Understand The Nuclear Atom — Understand Structure 1.2—The nuclear atom - IB Chemistry

SECTION 1

Historical Context & Motivation

For centuries, people assumed matter was continuous — you could keep dividing it forever. The ancient Greeks proposed the idea of the atom (from the Greek atomos, meaning "indivisible"), but it took more than two thousand years before scientists could actually test what atoms look like on the inside. The journey from the idea of a solid, featureless sphere to our modern understanding of the nuclear atom is one of the most dramatic stories in science.

1803 • Dalton's Atomic Theory

John Dalton proposed that all matter is made of tiny, indivisible atoms. Each element has its own type of atom, and chemical reactions rearrange atoms without creating or destroying them.

1897 • Discovery of the Electron

J.J. Thomson used cathode ray tubes to discover the electron — a negatively charged particle much smaller than an atom. This proved atoms were not indivisible after all.

1904 • The Plum Pudding Model

Thomson proposed that an atom was a sphere of positive charge with electrons embedded in it, like plums in a pudding. This was the first model to include subatomic particles.

1911 • Rutherford's Gold Foil Experiment

Ernest Rutherford, Hans Geiger, and Ernest Marsden fired alpha particles at gold foil. Most passed through, but some bounced back — revealing a tiny, dense, positively charged nucleus.

1932 • Discovery of the Neutron

James Chadwick discovered the neutron — a neutral particle in the nucleus. This completed our picture of the three subatomic particles: protons, neutrons, and electrons.

The central question that drove this era of discovery was deceptively simple: What is the internal structure of an atom? Rutherford's experiment shattered the plum pudding model and introduced the concept of the nuclear atom — an atom with a dense, central nucleus surrounded by a cloud of electrons. Understanding this structure is the foundation for all of chemistry.

SECTION 2

Core Principles of the Nuclear Atom

The nuclear model of the atom rests on several key ideas that emerged from decades of experimental work. Every atom has three types of subatomic particles — protons, neutrons, and electrons — arranged in a specific way. Together, these particles determine the identity of an element, its mass, and its chemical behaviour.

The Nucleus

The nucleus is the tiny, dense centre of the atom. It contains protons (positive charge) and neutrons (no charge). Despite holding nearly all the atom's mass, the nucleus is only about 10⁻¹⁵ m across — roughly 100,000 times smaller than the atom itself.

Protons & Atomic Number

Each proton carries a charge of +1 and a relative mass of approximately 1 u. The number of protons defines the atomic number (Z), which determines the element's identity.

Neutrons & Mass Number

Neutrons have no electric charge and a relative mass of approximately 1 u. The total number of protons and neutrons gives the mass number (A). Atoms of the same element with different neutron counts are called isotopes.

Electrons

Electrons carry a charge of −1 and have negligible mass (about 1/1836 of a proton). They occupy the space around the nucleus in energy levels. In a neutral atom, the number of electrons equals the number of protons.

Mostly Empty Space

An atom is overwhelmingly empty space. If the nucleus were the size of a marble at centre-field in a football stadium, the nearest electrons would be orbiting in the upper stands. This vast emptiness is why most alpha particles passed straight through Rutherford's gold foil.

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KEY TAKEAWAY

Think of an atom like a sports stadium. The nucleus is a tiny marble sitting at centre-field — it's incredibly small but holds almost all the mass. The electrons are like a few bees buzzing around in the vast open air of the stadium. All the "stuff" is concentrated in the marble, and the rest is essentially empty space. That's why Rutherford was so surprised when some alpha particles bounced straight back — it was like firing a ball and having it ricochet off something you could barely see.

SECTION 3

Visualising the Nuclear Atom

The diagram below shows the nuclear model of the atom, illustrating the arrangement of protons, neutrons, and electrons. Notice the enormous difference in scale between the nucleus and the electron cloud — in reality, this difference is even more extreme than any diagram can capture.

A simplified nuclear atom model showing protons (red) and neutrons (violet) packed together in the nucleus, with electrons (cyan) occupying energy levels at much greater distances. In reality, the nucleus is approximately 100,000 times smaller than the atom.

