Journey Inside the Atom
A simple guide to understanding the tiny building blocks that make up everything around us.
Everything you see, touch, or feel is made of matter. Matter is made of tiny particles called atoms. Both living things like our bodies and non-living things like a house are built from atoms. These particles are so small that we cannot see them with our eyes. For a long time, people wondered if atoms could be divided further. This chapter tells the story of how scientists discovered what lies inside the atom.
1. The Early Idea of Atoms
Ancient Thinkers
More than 2,000 years ago, thinkers in India and Greece asked the same question: What is everything made of?
Acharya Kanada in India suggested that if you keep dividing matter, you finally reach tiny particles that cannot be split any further. He called these particles paramanu. His ideas are written in the Sanskrit text Vaisesika Sutras. A paramanu is too small to be seen and cannot be divided. Combinations of paramanus form larger objects.
In Greece, philosophers Leucippus and Democritus had a similar idea. They called the smallest indivisible particle atomos, which means "uncuttable" or "indivisible" in Greek.
John Dalton's Atomic Theory
In 1808, John Dalton proposed the first scientific atomic theory based on experiments. He stated that:
- All matter is made of tiny particles called atoms.
- Atoms are indivisible and cannot be broken into smaller pieces.
- Atoms are the fundamental building blocks of matter.
Dalton's theory was the starting point for modern atomic science. However, later discoveries showed that atoms are not truly indivisible. They are made of even smaller particles.
2. The First Subatomic Particle: The Electron
J.J. Thomson's Discovery
In 1897, J.J. Thomson studied how electricity passes through gases at very low pressure. He used a glass tube with two electrodes and applied high voltage. He saw rays traveling from the negative electrode (cathode) to the positive electrode (anode). These are called cathode rays.
Thomson studied these rays using electric and magnetic fields. He found that cathode rays are streams of very small, negatively charged particles. These particles are much lighter than atoms. They were later named electrons. The charge of an electron is −1.602 × 10−19 C. For convenience, we write it as −1.
Thomson's Model of the Atom
After discovering electrons, Thomson faced a puzzle. Atoms are neutral, so there must be positive charge to balance the negative electrons. He proposed that an atom is a sphere of positive charge with electrons embedded in it, like seeds in a watermelon or plums in a pudding.
This model was the first real attempt to explain how positive and negative charges stay balanced in an atom. However, it was later proven wrong by new experiments.
3. Rutherford's Nuclear Model
The Gold Foil Experiment
In 1911, Ernest Rutherford, along with his students Geiger and Marsden, performed a famous experiment. They aimed a narrow beam of fast-moving alpha particles (positively charged particles from radioactive elements) at a very thin sheet of gold foil.
According to Thomson's model, the positive charge was spread evenly. So the scientists expected all alpha particles to pass straight through or bend only slightly. But the results were surprising:
- Most alpha particles passed straight through without any deflection.
- Some alpha particles were deflected at large angles.
- A very few alpha particles even bounced back.
What Rutherford Concluded
From this experiment, Rutherford proposed a new model of the atom:
- Most of the atom is empty space, because most alpha particles passed through.
- The positive charge and almost all the mass of the atom are concentrated in a very tiny central region called the nucleus.
- Electrons revolve around the nucleus, similar to how planets orbit the Sun. This is called the planetary model.
Rutherford calculated that the nucleus is about 100,000 times smaller than the atom. If an atom were the size of a cricket ground (about 100 m wide), the nucleus would be as small as a pepper grain at the centre.
Discovery of the Proton
Rutherford showed that the positive charge of the nucleus comes from particles called protons. A proton is much heavier than an electron and carries a charge of +1. In a neutral atom, the number of protons equals the number of electrons. For example, a helium atom has 2 protons and 2 electrons. A sodium atom has 11 protons and 11 electrons.
Why Rutherford's Model Was Not Complete
Rutherford's model could not explain why atoms are stable. According to physics, a charged particle moving in a circle is constantly accelerating. An accelerating electron should lose energy and spiral into the nucleus. If this happened, atoms would collapse, and matter would not exist. But atoms are stable. This meant Rutherford's model needed improvement.
4. Bohr's Model of the Atom
In 1913, Niels Bohr proposed a new model to fix the problems in Rutherford's model. His main ideas were:
- Electrons move only in certain fixed circular paths called stationary states, orbits, or shells.
