Thomson's Model · Rutherford's Experiment · Bohr's Shells · Isotopes · Isobars
Everything that you see, observe, or feel around is matter. You have learnt that matter consists of tiny particles called atoms. Both living beings (like us) and non-living things (like a house) are ultimately composed of atoms. These atoms are so tiny that they cannot be seen with the naked eye.
Is an atom truly the smallest unit of matter, or can it be divided even further? Scientists have been exploring whether atoms are divisible. If so, what are their constituents and how are these arranged?
A human body is made of organs, which contain tissues, which are built from cells, which are made of molecules -- all the way down to atoms!
Let us go back more than 2,000 years to ancient India and ancient Greece. Profound thinkers pondered over the same fundamental question: What is everything made up of?
Suggested that if matter (dravya) is divided repeatedly, you will reach the smallest particles that can no longer be divided. He called them parmanus. His ideas are recorded in the Vaisesika Sutras.
Proposed a similar idea. They called these indivisible particles atomos (Greek for "indivisible").
Concept of atoms as imaginary ideas, not based on experiments.
John Dalton proposed that all matter is composed of indivisible particles called atoms -- the first scientific description based on experiments. Atoms are the fundamental building blocks that cannot be broken down into smaller parts.
Until the late 19th century, atoms were thought to be the smallest, indivisible units. However, scientists discovered that certain elements emit radiation -- invisible energy and particles. This is called radioactivity, proving atoms were NOT indivisible!
In 1897, J. J. Thomson studied the conduction of electric current through gases at very low pressure. He used a glass tube with two electrodes (a cathode ray tube) and applied a high voltage. He observed rays moving from the cathode (negative electrode) to the anode (positive electrode) -- called cathode rays.
Electrons are a fundamental component of all atoms, present in every element. The charge of an electron (−1.602 × 10⁻¹⁹ C) is taken as −1.
Thomson faced a puzzle: atoms are neutral, so where is the positive charge? He proposed the atom to be a sphere of positive charge with electrons distributed throughout it.
Atom is a sphere of uniformly spread positive charge
Electrons scattered throughout like plums in a pudding
Red pulp = positive charge, Seeds = electrons
In 1911, Geiger and Marsden, working under Ernest Rutherford, tested Thomson's model through the famous gold foil experiment (also called the alpha-ray scattering experiment).
They aimed a narrow beam of alpha particles (tiny, positively charged particles from radioactive elements) at an extremely thin sheet of gold foil.
Alpha particles should pass straight through or be deflected only slightly (since positive charge is spread evenly).
From the experiment, Rutherford concluded:
The positive charge is NOT spread throughout but is concentrated in an extremely small, dense region called the nucleus. It contains all the positive charge and most of the mass.
Most of an atom is empty space, which is why most alpha particles passed through without deflection.
Electrons revolve around the nucleus, like planets orbiting the Sun. Hence: the planetary model.
The nucleus is about 10⁵ (one lakh) times smaller than the atom!
Diameter of atom ≈ 10⁻¹⁰ m, Diameter of nucleus ≈ 10⁻¹⁵ m
Imagine: If an atom were the size of a cricket ground (~100 m across), the nucleus would be just a tiny black pepper grain (a few mm) at the centre!
If a negatively charged electron keeps accelerating around the nucleus (changing direction constantly), it should lose energy, spiral inward, and eventually fall into the nucleus. Atoms would collapse! But in reality, atoms are stable. Rutherford's model could not explain atomic stability.
Rutherford showed that the nucleus carries positive charge from particles called protons. Protons are much heavier than electrons and possess a charge equal and opposite to that of electrons (+1). For an atom to be electrically neutral, the number of protons must equal the number of electrons.
To explain why atoms are stable, Niels Bohr proposed a new model of the atom in 1913.
Electrons do not move randomly but follow fixed circular paths called stationary states, orbits, or shells. Each shell has a definite energy, so they are also called energy levels.
Shells are represented by letters K, L, M, N, ... or numbers n = 1, 2, 3, 4, ...
While moving in a fixed shell, an electron does not lose energy. This explains atomic stability!