In the diagram above, the nucleus is drawn much larger than its actual proportion, so that you can see the individual protons and neutrons. The dashed concentric circles represent electron energy levels (also called shells). Electrons closer to the nucleus have lower energy, while those further out have higher energy. Remember: the atom is mostly empty space between the tiny nucleus and the electrons.

SECTION 4

The Mathematical Framework

Even though the nuclear atom is a qualitative model, several important numerical relationships let us describe any atom precisely. The IB syllabus expects you to calculate the number of subatomic particles in atoms and ions using two key values: the atomic number and the mass number.

MASS NUMBER

A = Z + N

A = mass number (total protons + neutrons), Z = atomic number (number of protons), N = number of neutrons. Rearranged: N = A − Z.

NUCLIDE NOTATION

ᴬ_Z X

In nuclide notation, the mass number A is written as a superscript and the atomic number Z as a subscript, both to the left of the element symbol X. For example, carbon-12 is written as ¹²₆C.

ELECTRONS IN IONS

electrons = Z − charge

For a neutral atom, electrons = Z. For a positive ion (cation), subtract the charge. For a negative ion (anion), add the magnitude of the charge. For example, Na⁺ has 11 − 1 = 10 electrons, while Cl⁻ has 17 + 1 = 18 electrons.

These relationships may look simple, but they are the backbone of problems involving isotopes, ions, and nuclear reactions. Always start by identifying Z and A from the periodic table or the nuclide notation, and then work out the number of each particle.

SECTION 5

Subatomic Particles in Detail

The three subatomic particles differ in charge, mass, and location within the atom. The table below summarises the key properties that you need to know for the IB exam. Pay special attention to the relative masses and charges — the IB often uses these relative values rather than the actual (absolute) values.

Property	Proton	Neutron	Electron
Symbol	p⁺ (or ¹₁H)	n⁰ (or ¹₀n)	e⁻
Location	Nucleus	Nucleus	Electron shells around nucleus
Relative charge	+1	0	−1
Relative mass	≈ 1	≈ 1	≈ 1/1836 (negligible)
Actual mass (kg)	1.673 × 10⁻²⁷	1.675 × 10⁻²⁷	9.109 × 10⁻³¹

Summary of subatomic particle properties

The three isotopes of hydrogen all have one proton (Z = 1), but they differ in their number of neutrons (N = 0, 1, and 2). This changes the mass number A but not the chemical identity of the element.

The hydrogen isotopes shown above are a perfect example of how isotopes work. Protium, deuterium, and tritium are all hydrogen because they each contain exactly one proton. However, they have different numbers of neutrons, which gives them different mass numbers and slightly different physical properties. The chemical behaviour of isotopes is virtually identical because chemistry is governed by electrons, not neutrons.

SECTION 6

Worked Example: Identifying Subatomic Particles

Let's work through a typical IB-style problem step by step. The goal is to determine the number of protons, neutrons, and electrons in a specific species.

How many protons, neutrons, and electrons are in the ion ⁵⁶₂₆Fe²⁺?

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Step 1 — Identify Z and A from the notation

The nuclide notation ⁵⁶₂₆Fe tells us that the mass number A = 56 (superscript) and the atomic number Z = 26 (subscript). The symbol Fe confirms this is iron.

A = 56, Z = 26
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Step 2 — Calculate the number of protons

The atomic number directly equals the number of protons. Since Z = 26, the ion contains 26 protons. Note: the charge of an ion does NOT change the number of protons — only electrons are gained or lost.

Protons = 26
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Step 3 — Calculate the number of neutrons

Use the formula N = A − Z. Substituting: N = 56 − 26 = 30. The nucleus contains 30 neutrons.

Neutrons = 56 − 26 = 30
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Step 4 — Calculate the number of electrons

The species has a 2+ charge (Fe²⁺), which means it has lost two electrons compared to a neutral atom. A neutral iron atom has 26 electrons (equal to Z). Therefore, Fe²⁺ has 26 − 2 = 24 electrons.

