Structure of Atoms

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A route into the idea

From 'solid sphere' to 'nucleus' to 'probability cloud' — what each model explained

Identify the experimental anomaly
Name the model that resolved it
State what the model could NOT explain
Name the next model that addressed the gap

Atomic models are not a ladder of truth — they are a sequence of better approximations. Each model solved a specific puzzle (multiple proportions, cathode rays, scattering, line spectra) but created new questions. The quantum mechanical model is the current best fit; it will be refined, not replaced.

Try the workshop →

Try an idea before you read. Explore how each model answered a question and created a new one. Make a prediction before opening each section. Explore the discovery →

Introduction

Matter is made up of fundamental building blocks, that is atoms and molecules. The existence of different kinds of matter is due to different atoms constituting them.

Are atoms really indivisible, as proposed by Dalton?

One of the first indication that atoms are not indivisible (means divisible), comes from studying static electricity and the conditions under which electricity is conducted by different substances.

Charged Particles in Matter

From different experiments, it was discovered that an atom is divisible and consists of charged particles.

By 1900, J. J. Thomson discovered a sub-atomic particle called an electron.

Even before the electron was identified, E. Goldstein in 1886 discovered the presence of new radiations in a gas discharge and called them canal rays. These rays were positively charged radiations which ultimately led to discovery of another sub-atomic particle called proton.

The electron is represented as ‘e’ while the proton as ‘p’.

The mass of proton is taken as one unit and its charge as plus one, whereas the mass of an electron is considered to be negligible and its charge is minus one.

Neutrons

In 1932, J. Chadwick discovered another sub-atomic particle which had no charge and a mass nearly equal to that of a proton. It was named as Neutrons. In general neutron is represented as ‘n’.

Neutrons are present in the nucleus of all the atoms, expect hydrogen.

The mass of an atom is given by the sum of the masses of protons and neutrons present in the nucleus.

Particle	Electron	Proton	Neutron
Symbol	e	p	n
Relative Charge	-1	+1	0
Nature	Negatively Charged	Positively Charged	Neutral
Discovery	JJ Thomson	E. Goldstein	Chadwick

The Structure of an Atom

Dalton’s atomic theory suggested that the atom is indivisible and indestructible. But the discovery of the two fundamental particles (electrons and protons) inside the atom, led to the failure of this aspect of Dalton’s atomic theory.

So, it became necessary to know how the electrons and protons are arranged within an atom. For explaining this, many model for the structure of atom were proposed.

(A) Thomson’s Model of an Atom

Thomson proposed the model of an atom to be similar to that of a Plum pudding (Christmas pudding) or Watermelon.

Thomson model is compared with watermelon as shown in the figure:

The positive charge in the atom is spread all over like the red edible part of the watermelon.
While the electrons are studded in the positively charged sphere, like the seeds in the watermelon.

Postulates of Thomson’ model of an Atom:

An atom consists of a positively charged sphere and the electrons are embedded in it.
The negative and positive charges are equal in magnitude. So, the atom as a whole is electrically neutral.

Although, Thomson’s model explained that atoms are electrically neutral, but it failed to explain the stability of an atom. This theory also fails to accounts for the position of neutrons in the atom.

There are no experimental evidences in its supports.

(B) Rutherford’s Model of an Atom

Ernest Rutherford was interested in knowing how the electrons are arranged within an atom. For this Rutherford designed an experiment, in which fast moving alpha, particles were made to fall on a thin gold foil.

He selected a gold foil because he wanted as thin a layer as possible. This gold foil was about 1000 atoms thick.
Alpha-particles are doubly charged helium ions. Since they have a mass of 4 u, the fast moving alpha - particles have a considerable amount of energy.
It was expected that alpha - particles would be deflected by the sub-atomic particles in the gold atoms. Since the alpha - particles were much heavier than the protons, he did not expect to see large deflections.

But, the alpha - particles scattering experiments gave totally unexpected results. The following were the observations:

(a) Most of the fast moving alpha - particles passed straight through the gold foil.

(b) Some of the alpha - particles were deflected by the foil by small angles.

