Welcome to the World of the Atom!

Have you ever wondered what you, your phone, and the air around you are actually made of? Everything in the universe is built from tiny building blocks called atoms. In this chapter, we are going to look inside the atom to see how it’s built, how it has different "versions" called isotopes, and how our ideas about atoms have changed over hundreds of years.

This is the foundation for everything else in the Radioactivity section, so getting a good grip on this now will make the later topics much easier!

1. What is an Atom?

An atom consists of a tiny central nucleus surrounded by even tinier particles called electrons. The nucleus is made up of two types of particles: protons and neutrons.

The Subatomic Particles

To understand atoms, you need to know the properties of the three main particles:

  • Protons: Found in the nucleus. They have a positive charge.
  • Neutrons: Found in the nucleus. They are neutral (no charge).
  • Electrons: Found orbiting the nucleus in shells. They have a negative charge.
Relative Mass and Charge

In Physics, we use "relative" numbers because the actual mass and charge of these particles are too small to use easily in everyday calculations.

Proton: Relative Mass = \(1\), Relative Charge = \(+1\)
Neutron: Relative Mass = \(1\), Relative Charge = \(0\)
Electron: Relative Mass = \(1/1835\) (very close to \(0\)), Relative Charge = \(-1\)

Quick Tip: Remember that in a neutral atom, the number of protons always equals the number of electrons. This means the total positive charge cancels out the total negative charge, leaving the atom with no overall charge.

2. The Scale of the Atom (Order of Magnitude)

Atoms are incredibly small, but the nucleus is even smaller! It is important to have an idea of the order of magnitude (the powers of 10) for these sizes:

  • The Atom: The radius of an atom is about \(10^{-10}\) metres.
  • The Nucleus: The radius of the nucleus is about \(10^{-15}\) metres.

Analogy: If an atom were the size of a massive football stadium, the nucleus would be like a small pea sitting right in the middle of the pitch. The rest of the stadium is mostly empty space where the electrons live!

3. Atomic Number, Mass Number, and Isotopes

When you look at a Periodic Table, you will see two numbers next to every element symbol. For example, Carbon is often shown with a \(12\) and a \(6\).

Atomic Number (\(Z\))

The Atomic Number is the number of protons in the nucleus. This is the most important number because it tells you which element you are looking at. If an atom has 6 protons, it must be Carbon.

Mass Number (\(A\))

The Mass Number is the total number of protons PLUS neutrons in the nucleus. We don't count electrons in the mass because they are so light they barely make a difference.

How to calculate neutrons:
\(Number\ of\ Neutrons = Mass\ Number - Atomic\ Number\)

What are Isotopes?

An isotope is a different "version" of an element. All isotopes of an element have:

  • The same number of protons (same atomic number).
  • A different number of neutrons (different mass number).

Example: Carbon-12 has 6 protons and 6 neutrons. Carbon-14 has 6 protons and 8 neutrons. They are both Carbon, but Carbon-14 is slightly heavier and is often radioactive!

4. Electrons and Ions

Electron Orbits

Electrons sit in fixed orbits (shells) at different distances from the nucleus. They can move between these orbits if they gain or lose energy:

  • Moving to a higher orbit (further from nucleus): The atom absorbs electromagnetic radiation.
  • Moving to a lower orbit (closer to nucleus): The atom emits (sends out) electromagnetic radiation.
Forming Ions

If an atom loses or gains an electron, it is no longer neutral. It becomes an ion. This process is called ionisation.

  • If an atom loses an electron, it becomes a positive ion (because it has more positive protons than negative electrons).
  • If an atom gains an electron, it becomes a negative ion.

5. Changing Models of the Atom

Our understanding of the atom has changed over time as scientists discovered new evidence. You need to know how the model evolved:

The Dalton Model (Early 1800s)

John Dalton thought atoms were solid, indivisible spheres—like tiny billiard balls. He believed each element had its own type of atom.

The Plum Pudding Model (1897)

After discovering the electron, J.J. Thomson suggested the "Plum Pudding" model. He thought the atom was a sphere of positive charge with negative electrons dotted inside it (like fruit in a pudding).

The Rutherford Model (1911)

Ernest Rutherford fired alpha particles at a thin gold foil. He was shocked to find that while most went through, some bounced back! This proved:

  • The atom is mostly empty space.
  • There is a tiny, dense, positively charged nucleus in the centre.

The Bohr Model (1913)

Niels Bohr suggested that electrons don't just float around; they orbit the nucleus in fixed shells at specific distances. This explained why atoms don't just collapse.

The Modern Model

Later, James Chadwick discovered the neutron, which explained why nuclei were heavier than expected based on just protons. This gave us the "nuclear atom" model we use today.

Key Takeaway: Science is always changing! When we get new evidence (like the gold foil experiment), we have to change our models to fit the facts.

Summary Checklist

[ ] Can you state the relative mass and charge of protons, neutrons, and electrons?
[ ] Do you know the typical size of an atom (\(10^{-10}\) m) vs a nucleus (\(10^{-15}\) m)?
[ ] Can you define an isotope?
[ ] Can you calculate the number of neutrons if given the mass and atomic numbers?
[ ] Do you understand how an atom becomes a positive ion?
[ ] Can you describe why Rutherford's experiment changed the "Plum Pudding" model?

Next Chapter: You will learn how unstable isotopes can decay by emitting radiation (Alpha, Beta, and Gamma)!