- Each shell has a fixed energy, so they are also called energy levels.
- Shells are named K, L, M, N, or numbered n = 1, 2, 3, 4, starting from the nucleus.
- While moving in a fixed shell, an electron does not lose energy.
- An electron can jump from one shell to another by absorbing or releasing energy equal to the difference between the two energy levels.
5. What Gives an Atom Its Mass?
Rutherford showed that almost all the mass is in the nucleus. Electrons are so light that their mass can be ignored. But scientists found a puzzle: a helium atom has 2 protons, yet it is about 4 times heavier than a hydrogen atom (which has 1 proton). This meant there must be something else in the nucleus adding mass without adding charge.
Discovery of the Neutron
In 1932, James Chadwick discovered a new particle in the nucleus. It has almost the same mass as a proton but carries no electric charge. It was named the neutron (symbol n or n0). Neutrons are present in all atoms except ordinary hydrogen.
| Particle | Symbol | Relative Charge | Location |
|---|---|---|---|
| Electron | e− | −1 | Outside nucleus |
| Proton | p+ | +1 | Nucleus |
| Neutron | n0 | 0 | Nucleus |
6. How Scientists Name Elements
In 1803, John Dalton used picture symbols for elements. Later, in 1813, Jöns Jacob Berzelius suggested using letters from Latin names. Today, the International Union of Pure and Applied Chemistry (IUPAC) approves all names and symbols.
- The first letter is always a capital (uppercase).
- The second letter (if any) is a small (lowercase) letter.
- Examples: Hydrogen = H, Aluminium = Al, Cobalt = Co.
- Some symbols come from Latin or Greek names: Iron = Fe (from ferrum), Gold = Au (from aurum), Sodium = Na (from natrium).
| Element | Symbol | Element | Symbol | Element | Symbol |
|---|---|---|---|---|---|
| Aluminium | Al | Copper (Cuprum) | Cu | Nitrogen | N |
| Argon | Ar | Fluorine | F | Oxygen | O |
| Barium | Ba | Gold (Aurum) | Au | Potassium (Kalium) | K |
| Boron | B | Hydrogen | H | Silicon | Si |
| Bromine | Br | Iodine | I | Silver (Argentum) | Ag |
| Calcium | Ca | Iron (Ferrum) | Fe | Sodium (Natrium) | Na |
| Carbon | C | Lead (Plumbum) | Pb | Sulfur | S |
| Chlorine | Cl | Magnesium | Mg | Uranium | U |
| Cobalt | Co | Neon | Ne | Zinc | Zn |
7. Atomic Number and Mass Number
Atomic Number (Z)
The atomic number is the number of protons in the nucleus of an atom. It is written as Z. The atomic number tells us which element it is. All atoms of the same element have the same atomic number. In a neutral atom, the number of protons equals the number of electrons.
Mass Number (A)
The mass number is the total number of protons and neutrons in the nucleus. It is written as A. Protons and neutrons together are called nucleons.
Since electrons are very light, their mass is ignored in these calculations. The mass of a neutron is almost equal to the mass of a proton.
| Element | Protons | Neutrons | Mass Number (A) |
|---|---|---|---|
| Hydrogen | 1 | 0 | 1 |
| Helium | 2 | 2 | 4 |
| Lithium | 3 | 4 | 7 |
Standard Notation
Scientists write the symbol of an element with its mass number on top and atomic number at the bottom:
For example, carbon has atomic number 6 and mass number 12. It is written as:
8. Arrangement of Electrons in Shells
Bohr and Bury gave rules to decide how electrons are arranged around the nucleus:
- The maximum number of electrons in a shell is given by the formula 2n2, where n is the shell number.
- K-shell (n=1): 2 × 12 = 2 electrons
- L-shell (n=2): 2 × 22 = 8 electrons
- M-shell (n=3): 2 × 32 = 18 electrons
- The outermost shell can hold a maximum of 8 electrons (or 2 if it is the only shell).
- Electrons fill shells starting from the one closest to the nucleus (K), then outward to L, M, N, and so on.
