The Atom Turned Out to Be Almost Entirely Empty

Because almost all of an atom's mass sits in a nucleus at its centre that is vanishingly small compared with the atom itself.

The electrons occupy the enormous volume around that nucleus, but they weigh almost nothing — a proton is about 1836 times heavier than an electron. So the atom is a tiny, dense, heavy centre surrounded mostly by space.

A table feels solid because the electron clouds of its atoms push back against yours, not because the space is filled. This page covers the first part of the ICSE Class 8 Chemistry chapter on atomic structure: atoms and molecules, the theory that started it, the particles inside, and the two models that describe how they are arranged.

An atom is the smallest particle of an element that can take part in a chemical reaction. A molecule is the smallest particle of a substance that can exist independently and still show all the properties of that substance.

The key difference is independent existence. A single atom of most elements cannot exist on its own for long — it joins up. A molecule can.

Molecules of elements contain atoms of only one kind:

- Monatomic — one atom: helium ($\text{He}$), neon ($\text{Ne}$), argon ($\text{Ar}$)
- Diatomic — two atoms: hydrogen (${\text{H}}_2$), oxygen (${\text{O}}_2$), nitrogen (${\text{N}}_2$), chlorine (${\text{Cl}}_2$)
- Triatomic — three atoms: ozone (${\text{O}}_3$)
- Polyatomic — many atoms: phosphorus (${\text{P}}_4$), sulphur (${\text{S}}_8$)

This number of atoms in one molecule is its atomicity.

Molecules of compounds contain atoms of two or more kinds: water (${\text{H}}_2\text{O}$), carbon dioxide (${\text{CO}}_2$), ammonia (${\text{NH}}_3$), sulphuric acid (${\text{H}}_2{\text{SO}}_4$).

Why the distinction matters in practice. $\text{O}$ and ${\text{O}}_2$ are not the same thing. $\text{O}$ is one atom of oxygen, which cannot exist freely; ${\text{O}}_2$ is a molecule of oxygen gas, which is what you breathe. Writing one where the other belongs changes the chemistry of every equation.

And the noble gases are the exception. Helium, neon and argon exist as single atoms, so for them the atom is the molecule. This is because their outermost shells are already complete, which is exactly why they do not need to join anything — an idea the third part of this chapter develops.

The postulates, which set out a systematic description of matter as particles:

- All matter is made of extremely small, indivisible particles called atoms.
- Atoms of the same element are identical in mass and in properties.
- Atoms of different elements differ in mass and in properties.
- Atoms combine in small whole-number ratios to form compounds.
- Atoms can be neither created nor destroyed in a chemical reaction.

What it explained well. The last two postulates account directly for the law of conservation of mass and for a compound always having a fixed composition — water being hydrogen and oxygen in a mass ratio of $1:8$ no matter where it comes from. Those successes are why the theory held up for so long.

Its limitations, each of which later work exposed:

- Atoms are not indivisible. They contain electrons, protons and neutrons, and can be split.
- Atoms of the same element are not always identical in mass. Isotopes of the same element have different masses — chlorine has atoms of mass $35$ and $37$.
- Atoms of different elements can have the same mass. Such atoms are called isobars.
- It does not explain why atoms combine, or the nature of the force holding them together.
- It does not distinguish clearly between an atom and a molecule.

The point worth taking from this. A theory being wrong in places does not make it worthless. Dalton's whole-number-ratio postulate is still used every time you write a formula, and his conservation postulate is still used every time you balance an equation. What replaced the theory was not a rejection of it but a correction of the parts that failed testing — which is how scientific models normally change.

Three subatomic particles, and they differ on every count that matters.

Electron — symbol $e^{-}$, relative charge $-1$, relative mass about $\frac{1}{1836}$ of a proton, located in shells revolving around the nucleus.

Proton — symbol $p^{+}$, relative charge $+1$, relative mass $1$, located inside the nucleus.

Neutron — symbol $n$, relative charge $0$, relative mass $1$, located inside the nucleus.

Why the atom is electrically neutral. In any neutral atom the number of electrons equals the number of protons, so the $-1$ charges exactly cancel the $+1$ charges. Sodium has $11$ of each, and its net charge is zero.

Why nearly all the mass is in the nucleus. Protons and neutrons each have relative mass $1$; electrons have about $\frac{1}{1836}$ of that. Take a sodium atom with $11$ protons, $12$ neutrons and $11$ electrons:

$$\text{mass of nucleons}=11+12=23$$

$$\text{mass of electrons}=\frac{11}{1836}\approx 0.006$$

So the electrons contribute less than a hundredth of one mass unit out of $23$. This is why the mass number counts only protons and neutrons and ignores electrons entirely.

Protons and neutrons together are called nucleons, for exactly that reason.

The particle that can actually change. Only electrons are gained, lost or shared in chemical reactions — and only from the outermost shell. The nucleus is untouched by ordinary chemistry, which is why the identity of an element never changes in a chemical reaction. That is the same fact the chapter on static electricity used to explain charging by friction: rubbing moves electrons and never protons.

Rutherford's nuclear model places all the positive charge and nearly all the mass in a tiny central nucleus, with electrons moving around it.

Its features:

- The atom has a small, dense, positively charged nucleus at its centre.
- Nearly the whole mass of the atom is concentrated there.
- Most of the atom is empty space.
- Electrons revolve around the nucleus, like planets around a sun.
- The nucleus is extremely small compared with the atom as a whole.

Its limitation. A charged particle moving in a curved path should continuously lose energy. If electrons did that, they would spiral inwards and crash into the nucleus, and no atom would be stable — yet atoms plainly are. The model could not explain why.

Bohr's model fixed that by restricting where electrons may be.

Its features:

- Electrons revolve only in certain fixed circular paths called orbits, shells or energy levels.
- Each shell has a definite, fixed energy, which is why they are called energy levels.
- An electron revolving within a shell does not radiate energy, so the atom is stable.
- Shells are named K, L, M, N outwards from the nucleus, numbered $n=1,2,3,4$.
- Energy increases as you move outwards — the K shell is lowest.
- An electron absorbs energy to jump to a higher shell and releases energy when it falls back.

How to picture the two together. Rutherford supplied the nucleus; Bohr supplied the shells. The combined picture — a dense nucleus of protons and neutrons, with electrons in fixed shells around it — is the model used for the rest of Class 8 chemistry, and it is enough to explain valency, ion formation and every formula you will write.

The boundary of the model. Bohr's neat circular orbits are a simplification; the modern description replaces them with regions of probability. But for predicting how elements combine, the shell picture gives the right answers, which is why it is the one taught and the one the next part of this chapter builds on.