K-shell (closest to nucleus) has least energy. Energy increases as we move further: K < L < M < N.
An electron can move to another shell by absorbing or releasing a fixed amount of energy equal to the difference between two levels.
Each shell can hold only a certain maximum number of electrons.
Atom as indivisible particle
Plum pudding: +ve sphere with embedded electrons
Nuclear / Planetary model with nucleus at centre
Electrons in fixed energy levels / shells
Quantum mechanics, electron clouds
Rutherford's model showed that most of an atom's mass is in the nucleus. But a puzzle remained: helium has 2 protons yet is 4 times heavier than hydrogen (which has 1 proton). So something else must be adding mass!
In 1932, James Chadwick (a student of Rutherford) discovered a new subatomic particle with mass nearly equal to a proton but no electrical charge. This neutral particle was named the neutron (symbol: n). Neutrons are found in the nucleus of all atoms except hydrogen.
Charge: −1
Mass: Negligible
Location: Orbits / Shells
Charge: +1
Mass: ~1 u
Location: Nucleus
Charge: 0
Mass: ~1 u
Location: Nucleus
| Subatomic Particle | Symbol | Relative Charge | Location |
|---|---|---|---|
| Electron | e⁻ | −1 | Shells around nucleus |
| Proton | p⁺ | +1 | Inside the nucleus |
| Neutron | n | 0 | Inside the nucleus |
John Dalton introduced the first pictorial symbols for elements in 1803. Later, in 1813, Berzelius suggested that symbols should be derived from Latin names. Today, IUPAC (International Union of Pure and Applied Chemistry) approves names and symbols.
The number of protons in the nucleus of an atom is its atomic number, designated by Z. It determines the identity of an element and its chemical behaviour. Since the atom is neutral, the number of protons equals the number of electrons.
Elements with different atomic numbers are distinct elements. The atomic number uniquely identifies an element.
The total number of protons + neutrons in the nucleus is the mass number (A). Protons and neutrons in the nucleus are collectively called nucleons.
Mass number (A) = Number of protons (Z) + Number of neutrons (n)
Number of neutrons = A − Z
| Element | Protons (p⁺) | Neutrons (n) | Mass Number (A) |
|---|---|---|---|
| Hydrogen | 1 | 0 | 1 |
| Helium | 2 | 2 | 4 |
| Lithium | 3 | 4 | 7 |
| Carbon | 6 | 6 | 12 |
| Oxygen | 8 | 8 | 16 |
| Sodium | 11 | 12 | 23 |
An element is written as: AZX (mass number on top, atomic number on bottom)
Example: Carbon = 126C (6 protons, 6 neutrons, mass number 12)
Solution:
Atomic number (Z) = 26, so protons = 26 and electrons = 26
Mass number (A) = 56, so neutrons = A − Z = 56 − 26 = 30
This element is Iron (Fe)!
Bohr and Bury suggested the following rules for electron distribution:
The maximum electrons in the outermost shell is always 8 (except for the first shell which can hold max 2).
Electrons fill shells in order: K first, then L, then M, then N. The L-shell is filled only after K is complete, and so on.
| Element | Symbol | Z | K | L | M |
|---|---|---|---|---|---|
| Hydrogen | H | 1 | 1 | - | - |
| Helium | He | 2 | 2 | - | - |
| Lithium | Li | 3 | 2 | 1 | - |
| Beryllium | Be | 4 | 2 | 2 | - |
| Boron | B | 5 | 2 | 3 | - |
| Carbon | C | 6 | 2 | 4 | - |
| Nitrogen | N | 7 | 2 | 5 | - |
| Oxygen | O | 8 | 2 | 6 | - |
| Fluorine | F | 9 | 2 | 7 | - |
| Neon | Ne | 10 | 2 | 8 | - |
| Sodium | Na | 11 | 2 | 8 | 1 |
| Magnesium | Mg | 12 | 2 | 8 | 2 |
| Aluminium | Al | 13 | 2 | 8 | 3 |
| Silicon | Si | 14 | 2 | 8 | 4 |
| Phosphorus | P | 15 | 2 | 8 | 5 |
| Sulfur | S | 16 | 2 | 8 | 6 |
| Chlorine | Cl | 17 | 2 | 8 | 7 |
| Argon | Ar | 18 | 2 | 8 | 8 |
The number of atoms of hydrogen or chlorine with which one atom of an element can combine is called its combining capacity (valency). The outermost shell is the valence shell, and its electrons are valence electrons.