Electrons = 26 − 2 = 24
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Step 5 — Summarise

The ion ⁵⁶₂₆Fe²⁺ contains 26 protons, 30 neutrons, and 24 electrons. The positive charge arises because there are two more protons than electrons, giving a net charge of +2.

⁵⁶₂₆Fe²⁺: 26 p⁺, 30 n⁰, 24 e⁻

SECTION 7

Comparing Atomic Models

The nuclear model was a massive improvement over earlier ideas, but it's important to understand both its strengths and limitations. Scientific models are always simplified representations of reality, and newer models build on older ones as evidence accumulates.

Model	Key Idea	Strengths	Limitations
Dalton (1803)	Atoms are solid, indivisible spheres unique to each element.	Explained conservation of mass and fixed ratios in compounds.	Could not explain electrical phenomena or the existence of subatomic particles.
Thomson (1904)	Electrons embedded in a uniform positive sphere ("plum pudding").	First model to include subatomic particles; explained electrical neutrality.	Disproved by Rutherford's experiment; no concentrated positive charge.
Rutherford (1911)	Dense positive nucleus with electrons in surrounding space.	Explained alpha scattering results; introduced the nuclear model.	Could not explain why electrons don't spiral into the nucleus or why atoms emit specific spectral lines.
Bohr (1913)	Electrons orbit the nucleus in quantised energy levels.	Explained hydrogen's line spectrum; introduced energy quantisation.	Only accurate for hydrogen; didn't explain multi-electron spectra.

Evolution of atomic models

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KEY TAKEAWAY

Think of each atomic model as a draft of a movie script. Dalton's script was a rough outline — it captured the basic plot (atoms exist, they combine in ratios) but left out the characters. Thomson added characters (electrons!) but got the setting wrong. Rutherford rewrote the setting (a dense nucleus at centre stage), and Bohr refined the script by giving each character specific stage directions (quantised orbits). Each draft is better than the last, but none is the final version — the quantum mechanical model is the current best draft, and it might be revised too.

SECTION 8

Connection to Advanced Theory

The nuclear atom model you've learned here is the starting point for more sophisticated ideas that you'll encounter later in the IB Chemistry course and beyond. Understanding the distinction between the simple nuclear model and the quantum mechanical model is important, because the IB expects you to know where each model applies and where it breaks down.

Feature	Nuclear Model (Rutherford)	Quantum Mechanical Model
Electron location	Electrons orbit like planets around the sun	Electrons exist in probability clouds (orbitals)
Energy levels	Not specified by Rutherford (Bohr added these)	Quantised energy levels with subshells (s, p, d, f)
Predictive power	Qualitative — explains scattering but not spectra	Quantitative — predicts spectra, bonding, and reactivity
Applicable atoms	All atoms (general structure)	All atoms (detailed electron behaviour)

Nuclear model vs. quantum mechanical model

As you move through the IB Chemistry course, you'll learn about electron configurations, where electrons fill orbitals in a specific order, and ionisation energy trends, which provide experimental evidence for energy levels. The nuclear model you've studied here gives you the essential vocabulary — protons, neutrons, electrons, atomic number, mass number, and isotopes — that all of these advanced topics build upon.

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IB Exam Tip

IB examiners frequently test whether you understand that scientific models evolve as new evidence is gathered. Be prepared to explain why a model was replaced (what evidence contradicted it) and what the new model explained that the old one could not. This connects to the Nature of Science (NOS) theme.

SECTION 9

Practice Problems

Question 1: CONCEPTUAL

In Rutherford's gold foil experiment, most alpha particles passed straight through the foil, a small number were deflected at large angles, and very few bounced straight back. Explain what each of these three observations tells us about the structure of the atom.

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Most alpha particles passing straight through shows that atoms are mostly empty space — there is very little solid matter to block them. The small number deflected at large angles indicates that the positive charge and mass are concentrated in a very small region (the nucleus), which repels the positively charged alpha particles. The very few that bounced back must have collided nearly head-on with this dense nucleus. Together, these observations disproved the plum pudding model and established the nuclear model of the atom.

