Atoms and Molecules

Everything around us — air, water, rocks, our own bodies — is made of matter. If we keep dividing a piece of matter into smaller and smaller parts, we eventually reach the tiniest particle that still keeps the properties of that substance. The atom is the smallest particle of an element that can take part in a chemical reaction. Atoms are extremely small — far too tiny to see even with an ordinary microscope — yet they are the basic building blocks from which all matter is made.

The first complete scientific theory about atoms was given by the scientist John Dalton, and is called Dalton's atomic theory. Its main ideas are: all matter is made of tiny particles called atoms; atoms cannot be created, destroyed, or divided (in a chemical reaction); atoms of the same element are identical in mass and properties, while atoms of different elements are different; and atoms combine in simple whole-number ratios to form compounds. This theory was a huge step forward, because it explained many observations about how substances combine.

Dalton's theory, though very useful, had some limitations, which became clear as science advanced. Dalton said the atom was indivisible (could not be divided), but later discoveries showed that atoms are themselves made of even smaller particles (protons, neutrons, and electrons). He also said atoms of the same element are always identical, but we now know that atoms of the same element can have different masses (these are called isotopes, studied later). So some of Dalton's ideas had to be modified.

Even with these limitations, Dalton's atomic theory remains the foundation of modern chemistry, because its core idea — that matter is made of atoms that combine in fixed ratios — is correct and is still used today. We now understand the atom in more detail, knowing it has an internal structure of subatomic particles. So the atom is the basic particle of an element, Dalton's theory first described atoms scientifically, and later work refined it — setting the stage for studying what atoms are actually made of.

Although Dalton thought the atom was indivisible, we now know that the atom is made of even smaller particles, called subatomic particles. There are three main subatomic particles: the proton, the neutron, and the electron. Each has its own charge and mass and is found in a particular part of the atom. Understanding these particles tells us how an atom is built and explains its properties.

The three particles differ in charge, mass, and location. The proton carries a positive (+) charge and has a relatively large mass; it is found in the nucleus (the dense centre of the atom). The neutron carries no charge (it is neutral) and has a mass about the same as a proton; it is also found in the nucleus. The electron carries a negative (−) charge and is very light (its mass is negligible compared with a proton); electrons move around the nucleus in shells (orbits). So the heavy, positively charged nucleus (protons and neutrons) is at the centre, with light, negatively charged electrons around it. An atom as a whole is electrically neutral, because it has equal numbers of protons and electrons, so the positive and negative charges balance.

Two important numbers describe an atom. The atomic number (Z) is the number of protons in the nucleus of an atom. It is the identity of the element — every atom of a given element has the same atomic number (for example, every hydrogen atom has Z = 1, and every carbon atom has Z = 6). The mass number (A) is the total number of protons and neutrons in the nucleus, since these particles make up almost all of the atom's mass. So A = Z + N, where N is the number of neutrons.

These numbers are written in a standard notation as $^A_Z X$, where X is the element's symbol, A (top) is the mass number, and Z (bottom) is the atomic number. From this notation we can read off the structure of the atom: Z gives the number of protons (and, in a neutral atom, the number of electrons), and the number of neutrons is found as N = A − Z. For example, for $^{12}_{6}C$ (carbon), Z = 6 (6 protons and 6 electrons) and A = 12, so the number of neutrons = 12 − 6 = 6. Understanding subatomic particles and these two numbers lets us describe exactly what any atom is made of.

We have learned that an atom is described by its atomic number (number of protons) and its mass number (protons + neutrons). It turns out that atoms can be related to each other in special ways depending on which of these numbers they share. Three such relationships are described by the terms isotopes, isobars, and isotones. These help us classify and compare atoms.

Isotopes are atoms of the same element that have the same atomic number but different mass numbers. In other words, they have the same number of protons (so they are the same element) but different numbers of neutrons (so their masses differ). A classic example is hydrogen, which has three isotopes: ordinary hydrogen ($^1_1H$, no neutrons), deuterium ($^2_1H$, one neutron), and tritium ($^3_1H$, two neutrons) — all have one proton (Z = 1) but different mass numbers. Isotopes of an element have the same chemical properties (because chemistry depends on electrons, which are equal in number) but slightly different physical properties due to their different masses.

