Hello, and welcome to today's chemistry lesson. We are going to explore one of the most fascinating topics in science: the structure of the atom. By the end of this lesson, you will understand what atoms are made of, how scientists discovered their components, and how we describe the arrangement of particles inside them.
Let us begin with a simple question: what is the smallest piece of matter you can imagine? Long ago, thinkers from India and Greece wondered about this too. In India, Maharishi Kanada proposed that matter is made of tiny, indestructible particles called paramanus, which combine to form larger particles called anu, now known as molecules. Meanwhile, the Greek philosopher Democritus gave these particles the name atom, from the word atomos, meaning indivisible.
In 1808, John Dalton developed this idea into a scientific theory. His atomic theory stated several important principles. First, matter consists of very small, indivisible particles called atoms. Second, atoms can neither be created nor destroyed. Third, all atoms of a given element are identical in every way, but they differ from atoms of other elements. Fourth, atoms combine in simple whole number ratios to form compounds. And finally, atoms are the smallest units of matter that take part in chemical reactions. Today, we know that some parts of Dalton's theory need modification, but his core insight that atoms participate in chemical reactions remains correct.
Now, here is where the story becomes truly exciting. Scientists discovered that atoms are not actually indivisible after all. They contain even smaller particles called sub-atomic or fundamental particles. There are three main types: electrons, protons, and neutrons.
Let us start with the electron. In 1897, J.J. Thomson was studying cathode rays, streams of particles produced in a special glass tube with a very low pressure of gas inside. When high voltage was applied, rays travelled from the negative electrode, called the cathode, toward the positive electrode. Thomson found that these rays bent toward a positive electric field, proving they carried negative charge. He named these particles electrons. An electron has a mass of 9.1 times ten to the power of minus 31 kilograms, which is incredibly small, only about one 1837th of a hydrogen atom. We represent an electron as ₋₁e⁰, where the superscript zero represents its negligible mass and the subscript negative one represents its one unit of negative charge.
But if atoms contain negative electrons, and atoms are electrically neutral, there must be positive particles too. These were discovered by E. Goldstein. He used a discharge tube with a perforated cathode, and observed rays travelling in the opposite direction to cathode rays, coming from the positive anode and passing through the holes in the cathode. These canal rays, as he called them, consisted of positively charged particles named protons. A proton carries one unit of positive charge, equal in magnitude but opposite in sign to the electron's charge. Its mass is 1.672 times ten to the power of minus 27 kilograms, roughly equal to a hydrogen atom. We write a proton as +1 p 1, where the superscript one represents one atomic mass unit and the subscript positive one represents one unit positive charge.
Thomson proposed the first model of atomic structure based on these discoveries. His plum pudding model pictured the atom as a sphere of positive charge with electrons embedded throughout, like plums in a pudding. However, this model could not explain many experimental observations, so it was eventually rejected.
The true nature of atomic structure was revealed by Ernest Rutherford in 1911. He directed fast-moving alpha particles, which are positively charged helium nuclei, at an extremely thin gold foil. Most particles passed straight through, some were slightly deflected, and a very few bounced back completely. This was astonishing. Rutherford concluded that most of an atom is empty space, with a tiny, dense, positively charged centre that he called the nucleus. This nucleus contains the protons, and it is so small that if an atom were the size of a stadium, the nucleus would be like a cricket ball at the centre.
Yet there was still a puzzle. The mass of most atomic nuclei was greater than could be accounted for by protons alone. In 1932, James Chadwick discovered the missing particle: the neutron. Neutrons reside in the nucleus alongside protons, have approximately the same mass as protons, but carry no electrical charge. We represent a neutron as 0 n 1, where the subscript zero shows it has no charge and the superscript one indicates its mass of approximately one atomic mass unit. Only ordinary hydrogen lacks neutrons; its nucleus is simply a single proton.
Now we can picture the modern atom. At the centre sits the nucleus, containing protons and neutrons held together by powerful nuclear forces. Surrounding this nucleus, electrons move in specific regions called shells or energy levels. Each shell is associated with a fixed amount of energy, with shells closer to the nucleus having lower energy. The atom remains neutral because the number of protons equals the number of electrons.