Question 2: BASIC CALCULATION

Determine the number of protons, neutrons, and electrons in the atom ³⁹₁₉K.

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The atomic number Z = 19 tells us that potassium has 19 protons. Since the atom is neutral (no charge is given), the number of electrons equals the number of protons: 19 electrons. The mass number A = 39 gives us the total of protons plus neutrons, so the number of neutrons is N = A − Z = 39 − 19 = 20 neutrons. Therefore, ³⁹₁₉K has 19 protons, 20 neutrons, and 19 electrons.

Question 3: INTERMEDIATE

Two species are represented as ⁴⁰₂₀Ca and ⁴⁰₁₈Ar. These species have the same mass number. Are they isotopes of each other? Explain your reasoning, and compare the number of each subatomic particle in the two atoms.

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No, these are not isotopes of each other. Isotopes must be atoms of the same element (same atomic number Z), differing only in their number of neutrons. Calcium-40 has Z = 20, so it is calcium, while argon-40 has Z = 18, so it is argon — they are different elements. Atoms with the same mass number but different atomic numbers are called isobars, not isotopes. Ca-40 has 20 protons, 20 neutrons (40 − 20), and 20 electrons. Ar-40 has 18 protons, 22 neutrons (40 − 18), and 18 electrons.

Question 4: APPLIED

A forensic scientist analyses a metal sample and determines it contains ions with 26 protons, 30 neutrons, and 23 electrons. Write the full nuclide notation for this ion, including its charge. Identify the element and explain how you determined the charge.

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The element with 26 protons is iron (Fe), since the atomic number Z = 26 defines the element. The mass number A = protons + neutrons = 26 + 30 = 56. In a neutral atom of iron there would be 26 electrons, but this ion has only 23 electrons. It has lost 3 electrons (26 − 23 = 3), giving it a charge of +3. The full nuclide notation is ⁵⁶₂₆Fe³⁺. This is the iron(III) ion, commonly found in rust and many biological molecules like haemoglobin.

Question 5: CRITICAL THINKING

Thomson's plum pudding model predicted that alpha particles would pass through gold foil with only minor deflections because the positive charge was spread thinly across the atom. Rutherford's experiment showed that a few particles were deflected by more than 90°. Explain why the plum pudding model could not account for these large-angle deflections, and discuss why it took a new model (rather than a modification of the old one) to explain the results.

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In the plum pudding model, positive charge is spread uniformly throughout a relatively large sphere. A positively charged alpha particle passing through such a diffuse charge distribution would experience weak, evenly distributed repulsive forces from all directions — these forces would largely cancel out, producing only slight deflections. Large-angle deflections (greater than 90°) require an extremely strong repulsive force concentrated in a very small region. This is only possible if the positive charge and most of the mass are packed into a tiny volume — the nucleus. The plum pudding model couldn't be simply tweaked because its fundamental assumption (uniform charge distribution) is incompatible with the observation. Rutherford had to propose a structurally different model where nearly all the mass and all the positive charge are concentrated in a nucleus roughly 10⁵ times smaller than the atom itself. This is an example of a paradigm shift in science — sometimes evidence requires an entirely new framework rather than a patch on the old one.

SUMMARY

Lesson Summary

Summary

The nuclear atom consists of a tiny, dense nucleus containing positively charged protons and neutral neutrons, surrounded by negatively charged electrons in energy levels. The atomic number (Z) equals the number of protons and defines the element, while the mass number (A) equals the total of protons plus neutrons. The number of neutrons is calculated as N = A − Z.

Isotopes are atoms of the same element with different numbers of neutrons (same Z, different A). Ions form when atoms gain or lose electrons, changing the electron count but not the proton or neutron count. Rutherford's gold foil experiment provided the critical evidence for the nuclear model by showing that most of the atom is empty space with a concentrated positive core. Scientific models evolve as new evidence emerges — from Dalton's solid sphere, to Thomson's plum pudding, to Rutherford's nuclear atom, each model improved our understanding of atomic structure.