion is an atom or a group of atoms that carries a positive or negative charge because it has lost or gained electrons.

A positively charged ion is called a cation; it forms when an atom loses one or more electrons, for example Na⁺, Ca²⁺, and Al³⁺. A negatively charged ion is called an anion; it forms when an atom gains one or more electrons, for example Cl^- and O^2-. Some ions, called polyatomic ions, are groups of atoms bonded together that carry a net charge and behave as a single unit, such as the hydroxide ion (OH^-), the ammonium ion (NH₄⁺), the carbonate ion (CO₃^2-), and the sulphate ion (SO₄^2-).

The combining capacity of an element or ion is called its valency. To write a compound's chemical formula, the symbols of the combining ions are placed side by side and their valencies (ignoring the sign) are criss-crossed and written as subscripts; the resulting ratio is then simplified to the smallest whole numbers. For example, to write the formula of aluminium oxide, aluminium has a valency of 3 (Al³⁺) and oxygen has a valency of 2 (O^2-); crossing over the valencies gives the formula Al₂O₃.

Once atomic masses are known, the mass of a whole molecule can be calculated. The molecular mass of a substance is the sum of the atomic masses of all the atoms present in one molecule of that substance, expressed in atomic mass units (u). For example, the molecular mass of water (H₂O) is calculated as 2 × (mass of H) + 1 × (mass of O) = 2(1u) + 16u = 18u.

Some compounds, especially those formed between metals and non-metals such as sodium chloride (NaCl) or calcium carbonate (CaCO₃), are made up of ions arranged in a large, repeating network rather than existing as separate molecules. For such ionic compounds, the term formula unit mass is used instead of molecular mass. It is calculated in exactly the same way, by adding up the atomic masses of all the atoms shown in the chemical formula. For example, the formula unit mass of sodium chloride is 23u (Na) + 35.5u (Cl) = 58.5u.

Because atoms and molecules are extremely tiny, chemists needed a convenient way to count large numbers of particles while working with masses that can actually be measured in a laboratory, in grams. This is done using the mole, the SI unit for the amount of a substance. One mole of any substance contains exactly 6.022 × 10²³ particles (atoms, molecules, or ions); this fixed number is called the Avogadro constant, named after the scientist Amedeo Avogadro.

The mass of one mole of a substance, expressed in grams, is called its molar mass, and it is numerically equal to the atomic, molecular, or formula unit mass of that substance in u. For example, since the molecular mass of water is 18u, the molar mass of water is 18 g, and 18 g of water contains 6.022 × 10²³ molecules. This relationship allows an easy conversion between the mass of a substance and the number of moles it contains: number of moles = given mass ÷ molar mass.