Niels Bohr explained why electrons do not spiral into the nucleus. According to his model from 1913, electrons revolve in fixed orbits or shells at very high speed without gaining or losing energy. The inward attraction of the nucleus is balanced by the motion of electrons, creating a stable structure.
Let us learn how we describe atoms precisely. The atomic number, symbol Z, is defined as the number of protons in the nucleus. In a neutral atom, this equals the number of electrons. The atomic number never changes for a given element; it is what makes each element unique. For example, oxygen has atomic number 8, meaning every oxygen atom has 8 protons.
The mass number, symbol A, is the sum of protons and neutrons in the nucleus. Since electrons have negligible mass, virtually all atomic mass resides in the nucleus. We can find neutrons by subtracting: number of neutrons equals mass number minus atomic number. Take sodium: with atomic number 11 and mass number 23, it has 12 neutrons. We write sodium symbolically as 11 Na 23 or 11 23 Na.
Electrons arrange themselves in shells according to specific rules. The maximum electrons in any shell is given by two n squared, where n is the shell number. This is known as the Bohr-Bury scheme. So the K shell holds 2 electrons, the L shell holds 8, the M shell holds 18, and the N shell holds 32. However, the outermost shell can never exceed 8 electrons; this is called the octet rule. If there is only one shell, it can hold only 2 electrons, called the duplet rule. Electrons fill inner shells completely before occupying outer ones.
Let us see how this works. Hydrogen, with 1 electron, has configuration 1. Neon, with 10 electrons, fills the first shell with 2 and the second with 8. Sodium, with 11 electrons, has configuration 2, 8, 1. Potassium, with 19 electrons, fills K and L shells with 2 and 8, but cannot put 9 in the M shell as that would exceed 8 in the outermost shell, so it places 8 in M and the final electron in N, giving 2, 8, 8, 1.
The outermost shell is called the valence shell, and its electrons are valence electrons. These determine how an atom combines with others, a property called valency. Valency is the combining capacity of an element, specifically the number of hydrogen atoms that combine with or can be displaced by one atom of that element.
Some elements show variable valency. Iron, for example, can form ferrous ions with valency 2, written as Fe²⁺, or ferric ions with valency 3, written as Fe³⁺. Copper similarly shows valencies of 1 and 2. Modern notation uses Roman numerals: iron two and iron three, where neither the name nor symbol changes.
When atoms gain or lose electrons, they become charged particles called ions. Positive ions, or cations, form when atoms lose electrons. Negative ions, or anions, form when atoms gain electrons. A group of atoms of different elements that behaves as a single unit with a positive or negative charge is called a radical. Positively charged radicals like NH4+ are called basic radicals or cations, while negatively charged radicals like SO4 2- and NO3- are called acid radicals or anions.
Finally, let us consider isotopes. These are atoms of the same element with the same atomic number but different mass numbers, meaning they have different numbers of neutrons. Hydrogen has three isotopes: protium or ordinary hydrogen with no neutrons, deuterium or heavy hydrogen with one neutron, and tritium or very heavy hydrogen with two neutrons. Carbon has isotopes with mass numbers 12, 13, and 14. Isotopes have identical chemical properties because they have the same electronic configuration, but they differ in physical properties like density and melting point due to their different masses.
Let us recap the key points from today's lesson. First, atoms are divisible and contain three fundamental particles: electrons, protons, and neutrons. Second, the atomic number Z equals the number of protons, and the mass number A equals protons plus neutrons. Third, electrons occupy shells with maximum capacity given by two n squared, with the outermost shell limited to eight electrons, or two for the first shell. Fourth, the valence shell determines chemical combining power, or valency. Fifth, ions form when atoms gain or lose electrons, and radicals are charged groups of atoms. Sixth, isotopes are atoms of the same element with different numbers of neutrons.
You have now journeyed from ancient philosophical ideas to our modern understanding of atomic structure. This knowledge forms the foundation for everything you will learn in chemistry. Keep exploring, keep questioning, and remember that science is a story of continuous discovery. Until next time, stay curious and keep learning.