Conclusion from alpha, particles scattering experiment are:

Most of the space inside the atom is empty because most of the alpha - particles passed through the gold foil without getting deflected.
Very few particles were deflected from their path, indicating that the positive charge of the atom occupies very little space.
A very small fraction of alpha - particles were deflected by 180o, indicating that all the positive charge and mass of the gold atom were concentrated in a very small volume within the atom.

From the data, it is was concluded that the radius of the nucleus is 105 times less than the radius of the atom.

On the basis of experiment, Rutherford put forward the nuclear model of an atom, which had the following features:

There is positively charged centre in an atom called the nucleus. Nearly all the mass of an atom resides in the nucleus.
The electrons revolve around the nucleus in circular paths.
The size of the nucleus is very small as compared to the size of the atoms.

Drawbacks of Rutherford’s model of the atom.

The revolution of the electron in a circular orbit is not expected to be stable. Any particle in a circular orbit would undergo acceleration radiating energy and would lose energy and finally fall into the nucleus. If this were so, the atom should be highly unstable and hence the matter would not exist in the form that we know.
Fails to explain the arrangement of the electrons in an atom.

(C) Bohr’s Model of an Atom

In order to overcome the objective raised against Rutherford’s model of an atom, Neil Bohr put forward the following postulates of the model of an atom:

Only certain special orbits known as discrete orbit of electrons are allowed inside the atom.
While revolving in discrete orbits the electrons do not radiate energy.

These orbits or shells are called energy levels. These orbits or shells are represented by the letters K, L, M, N, … or the numbers, n = 1, 2, 3, 4, …

Summary of three different models of an Atom:

How are electrons distributed in different orbits (shells)?

The distribution of electrons into different orbits was suggested by Bohr and Bury. The following rules are followed for writing the number of electrons in different energy levels or shells:

The maximum number of electrons present in a shell is given by the formula

2n2, where n = orbit number or energy level index,

n = 1, 2, 3, ……

Hence maximum number of electrons in different shells are as follows:

First orbit or K-shell = 2 × 12 = 2

Second orbit or L-shell = 2 × 22 = 8

Third orbit or M-shell = 2 × 23 = 18

Fourth orbit or N-shell = 2 × 24 = 32 and so on …

The maximum number of electrons that can be accommodated in the outermost orbit is 8.

Electrons are not accommodated in a given shell, unless the inner shells are filled. That is, the shells are filled in a step-wise manner.

Atomic structure of the first eighteen elements:

Valency

Valency is the combining capacity of an atom.
The electrons present in the outermost shell of an electrons are known as the valence electrons.
The outermost shell of an atom can be accommodated with a maximum of 8 electrons; hence the shell is completely filled and show little chemical activity. In other words, their combining capacity is zero.
Of these inert elements, the helium atom has 2 electrons in its outermost shell and all other elements (neon, argon, krypton, xenon and radon) have atoms with 8 electrons in the outermost shell.
The combining capacity of the atoms of elements, ie., their tendency to react and form molecules with atoms of the same or different elements, can be explained based on their attempt to attain a fully-filled outermost shell

An outermost shell, which had eight electrons was said to possess an octet. This is done by sharing, gaining or losing electrons.
The number of electrons shared, gained or lost so as to make the octet of electrons in the outermost shell, gives the combining capacity of the element, i.e., valency.
For example, hydrogen/ lithium/ sodium atoms contain one electron each in their outermost shell, therefore each one of them can lose one electron. So, they have valence of one.
If the number of electrons in the outermost shell of an atom is close to its full capacity, then valency is determined in a different way.
For example, the fluorine atom has 7 electrons in the outermost shell and so its valency could be 7. But for fluorine to gain one electron is easier instead of losing 7 electrons. Hence, its valency is determined by subtracting seven electrons from the octet and thus its valency would be one.

Atomic Number and Mass Number

Atomic Number:

Atomic number is defined as the total number of protons present in the nucleus of an atom. It is denoted by ‘Z’.

All atoms of an element have the same atomic number, Z. in fact the elements are defined by the number of protons they possess.

For example, for Hydrogen, Z = 1, because it possesses only one electron.

Similarly, for Carbon, Z = 6, as it possesses six protons.