Electronic Configuration of First 18 Elements
The way electrons are divided among shells is called the electronic configuration. Here is the data for the first eighteen elements:
| Element | Symbol | Atomic No. | Protons | Neutrons | Electrons | K (n=1) | L (n=2) | M (n=3) | N (n=4) |
|---|---|---|---|---|---|---|---|---|---|
| Hydrogen | H | 1 | 1 | 0 | 1 | 1 | – | – | – |
| Helium | He | 2 | 2 | 2 | 2 | 2 | – | – | – |
| Lithium | Li | 3 | 3 | 4 | 3 | 2 | 1 | – | – |
| Beryllium | Be | 4 | 4 | 5 | 4 | 2 | 2 | – | – |
| Boron | B | 5 | 5 | 6 | 5 | 2 | 3 | – | – |
| Carbon | C | 6 | 6 | 6 | 6 | 2 | 4 | – | – |
| Nitrogen | N | 7 | 7 | 7 | 7 | 2 | 5 | – | – |
| Oxygen | O | 8 | 8 | 8 | 8 | 2 | 6 | – | – |
| Fluorine | F | 9 | 9 | 10 | 9 | 2 | 7 | – | – |
| Neon | Ne | 10 | 10 | 10 | 10 | 2 | 8 | – | – |
| Sodium | Na | 11 | 11 | 12 | 11 | 2 | 8 | 1 | – |
| Magnesium | Mg | 12 | 12 | 12 | 12 | 2 | 8 | 2 | – |
| Aluminium | Al | 13 | 13 | 14 | 13 | 2 | 8 | 3 | – |
| Silicon | Si | 14 | 14 | 14 | 14 | 2 | 8 | 4 | – |
| Phosphorus | P | 15 | 15 | 16 | 15 | 2 | 8 | 5 | – |
| Sulfur | S | 16 | 16 | 16 | 16 | 2 | 8 | 6 | – |
| Chlorine | Cl | 17 | 17 | 18 | 17 | 2 | 8 | 7 | – |
| Argon | Ar | 18 | 18 | 22 | 18 | 2 | 8 | 8 | – |
9. Valency: The Combining Capacity
Atoms combine with each other to form molecules. The combining capacity of an atom is called its valency. It tells us how many hydrogen or chlorine atoms can combine with one atom of an element. Hydrogen and chlorine both have a valency of 1.
Examples:
- In water (H2O), oxygen combines with 2 hydrogen atoms. So, the valency of oxygen is 2.
- In ammonia (NH3), nitrogen combines with 3 hydrogen atoms. So, the valency of nitrogen is 3.
- In magnesium chloride (MgCl2), magnesium combines with 2 chlorine atoms. So, the valency of magnesium is 2.
How Valency Depends on Electrons
The outermost shell of an atom is called the valence shell, and the electrons in it are called valence electrons. If the outermost shell has 8 electrons, it is called an octet. Atoms with a complete octet (or 2 electrons in the case of helium) are very stable and do not react easily. Atoms with incomplete outer shells try to lose, gain, or share electrons to complete their octet.
The number of electrons gained, lost, or shared to complete the octet is the valency of the element.
- If valence electrons are less than 4, the atom tends to lose electrons. Valency = number of valence electrons.
- If valence electrons are more than 4, the atom tends to gain electrons. Valency = 8 − number of valence electrons.
- If valence electrons = 4, the atom tends to share electrons. Valency = 4.
Examples:
- Sodium (2, 8, 1): Has 1 valence electron. It loses 1 electron to get an octet. Valency = 1.
- Oxygen (2, 6): Has 6 valence electrons. It gains 2 electrons to complete the octet. Valency = 2.
- Carbon (2, 4): Has 4 valence electrons. It shares 4 electrons. Valency = 4.
10. Isotopes and Isobars
Isotopes
Dalton believed that all atoms of an element are identical. Later, scientists found that atoms of the same element can have different numbers of neutrons. These atoms have the same atomic number but different mass numbers. They are called isotopes.
Hydrogen has three natural isotopes:
- Protium (11H): 1 proton, 0 neutrons, 1 electron. Most common (~99.98%).
- Deuterium (21H): 1 proton, 1 neutron, 1 electron (~0.015%).
- Tritium (31H): 1 proton, 2 neutrons, 1 electron (very rare).
Carbon also has three isotopes:
- Carbon-12 (126C): 6 protons, 6 neutrons, 6 electrons. Most abundant.