An atom with 8 electrons in its outermost shell (or 2 for helium) is stable and largely unreactive. This is called an octet.
Atoms with incomplete valence shells are usually more reactive. They lose, gain, or share electrons to complete their octet.
The atom tends to lose electrons. Valency = number of valence electrons.
Example: Sodium (2, 8, 1) loses 1 electron. Valency = 1.
The atom tends to gain electrons. Valency = 8 − valence electrons.
Example: Oxygen (2, 6) gains 2 electrons. Valency = 2.
The atom shares electrons with others.
Example: Carbon (2, 4) shares 4 electrons. Valency = 4.
H₂O (Water): Oxygen combines with 2 hydrogen atoms → Valency of O = 2
NH₃ (Ammonia): Nitrogen combines with 3 hydrogen atoms → Valency of N = 3
MgCl₂: Magnesium combines with 2 chlorine atoms → Valency of Mg = 2
Dalton proposed that all atoms of an element are identical. But scientists later discovered that atoms of the same element can have the same number of protons (same Z) yet different numbers of neutrons (different A). These "twin atoms" are called isotopes.
Naturally occurring hydrogen is a mixture of three isotopes:
1 proton, 0 neutrons
1 proton, 1 neutron
1 proton, 2 neutrons
Carbon has three isotopes: 126C (most abundant), 136C, and 146C. Each has 6 protons and 6 electrons, but they differ in neutrons (6, 7, and 8 respectively).
Uranium isotope used as fuel in nuclear reactors to generate electricity.
Radioactive cobalt used in radiation therapy for cancer.
Iodine isotope used to treat goitre and thyroid cancer.
Used in archaeology and geology to determine the age of ancient fossils and artefacts.
When an element has multiple isotopes, we calculate a weighted average atomic mass based on the natural abundance of each isotope.
Chlorine has two isotopes: 35Cl (75%) and 37Cl (25%)
Simple average: (35 + 37) ÷ 2 = 36 u (not accurate!)
Weighted average: (35 × 75/100) + (37 × 25/100) = 26.25 + 9.25 = 35.5 u
This doesn't mean any single atom has mass 35.5 u. It means if you take 1 million Cl atoms, their weighted average mass would be 35.5 u.
Isotopes have the same atomic number but different mass numbers. But what if atoms have the same mass number but different atomic numbers? These are called isobars.
| Element | Atomic Number (Z) | Mass Number (A) | Neutrons |
|---|---|---|---|
| Argon (Ar) | 18 | 40 | 22 |
| Potassium (K) | 19 | 40 | 21 |
| Calcium (Ca) | 20 | 40 | 20 |
All three have mass number 40, but they are different elements with different atomic numbers and different chemical properties.
Isotopes: Same Z, different A (same element, different mass)
Isobars: Same A, different Z (different elements, same mass)
Iso-topes = same type (element)
Iso-bars = same baric (mass)
Test your understanding with these interactive experiments and activities! Each one brings a key concept from this chapter to life.
Recreate Rutherford's famous experiment! Click "Fire!" to launch alpha particles at the gold foil. Watch how most pass through, some deflect, and very few bounce back.
Select an element to see its electrons fill up the K, L, and M shells following the 2n² rule.
Enter subatomic particle counts to calculate atomic number and mass number, or enter an element symbol to look up all values.
Sort these atom pairs into Isotopes (same atomic number, different mass number) or Isobars (same mass number, different atomic number). Click a pair to place it, then click "Check" to validate!
Test your knowledge with 8 quick-fire questions! Select the correct answer for each.
Atom is a sphere of positive charge with electrons embedded throughout (plum pudding model). First model to include subatomic particles.
Atom is mostly empty space with a dense, positively charged nucleus at centre. Electrons orbit like planets. Could not explain atomic stability.