Mass Number:

The mass number is defined as the sum of total number of protons and neutrons present in the nucleus of an atom. It is denoted by ‘A’.

The protons and neutrons together are called nucleons.

For example, mass of carbon is 12 u, i.e., 6 protons + 6 neutrons, 6 u + 6 u = 12 u

Similarly, mass of aluminium is 27 u, i.e., 13 protons + 14 neutrons.

In the notation for an atom, the atomic number, mass number and symbol of the element are to be written as:

For example, nitrogen is written as

Isotopes:

Isotopes are defined as the atoms of the same element, having the same atomic number but different mass numbers.

The chemical properties of isotopes are similar but their physical properties are different.

Many elements consist of a mixture of isotopes. Each isotope of an element is a pure substance.

For example, Chlorine occurs in two isotopic forms in nature, with masses 35 u and 37 u in the ratio of 3:1.

The average mass is taken of all the naturally occurring atoms of that element. If an element has no isotope, then the mass of its atom is the sum of protons and neutrons in it. But if an element occurs in isotopic form, then percentage of each isotopic from is to be known and then average mass is calculated.

So, for certain amount of chlorine taken it will contain both the isotopes of chlorine and the average mass is 35.5 u.

Application:

(i) An isotope of Uranium is used as a fuel in nuclear reactors.

(ii) An isotope of cobalt (Co60) is used in the treatment of cancer.

(iii) An isotope of iodine is used in the treatment of goitre.

Isobars:

Atoms of different elements with different atomic numbers, which have the same mass number, are known as isobars. Or

The total number of nucleons (protons + neutrons) is the same in the atoms.

Key takeaways

Atoms are not indivisible; they are complex structures composed of protons, neutrons, and electrons.
Rutherford's gold foil experiment proved that atoms are mostly empty space, containing a dense, positively charged nucleus.
Bohr's model introduced discrete energy shells (K, L, M, N) where electrons reside without radiating energy and collapsing.
The atomic number defines an element via its proton count, while the mass number is the sum of protons and neutrons.
Valence electrons in the outermost shell dictate an atom's chemical reactivity and its drive to form bonds.

Test yourself

What did Rutherford's gold foil experiment prove?

It proved that atoms consist mostly of empty space with a tiny, incredibly dense, positively charged nucleus at the center.

How do isotopes differ from one another?

Isotopes are atoms of the same element that have the same number of protons but a different number of neutrons, resulting in different mass numbers.

What is the maximum number of electrons the innermost shell (K-shell) can hold?

The K-shell can hold a maximum of 2 electrons.

Play with the idea

Structure of the atom: test the model, count the particles, write the decay

These scenarios explore the evolution of atomic models, subatomic particles, atomic structure notation, electron distribution, and radioactivity. Each question uses fictional but structurally accurate experimental evidence. The test is whether you can connect the evidence to the correct model or calculation.

Situation 1

In the gold foil experiment, most α-particles pass through with little deflection, but ~1 in 8000 deflects at angles > 90°. A student claims: 'This proves atoms are mostly empty space with a tiny dense nucleus.' Which assessment does the evidence support?

Correct — the rare large-angle scattering requires a concentrated positive charge; the empty space explains why most pass through.Incorrect — the deflection could be caused by multiple small scatterings in the plum pudding.Incomplete — the experiment does not tell us about electron arrangement.

Explore the reasoning for every approach

Correct — the rare large-angle scattering requires a concentrated positive charge; the empty space explains why most pass through.

Yes. Thomson's plum pudding model (diffuse positive charge) could not produce large-angle deflections. The observation that most α-particles pass through (empty space) and a few scatter at >90° (tiny dense nucleus) is exactly what Rutherford's nuclear model predicts. The nucleus occupies ~1/100,000 of atomic volume but contains >99.9% of mass.

Incorrect — the deflection could be caused by multiple small scatterings in the plum pudding.

Multiple small scatterings would produce a distribution of small angles, not the observed rare large-angle events. The probability of many small scatterings aligning to produce >90° deflection is negligible. Single scattering from a concentrated charge is the only viable mechanism.