- Carbon-13 (136C): 6 protons, 7 neutrons, 6 electrons.
- Carbon-14 (146C): 6 protons, 8 neutrons, 6 electrons. Used in carbon dating.
All isotopes of an element have the same number of electrons and the same electronic configuration. Therefore, they have the same chemical properties. However, they differ in physical properties like mass, boiling point, and melting point because they have different numbers of neutrons.
Uses of Isotopes
- Uranium-235: Used as fuel in nuclear power plants.
- Cobalt-60: Used in radiation treatment for cancer.
- Iodine-131: Used to treat thyroid disorders and goitre.
- Carbon-14: Used to find the age of ancient fossils and artefacts (carbon dating).
Average Atomic Mass
Elements in nature are usually a mixture of isotopes. The average atomic mass is calculated using the masses of all isotopes and how common (abundant) each one is. This is called a weighted average.
Chlorine has two main isotopes: 35Cl (75% abundance) and 37Cl (25% abundance).
Simple average = (35 + 37) ÷ 2 = 36 u
But this is wrong because the isotopes are not equally common. The correct weighted average is:
(35 × 75/100) + (37 × 25/100) = 26.25 + 9.25 = 35.5 u
So, the average atomic mass of chlorine is 35.5 u.
Isobars
Sometimes, atoms of different elements can have the same mass number but different atomic numbers. These are called isobars.
For example, calcium (atomic number 20), potassium (atomic number 19), and argon (atomic number 18) all have mass number 40. They have different numbers of protons, but the same total number of protons + neutrons.
- Isotopes: Same atomic number, different mass number (same element).
- Isobars: Same mass number, different atomic number (different elements).
11. The Journey of Atomic Models
Our understanding of the atom has changed many times over the years. Each new model was built on the discoveries before it.
Today, we know that electrons do not move in exact paths like planets. Instead, they exist as a probability cloud around the nucleus. We can predict the region where an electron is most likely to be found, but not its exact path. You will learn more about this quantum mechanical model in higher classes.
12. Solutions to Exercises
Section A: Conceptual Questions
(i) The experiment showed the existence of neutrons.
(ii) The results disproved the plum pudding model and led to the nucleus idea.
(iii) Large deflection of a few alpha particles showed that mass and positive charge are packed into a tiny centre.
(iv) The deflection showed that electrons move around the nucleus.
(i) Incorrect — Neutrons were discovered later by Chadwick in 1932. This experiment only revealed the nucleus.
(ii) Correct — The unexpected bouncing back of alpha particles proved that positive charge is concentrated, not spread out like pudding.
(iii) Correct — Only a dense, tiny nucleus could make heavy alpha particles bounce back.
(iv) Incorrect — The experiment showed the existence of a nucleus; it did not directly show how electrons move.
(i) Electrons lose energy in fixed orbits and fall into the nucleus.
(ii) Electrons can exist anywhere with no fixed energy.
(iii) Electrons revolve in fixed energy orbits without losing energy.
(iv) Electrons can be found between energy levels.
(i) Incorrect — Bohr stated that electrons in fixed orbits do not lose energy. This was the main improvement over Rutherford's model.
(ii) Incorrect — Electrons can only exist in allowed shells with fixed energy values.
(iii) Correct — This is Bohr's main postulate. Stationary states keep the atom stable.
(iv) Incorrect — Electrons cannot exist between shells; they jump directly from one level to another.
X: 18 protons, 19 neutrons
Y: 17 protons, 18 neutrons
Z: 17 protons, 20 neutrons
Explain the relation between (i) Y and Z, (ii) Z and X.
(i) Y and Z have the same number of protons (17) but different neutrons (18 vs 20). So they are isotopes of the same element (chlorine).
(ii) Z and X have different protons (17 vs 18) but the same mass number (17+20=37 and 18+19=37). So they are isobars.
(i) Bohr's fixed orbits
(ii) Thomson's plum pudding
(iii) Rutherford's dense nucleus
(iv) Dalton's indivisible atom
1. Dalton (1808) — Indivisible atom
2. Thomson (1897) — Plum pudding model
3. Rutherford (1911) — Nuclear model
4. Bohr (1913) — Fixed energy orbits
Reason (R): The number of electrons equals the number of protons in an atom.
Choose the correct option.
Atomic number = 12, so protons = 12 and electrons = 12.