Electrons move in fixed energy levels (K, L, M, N shells). No energy loss in stationary states. Explains atomic stability.
Electron (−1, in shells), Proton (+1, in nucleus), Neutron (0, in nucleus). Atom mass comes mainly from protons + neutrons.
Number of protons in the nucleus. Uniquely identifies an element. Z = number of protons = number of electrons.
Total protons + neutrons (nucleons). A = Z + number of neutrons.
Max electrons per shell: 2n². Outermost shell max: 8. Fill in order: K, L, M, N.
Combining capacity. Atoms lose/gain/share electrons to complete octet. Valence electrons decide reactivity.
Same atomic number, different mass numbers. Same chemical properties, different physical properties.
Same mass number, different atomic numbers. Different elements entirely.
1. Dalton's Model (1808): Proposed that all matter is made of indivisible atoms. Atoms of the same element are identical; atoms of different elements differ. This was the first scientific atomic theory but could not explain subatomic particles.
2. Thomson's Model (1897): After discovering electrons, Thomson proposed the plum pudding model -- a sphere of positive charge with electrons embedded throughout. Limitation: Could not explain the results of the gold foil experiment (large-angle deflections).
3. Rutherford's Model (1911): Through the gold foil experiment, proposed the nuclear/planetary model -- a tiny, dense, positively charged nucleus at the centre with electrons orbiting around it. Most of the atom is empty space. Limitation: Could not explain atomic stability (electrons should spiral into the nucleus).
4. Bohr's Model (1913): Proposed that electrons move in fixed energy levels (shells) without losing energy. Each shell has definite energy. Electrons can jump between shells by absorbing or releasing energy. This successfully explained atomic stability and many experimental observations.
Each model improved upon the previous one as new experimental evidence was discovered, showing how science moves forward through curiosity and experimentation.
Setup: Geiger and Marsden, under Rutherford, aimed a narrow beam of alpha particles (positively charged) at an extremely thin gold foil.
Observations:
• Most alpha particles passed straight through without deflection → Atom is mostly empty space.
• Some particles were deflected at large angles → They came close to a concentrated positive charge that repelled them.
• A very few bounced back (nearly 180°) → They hit something very dense and positively charged head-on.
Conclusions:
• The positive charge and most of the mass are packed into a tiny, dense nucleus at the centre.
• The nucleus is about 10⁵ times smaller than the atom.
• Electrons orbit the nucleus in the surrounding empty space (planetary model).
• This completely disproved Thomson's plum pudding model, which predicted uniform deflection.
Bohr-Bury Rules:
1. Maximum electrons in a shell = 2n² (where n is the shell number): K=2, L=8, M=18, N=32.
2. The outermost shell can hold a maximum of 8 electrons.
3. Electrons fill shells in order starting from the innermost: K → L → M → N.
Sodium (Z = 11): Total electrons = 11
• K-shell: 2 (full) → L-shell: 8 (full) → M-shell: 1
• Electronic configuration: 2, 8, 1
• Valence electrons = 1, Valency = 1
Chlorine (Z = 17): Total electrons = 17
• K-shell: 2 (full) → L-shell: 8 (full) → M-shell: 7
• Electronic configuration: 2, 8, 7
• Valence electrons = 7, Valency = 8 − 7 = 1
Isotopes: Atoms of the same element with the same atomic number (Z) but different mass numbers (A) due to different numbers of neutrons.
Example -- Carbon isotopes: 126C (6 neutrons), 136C (7 neutrons), 146C (8 neutrons). All have 6 protons and 6 electrons. They have same chemical properties but different physical properties.
Isobars: Atoms of different elements with the same mass number (A) but different atomic numbers (Z).
Example: 4018Ar, 4019K, and 4020Ca all have A = 40 but are different elements with different chemical properties.
Applications of Isotopes:
• 23592U is used as fuel in nuclear reactors for electricity generation.
• 6027Co is used in radiation therapy for cancer treatment.
• 13153I is used to treat goitre and thyroid cancer.
• 146C is used in carbon dating to determine the age of fossils and artefacts in archaeology.