Incomplete — the experiment does not tell us about electron arrangement.

True but irrelevant to the claim. The claim is about the nucleus and empty space, which the evidence directly supports. The electron arrangement was a separate question (addressed later by Bohr and quantum mechanics).

Situation 2

Chlorine has two stable isotopes: ³⁵Cl (75.78% abundance) and ³⁷Cl (24.22% abundance). A student calculates the average atomic mass as (35 + 37) / 2 = 36 u. Which assessment does the evidence support?

Correct — the average of the two mass numbers is 36.Incorrect — the weighted average by abundance gives ~35.5 u, not 36 u.Cannot determine — need the exact isotopic masses, not just mass numbers.

Explore the reasoning for every approach

Correct — the average of the two mass numbers is 36.

Incorrect. The average atomic mass is a WEIGHTED average by natural abundance, not a simple average. Correct calculation: (35 × 0.7578) + (37 × 0.2422) = 26.523 + 8.9614 = 35.484 u ≈ 35.5 u. The simple average assumes equal abundance, which is false.

Incorrect — the weighted average by abundance gives ~35.5 u, not 36 u.

Yes. The periodic table lists 35.5 u for chlorine because the more abundant isotope (³⁵Cl at 75.78%) pulls the average closer to 35. The student's simple average ignores abundance data. This is why atomic masses are rarely whole numbers.

Cannot determine — need the exact isotopic masses, not just mass numbers.

Mass numbers (35, 37) are close enough for Class 9 precision. Actual isotopic masses are 34.9689 u and 36.9659 u, giving 35.45 u. The principle — weighted average by abundance — is what matters. The student's error is using unweighted average.

Situation 3

A radioactive nuclide ²¹⁰₈₄Po decays by α-emission to ²⁰⁶₈₂Pb. Another nuclide ¹⁴₆C decays by β⁻-emission to ¹⁴₇N. A student writes the equations: (1) ²¹⁰₈₄Po → ²⁰⁶₈₂Pb + ⁴₂α. (2) ¹⁴₆C → ¹⁴₇N + ⁰₋₁β. Are both equations balanced?

Yes — both conserve mass number and atomic number.No — β-decay should decrease Z because an electron is lost.No — the α-particle should be written as ⁴₂He, not ⁴₂α.

Explore the reasoning for every approach

Yes — both conserve mass number and atomic number.

Yes. Equation (1): A: 210 = 206 + 4 ✓; Z: 84 = 82 + 2 ✓. Equation (2): A: 14 = 14 + 0 ✓; Z: 6 = 7 + (−1) ✓. Both are correctly balanced. α-decay reduces A by 4, Z by 2. β⁻-decay keeps A same, increases Z by 1 (neutron → proton + electron).

No — β-decay should decrease Z because an electron is lost.

Incorrect. In β⁻-decay, a neutron in the nucleus transforms: n → p + e⁻. The proton stays (Z increases by 1); the electron is ejected. The atomic number of the daughter nucleus INCREASES. The emitted electron is not an orbital electron — it comes from the nucleus.

No — the α-particle should be written as ⁴₂He, not ⁴₂α.

Both notations are accepted. ⁴₂α emphasises it's an α-particle; ⁴₂He emphasises it's a helium nucleus. They are the same particle (2 protons, 2 neutrons). The equation is balanced either way.

Investigate before you memorise

Structure of the atom: from solid spheres to probability clouds

Explore how each model answered a question and created a new one. Make a prediction before opening each section.

Open the model evolution timeline

Five models — each replaced the previous by explaining new evidence

No model is "wrong" — each was the best fit for its evidence. The quantum mechanical model is current but not final.