Neutrons = Mass number − Atomic number = 24 − 12 = 12 neutrons.
Electronic configuration: K=2, L=8, M=2 (or 2, 8, 2).
(a) Lithium (Li): Total e− = 3, Valence e− = 1, Valency = 1, Protons = 3, Z = 3.
(b) Boron (B): Total e− = 5, Valence e− = 3, Valency = 3, Protons = 5, Z = 5.
(c) Aluminium (Al): Total e− = 13, Valence e− = 3, Valency = 3, Protons = 13, Z = 13.
(d) Fluorine (F): Total e− = 9, Valence e− = 7, Valency = 1, Protons = 9, Z = 9.
Neutrons = 70 − 31 = 39 neutrons.
In a neutral atom, electrons = 79.
Neutrons = 197 − 79 = 118 neutrons.
This element is gold (Au).
| Atomic No. | Mass No. | Neutrons | Protons | Electrons | Name |
|---|---|---|---|---|---|
| 5 | ? | 6 | ? | ? | ? |
| ? | 14 | ? | ? | 7 | Nitrogen |
| ? | 24 | ? | 12 | ? | ? |
| 15 | ? | 16 | ? | ? | ? |
| ? | 1 | 0 | ? | ? | ? |
Row 1: Z=5, so protons=5, electrons=5. Mass no. = 5+6 = 11. Element = Boron (B).
Row 2: Electrons=7, so protons=7, Z=7. Neutrons = 14−7 = 7.
Row 3: Protons=12, so Z=12, electrons=12. Neutrons = 24−12 = 12. Element = Magnesium (Mg).
Row 4: Z=15, so protons=15, electrons=15. Mass no. = 15+16 = 31. Element = Phosphorus (P).
Row 5: Mass no.=1, neutrons=0, so protons=1, Z=1, electrons=1. Element = Hydrogen (H).
(i) Electrons and protons? (ii) Atomic number? (iii) Identify X.
(iv) Electronic configuration? (v) Valence electrons? (vi) Mass number if 2 neutrons are added? (vii) Relation with new atom?
(i) Protons = 35 − 18 = 17. Electrons = 17 (neutral atom).
(ii) Atomic number = 17.
(iii) The element is Chlorine (Cl).
(iv) Electronic configuration: 2, 8, 7.
(v) Valence electrons = 7.
(vi) New mass number = 35 + 2 = 37.
(vii) The new atom is an isotope of X (same atomic number, different mass number).
(i) Atomic number (ii) Atomic mass (iii) Mass number (iv) Overall charge?
(i) Atomic number remains 12 because it depends only on the number of protons.
(ii) Atomic mass increases because the new particles are much heavier than electrons, though still small compared to the nucleus.
(iii) Mass number remains 24 because mass number counts only protons and neutrons in the nucleus.
(iv) Overall charge remains neutral because the new particles have the same negative charge as electrons, balancing the 12 positive protons.
Section B: Numerical Problems
Mass number = 56, so neutrons = 56 − 26 = 30 neutrons.
This element is iron (Fe).
Neutrons = 23 − 11 = 12 neutrons. This is sodium-23 (Na).
(i) 126C (ii) 199F (iii) 2814Si
(i) Carbon: Z=6, configuration 2,4. Valence electrons = 4.
(ii) Fluorine: Z=9, configuration 2,7. Valence electrons = 7.
(iii) Silicon: Z=14, configuration 2,8,4. Valence electrons = 4.
Z=12 (Magnesium): 2, 8, 2
Z=16 (Sulfur): 2, 8, 6
Z=18 (Argon): 2, 8, 8
Neutrons = 23 − 11 = 12 neutrons.
Sample riddle: "I have 6 protons and 6 neutrons. I am the basis of life. Who am I?" Answer: Carbon-12.
Mass numbers are 23 (11+12) and 24 (11+13).
They are isotopes of sodium.
Average mass = (79 × 49.7/100) + (81 × 50.3/100)
= 39.263 + 40.743 = 80.006 u (approximately 80.0 u).
Section C: Think and Explain
(i) If positive charge is less than negative charge, the model becomes a negative ion (anion). It is not neutral.
(ii) If the clay itself carries negative charge, the total negative charge becomes even larger. The model is not neutral and does not represent a neutral atom.