Isotopes have many useful applications. In medicine, certain isotopes (radioactive ones) are used to diagnose and treat diseases — for example, in cancer treatment and in medical imaging. In dating, the isotope carbon-14 is used to find the age of ancient objects and fossils (carbon dating). Isotopes are also used as tracers and in nuclear power. So although isotopes are just atoms of the same element with different masses, they have important real-world uses.

Two related terms compare atoms of different elements. Isobars are atoms of different elements that have the same mass number but different atomic numbers — they have different numbers of protons but happen to have the same total of protons and neutrons. Isotones are atoms that have the same number of neutrons but different atomic numbers (and so different mass numbers). To summarise: isotopes share the atomic number (same element), isobars share the mass number, and isotones share the number of neutrons. Recognising these relationships deepens our understanding of how atoms can be similar or different.

The electrons in an atom are not scattered randomly; they are arranged in definite energy levels or shells around the nucleus. These shells are named, from the innermost outward, K, L, M, N, … The arrangement of electrons in these shells is called the electronic configuration of the atom. Knowing the electronic configuration helps us understand how an atom will behave chemically — that is, how it will combine with other atoms.

There are simple rules for filling the shells with electrons. Electrons fill the shells starting from the innermost (lowest energy) shell. Each shell can hold only a certain maximum number of electrons: the K shell holds up to 2 electrons, the L shell up to 8, the M shell up to 8 (for the first 18 elements), and so on. So the filling order is 2, 8, 8, … For example, the electronic configuration of sodium (atomic number 11) is 2, 8, 1 — two electrons in K, eight in L, and one in M. We fill each inner shell before moving to the next.

The electrons in the outermost shell are especially important and are called the valence electrons. These outermost electrons are the ones involved when atoms combine with one another, so they determine the chemical behaviour of the atom. Atoms are most stable when their outermost shell is full (8 electrons, or 2 for the K shell) — this is why the noble gases, which already have full outer shells, are very unreactive.

The combining capacity of an atom is described by its valency. Valency is the combining capacity of an atom, and it is related to the number of valence electrons. An atom tends to gain, lose, or share electrons to complete its outermost shell, and the valency tells us how many electrons it will do this with. For elements with 1, 2, or 3 valence electrons, the valency is usually equal to the number of valence electrons (which the atom tends to lose); for those with 5, 6, or 7, the valency is usually 8 minus the number of valence electrons (the number it tends to gain). For example, sodium (2, 8, 1) has 1 valence electron and a valency of 1, while oxygen (2, 6) has 6 valence electrons and a valency of 8 − 6 = 2. Understanding electronic configuration and valency lets us predict how atoms combine to form compounds.

When atoms combine, they form molecules. A molecule is the smallest particle of a substance (element or compound) that can exist independently and shows the properties of that substance. The composition of a molecule is shown by its molecular formula, which uses symbols and numbers to tell us which atoms, and how many of each, are present. For example, the formula H₂O tells us a water molecule has 2 hydrogen atoms and 1 oxygen atom, and CO₂ tells us a carbon dioxide molecule has 1 carbon and 2 oxygen atoms.

Just as atoms have mass, molecules have mass too. The molecular mass of a substance is the sum of the atomic masses of all the atoms in its molecule. To find it, we add up the atomic mass of each atom in the formula. For example, for water (H₂O), taking the atomic mass of hydrogen as 1 and oxygen as 16: molecular mass = (2 × 1) + 16 = 18. So one molecule of water has a molecular mass of 18 (in atomic mass units). Molecular mass lets us compare the masses of different molecules.

Because atoms and molecules are so incredibly tiny, chemists count them in very large groups, using a unit called the mole. The mole is the unit used to count particles (atoms, molecules) in chemistry, just as a "dozen" means 12. One mole of any substance contains a fixed, very large number of particles, known as Avogadro's number, which is 6.022 × 10²³. So one mole contains 6.022 × 10²³ particles — whether atoms, molecules, or other particles. This huge number is needed because individual atoms are so small.

The mole concept connects the tiny world of atoms to the amounts we can actually measure. Importantly, one mole of a substance has a mass equal to its molecular mass in grams. For example, one mole of water (molecular mass 18) has a mass of 18 grams and contains 6.022 × 10²³ water molecules. This lets chemists weigh out a known number of particles simply by weighing a substance. So molecular formulas tell us what a molecule contains, molecular mass tells us how heavy it is, and the mole (with Avogadro's number) lets us count and weigh atoms and molecules — completing our study of atoms and molecules.