Model	Description	Key Evidence	Limitation	Diagram
Dalton (1803)	Solid indivisible spheres	Law of multiple proportions	No internal structure; cannot explain electricity, radioactivity	Solid sphere
Thomson (1897)	Plum pudding — positive sphere with embedded electrons	Cathode rays (electrons); charge-to-mass ratio	Cannot explain Rutherford scattering; no nucleus	Positive sphere with negative electrons inside
Rutherford (1911)	Nuclear model — tiny dense positive nucleus, electrons orbit	Gold foil experiment: most α-particles pass through, few deflect at large angles	Orbiting electrons should radiate energy and spiral in (classical EM); cannot explain line spectra	Central nucleus, electrons in orbits
Bohr (1913)	Quantised orbits — electrons in fixed energy levels, jump by absorbing/emitting photons	Hydrogen line spectrum (Balmer series); Rydberg formula	Only works for hydrogen/one-electron systems; violates uncertainty principle	Nucleus, discrete electron shells (K, L, M...)
Quantum Mechanical (1926–)	Electrons as probability clouds (orbitals); no fixed paths	Schrödinger equation; Heisenberg uncertainty; electron diffraction	Mathematically complex; visualisation requires analogy	Nucleus, orbital shapes (s, p, d, f)

Why did Rutherford reject the plum pudding model?

In the gold foil experiment, most α-particles passed through with little deflection (atom is mostly empty space), but ~1 in 8000 deflected at angles > 90°. Thomson's diffuse positive charge could not produce such large deflections — only a tiny, dense, massive positive centre (nucleus) could. The nucleus contains >99.9% of the atom's mass in ~1/100,000 of its volume.

Open the subatomic particle evidence

Electron, proton, neutron — how we found them

Each particle was discovered by a specific experiment. The evidence trail: cathode rays → nucleus → neutron.

Particle	Discoverer	Charge	Mass	Location	Key Experiment
Electron (e⁻)	J.J. Thomson (1897)	−1 (relative)	9.11×10⁻³¹ kg (1/1837 u)	Outside nucleus (orbitals)	Cathode ray tube; deflection by electric/magnetic fields
Proton (p⁺)	E. Rutherford (1917)	+1	1.67×10⁻²⁷ kg (1.007 u)	Inside nucleus	α-particle scattering; hydrogen nucleus identified as fundamental particle
Neutron (n⁰)	J. Chadwick (1932)	0	1.67×10⁻²⁷ kg (1.009 u)	Inside nucleus	Be-9 + α → C-12 + n; neutral particle with proton-like mass

How did Chadwick prove the neutron exists?

Chadwick bombarded beryllium-9 with α-particles. A neutral radiation emerged that could knock protons out of paraffin wax. The energy transfer matched a neutral particle with mass ≈ proton. It was not γ-rays (which would not transfer so much momentum). The reaction: ⁹₄Be + ⁴₂α → ¹²₆C + ¹₀n.

Open the atomic structure decoder

Z, A, isotopes, isobars, isotones — the naming system

Every nuclide is defined by two numbers: Z (protons) and A (protons + neutrons). The third (neutrons = A − Z) is derived.

Term	Definition	Symbol	Example
Atomic number (Z)	Number of protons in nucleus = number of electrons in neutral atom	Z	Carbon: Z = 6 (6 protons, 6 electrons)
Mass number (A)	Total protons + neutrons in nucleus	A	Carbon-12: A = 12 (6p + 6n); Carbon-14: A = 14 (6p + 8n)
Isotopes	Same Z, different A (different neutron number)		¹²C, ¹³C, ¹⁴C — all carbon, different masses
Isobars	Different Z, same A		⁴⁰Ca (Z=20) and ⁴⁰Ar (Z=18) — both A=40
Isotones	Same neutron number, different Z		¹⁴C (6p, 8n) and ¹⁶O (8p, 8n) — both 8 neutrons
Nuclide notation	ᴬᶻX (X = element symbol, Z = atomic number, A = mass number)		²³⁵₉₂U, ¹⁴₆C, ⁴₀₁₈Ar

Why do isotopes have similar chemical properties but different physical properties?

Chemical properties depend on electrons (Z), which are identical for isotopes. Physical properties (mass, density, diffusion rate, boiling point) depend on nuclear mass (A). Heavy water (D₂O, where D = ²H) boils at 101.4 °C vs 100 °C for H₂O. Radioactive isotopes decay; stable ones do not.

Open the electron distribution rules

Shells, orbitals, and the octet

Electrons fill by energy (Aufbau), not just shell number. The 4s fills before 3d. Valence electrons drive chemistry.

