Comprehensive theory, key formulas, diagrams, and memory aids for Atoms and Molecules.
Ancient Indian and Greek philosophers have always wondered about the unknown and unseen form of matter. The idea of divisibility of matter was considered long back in India, around 500 BC. An Indian philosopher, Maharishi Kanad, postulated that if we go on dividing matter (padarth), we shall get smaller and smaller particles. Ultimately, a stage will come when we shall come across the smallest particles beyond which further division will not be possible. He named these particles Parmanu. Later, Greek philosophers Democritus and Leucippus called these indivisible particles "atoms" (meaning indivisible).
By the end of the eighteenth century, scientists recognized the difference between elements and compounds and naturally became interested in finding out how and why elements combine. This chapter explores the fundamental laws governing these combinations and introduces the concepts of atoms, molecules, and the mole.
The foundation of chemical sciences was laid by Antoine L. Lavoisier by establishing two important laws of chemical combination. These laws were formulated after much experimentation by Lavoisier and Joseph L. Proust.
This law states that mass can neither be created nor destroyed in a chemical reaction. During any chemical change, the total mass of the reactants (the substances that react) is exactly equal to the total mass of the products (the substances formed). For example, if you burn carbon in oxygen to form carbon dioxide in a closed container, the weight of the container before and after the reaction remains exactly the same.
Proust noted that many compounds were composed of two or more elements and each such compound had the same elements in the same proportion, irrespective of where the compound came from or who prepared it. This law states that in a chemical substance, the elements are always present in definite proportions by mass. For example, in water ($H_2O$), the ratio of the mass of hydrogen to the mass of oxygen is always $1:8$, whatever the source of water. Thus, if 9 g of water is decomposed, 1 g of hydrogen and 8 g of oxygen are always obtained.
To provide an explanation for these laws, British chemist John Dalton proposed his atomic theory in 1808. His theory provided a theoretical basis for the laws of chemical combination.
Postulates of Dalton's Atomic Theory: 1. All matter is made of very tiny particles called atoms, which participate in chemical reactions. 2. Atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction (explains the Law of Conservation of Mass). 3. Atoms of a given element are identical in mass and chemical properties. 4. Atoms of different elements have different masses and chemical properties. 5. Atoms combine in the ratio of small whole numbers to form compounds. 6. The relative number and kinds of atoms are constant in a given compound (explains the Law of Constant Proportions).
Have you ever observed a mason building a wall? From walls to a room and then a building, the building block is the brick. Similarly, the building blocks of all matter are atoms. Atoms are incredibly small. Their radius is measured in nanometres ($1 \text{ nm} = 10^{-9} \text{ m}$).
Dalton's theory proposed that each element had a characteristic atomic mass. Because atoms are so small, determining their absolute mass was difficult. Scientists started evaluating relative atomic masses. The universally accepted mass reference is the carbon-12 isotope. One atomic mass unit (amu or simply 'u') is a mass unit equal to exactly one-twelfth ($1/12^{th}$) the mass of one atom of carbon-12. Using this, Hydrogen is $1\text{u}$, Oxygen is $16\text{u}$, and Carbon is $12\text{u}$.
A molecule is in general a group of two or more atoms that are chemically bonded together (held together by attractive forces). A molecule can be defined as the smallest particle of an element or a compound that is capable of an independent existence and shows all the properties of that substance.
The molecules of an element are constituted by the same type of atoms. For example, a molecule of oxygen consists of two atoms of oxygen and is known as a diatomic molecule ($O_2$). If 3 atoms of oxygen unite, we get ozone ($O_3$). The number of atoms constituting a molecule is known as its atomicity. * Argon (Ar), Helium (He): Monoatomic * Oxygen ($O_2$), Hydrogen ($H_2$): Diatomic * Phosphorus ($P_4$): Tetra-atomic * Sulphur ($S_8$): Poly-atomic
Atoms of different elements join together in definite proportions to form molecules of compounds. For example, in water ($H_2O$), two atoms of hydrogen combine with one atom of oxygen.
Compounds composed of metals and non-metals contain charged species. The charged species are known as ions. An ion can be a negatively or positively charged particle. * A negatively charged ion is called an anion (e.g., $Cl^-$). * A positively charged ion is called a cation (e.g., $Na^+$). When sodium and chlorine combine, they form sodium chloride ($NaCl$), an ionic compound.
The chemical formula of a compound is a symbolic representation of its composition. To write it, you need to learn the symbols and combining capacity of the elements. The combining power (or capacity) of an element is known as its valency. Valency tells us how atoms of an element will combine with the atom(s) of another element.
Rules for writing chemical formulae: 1. The valencies or charges on the ion must balance. 2. When a compound consists of a metal and a non-metal, the name or symbol of the metal is written first (e.g., Calcium oxide, $CaO$). 3. In compounds formed with polyatomic ions, the number of ions present in the molecule is indicated by enclosing the formula of the ion in a bracket and writing the number of ions outside the bracket (e.g., $Mg(OH)_2$).
Example: Formula for Hydrogen chloride. Symbol: H Cl Valency: 1 1 Formula: $HCl$ (criss-cross the valencies)
The molecular mass of a substance is the sum of the atomic masses of all the atoms in a molecule of the substance. It is expressed in atomic mass units (u). For example, the molecular mass of water ($H_2O$): $(2 \times \text{Atomic mass of H}) + (1 \times \text{Atomic mass of O}) = (2 \times 1) + (1 \times 16) = 18\text{u}$.
When dealing with chemical reactions, we cannot count individual atoms. We need a unit that bridges the gap between atomic mass and macroscopic mass (grams). This unit is the mole.
One mole of any species (atoms, molecules, ions, or particles) is that quantity in number having a mass equal to its atomic or molecular mass in grams. The number of particles present in 1 mole of any substance is fixed, with a value of $6.022 \times 10^{23}$. This is an experimentally obtained value and is called the Avogadro Constant or Avogadro Number (represented by $N_A$).
The mass of 1 mole of a substance is called its molar mass.
The discovery of the laws of chemical combination led to Dalton's atomic theory, which established the atom as the fundamental building block of matter. Atoms combine to form molecules and ions, which are represented concisely using chemical formulae based on valency. The concept of the mole is a powerful tool in chemistry, allowing scientists to count atoms and molecules by weighing them, thus linking the microscopic world to the macroscopic world.