Rule	Formula / Detail	K Shell (n=1)	L Shell (n=2)	M Shell (n=3)	N Shell (n=4)
Maximum electrons per shell (Bohr-Bury)	2n² where n = shell number (1, 2, 3, 4...)	n=1 (K): 2	n=2 (L): 8	n=3 (M): 18	n=4 (N): 32
Aufbau principle	Fill lowest energy orbitals first (1s, 2s, 2p, 3s, 3p, 4s, 3d...)	1s²	2s² 2p⁶	3s² 3p⁶ 4s² 3d¹⁰	4p⁶ 5s² 4d¹⁰...
Valence electrons	Electrons in outermost shell; determine chemical properties	Noble gases (except He): 8	Group 1: 1; Group 2: 2; Group 13: 3...
Octet rule	Atoms tend to gain/lose/share to achieve 8 valence electrons (2 for He)	He: 2 (duplet)	Ne, Ar, Kr...: 8

Why does calcium (Z=20) have configuration 2,8,8,2 not 2,8,10?

The 4s orbital is lower in energy than 3d. After argon (2,8,8), the next electron goes into 4s, not 3d. Calcium: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s². The 3d fills after 4s (scandium onwards). The Bohr-Bury 2n² rule gives maximum capacity, not filling order. Energy order: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p...

Open the radioactivity lab

Alpha, beta, gamma — three radiations, three penetrations

Each type is a different particle/wave with distinct charge, mass, and penetration. They change Z and A differently.

Type	Description	Charge	Mass	Penetration	Nuclear Equation
Alpha (α)	Helium nucleus (2p+2n)	+2	4 u	Stopped by paper	²³⁸₉₂U → ²³⁴₉₀Th + ⁴₂α
Beta (β)	Fast electron (n → p + e⁻)	−1	~0	Stopped by aluminium foil	¹⁴₆C → ¹⁴₇N + ⁰₋₁β
Gamma (γ)	High-energy photon	0	0	Stopped by thick lead/concrete	Often accompanies α/β decay; nucleus loses energy

Why does β-decay increase Z by 1 but leave A unchanged?

A neutron in the nucleus transforms into a proton + electron (β⁻). The proton stays in nucleus (Z increases by 1); the electron is ejected at high speed. Total nucleons (A) unchanged. Example: ¹⁴₆C → ¹⁴₇N + ⁰₋₁β. The new element has Z+1, same A.

Open the identification game

Identify the particle, the model, or the nuclide

Q: A particle with charge +1, mass ~1 u, found in nucleus. Prediction: Proton.
Q: Neutral particle, mass ~1 u, discovered 1932. Prediction: Neutron.
Q: Nuclide with Z=17, A=35. How many neutrons? Prediction: 35 − 17 = 18 neutrons. It's ³⁵₁₇Cl.
Q: Element X has isotopes of mass 35 and 37 in 3:1 ratio. Average atomic mass? Prediction: (35×3 + 37×1) / 4 = 35.5 u.
Q: α-particle emitted by ²²⁶₈₈Ra. What is the daughter nuclide? Prediction: ²²²₈₆Rn (Z−2, A−4).
Q: β⁻ emitted by ¹³¹₅₃I. What is the daughter? Prediction: ¹³¹₅₄Xe (Z+1, A unchanged).

Why is the electron's mass ignored in mass number calculations?

Electron mass is 1/1837 u ≈ 0.0005 u. For any atom, the mass of all electrons combined is < 0.1% of total mass. Mass number counts nucleons (protons + neutrons) only. Atomic mass (in u) includes electrons but is still dominated by the nucleus.

Based on NCERT Class 9 Science Chapter 4: Structure of the Atom. Game scenarios and challenge questions are original teaching examples.

Try it

Structure of Atoms - Class 9 Science Notes

Test your understanding of atomic structure with this interactive quiz.

1Rutherford described his gold foil experiment result as 'almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you.' Why was this result so surprising?

2Carbon-12 has 6 protons and 6 neutrons. Carbon-14 (used in radiocarbon dating) also has 6 protons but 8 neutrons. Based on the text, what term best describes the relationship between these two atoms?

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