1. What is an Atom?
An atom is the basic unit of an element that retains the chemical identity of that element. Everything around us is made of matter, and matter is built from atoms and molecules.
Atoms are extremely small. NCERT gives the approximate diameter of an atom as about 10−10 m. The nucleus is much smaller, with a diameter of about 10−15 m. Thus, most of the volume of an atom is empty space. citeturn0search12
2. Subatomic Particles
Atoms are not indivisible. They contain smaller particles called subatomic particles: electrons, protons and neutrons.
| Particle | Charge | Relative mass | Location |
|---|---|---|---|
| Electron (e−) | −1 | Very small compared with proton/neutron | Outside nucleus, in electron shells/regions |
| Proton (p+) | +1 | About 1 u | Nucleus |
| Neutron (n0) | 0 | About 1 u | Nucleus |
Who discovered what?
- Electron: J. J. Thomson’s cathode-ray experiments established the existence of negatively charged particles.
- Proton: Positive rays and later work associated with the hydrogen nucleus established the proton as the positively charged nuclear particle.
- Neutron: James Chadwick discovered the neutron in 1932.
3. Thomson’s Model of the Atom
J. J. Thomson proposed a model in which the atom was considered to be a positively charged sphere with electrons embedded in it.
This model is popularly described as the “plum pudding” model. The positive charge was assumed to be spread throughout the atom, with negatively charged electrons embedded inside.
Why was it important?
It was an early attempt to include the newly discovered electron in a model of the atom and explain why atoms are electrically neutral.
Why was it rejected?
Rutherford’s gold foil experiment showed that most of the positive charge and mass are concentrated in a tiny nucleus. This was inconsistent with a diffuse positive charge spread throughout the atom.
4. Rutherford’s Gold Foil Experiment
Ernest Rutherford and his collaborators studied how α-particles were scattered when directed at a very thin sheet of gold foil.
Expected result from Thomson’s model
If positive charge were spread throughout the atom, α-particles should mostly pass through with only small deflections.
Actual observations
Passed straight through the foil.
Were deflected through small angles.
Were deflected through large angles or bounced back.
What did this mean?
- Most of the atom is empty space.
- Positive charge is concentrated in a very small region.
- The central region is dense and contains most of the atom’s mass.
- The central region is called the nucleus.
5. Rutherford’s Nuclear / Planetary Model
From the gold foil experiment, Rutherford proposed the nuclear model. The nucleus is extremely small, dense and positively charged, while electrons occupy the region around it. NCERT describes this as the planetary model because electrons were pictured as moving around the nucleus somewhat like planets around the Sun. citeturn0search12
Contains positive charge and most of the atom’s mass; extremely small compared with the atom.
Most of the atom’s volume is empty space, with electrons occupying the surrounding region.
Scale comparison
NCERT gives the atom’s diameter as approximately 10−10 m and the nucleus as approximately 10−15 m. The nucleus is therefore about 100,000 times smaller in diameter than the atom. citeturn0search12
6. Bohr’s Model of the Atom
Rutherford’s model explained the nucleus but did not adequately explain the stability of atoms and the arrangement of electrons. Niels Bohr proposed that electrons occupy specific allowed energy levels or shells around the nucleus.
Main ideas at this level
- Electrons occupy definite energy levels called shells.
- Electrons in a shell have a definite energy associated with that shell.
- Electrons can move between allowed energy levels by absorbing or releasing energy.
7. Electron Shells
Electrons are arranged in shells around the nucleus. They are commonly labelled K, L, M, N... from the shell closest to the nucleus outward.
| Shell | Principal shell number | Maximum capacity (2n²) |
|---|---|---|
| K | n = 1 | 2 |
| L | n = 2 | 8 |
| M | n = 3 | 18 |
| N | n = 4 | 32 |
For the first few elements studied at this level, electron distributions can be written in forms such as 2, 1 for lithium, 2, 8, 1 for sodium and 2, 8, 7 for chlorine.
8. Electronic Configuration
The distribution of electrons among shells is called the electronic configuration of an atom.
| Element | Atomic number | Electron configuration |
|---|---|---|
| Hydrogen | 1 | 1 |
| Helium | 2 | 2 |
| Carbon | 6 | 2, 4 |
| Oxygen | 8 | 2, 6 |
| Neon | 10 | 2, 8 |
| Sodium | 11 | 2, 8, 1 |
| Magnesium | 12 | 2, 8, 2 |
| Chlorine | 17 | 2, 8, 7 |
| Argon | 18 | 2, 8, 8 |
The electrons in the outermost shell are called valence electrons. They play a major role in chemical bonding and the formation of ions.
9. Valency
Valency describes the combining capacity of an atom. At this level, it can be understood by considering how many electrons an atom tends to lose, gain or share to achieve a more stable outer-shell arrangement.
| Element | Configuration | Valency (common school-level value) |
|---|---|---|
| Na | 2,8,1 | 1 |
| Mg | 2,8,2 | 2 |
| Al | 2,8,3 | 3 |
| C | 2,4 | 4 |
| O | 2,6 | 2 |
| Cl | 2,8,7 | 1 |
| Ne | 2,8 | 0 |
Stable outer shell
Noble gases such as helium, neon and argon have stable outer electron arrangements. Other atoms may lose, gain or share electrons in chemical interactions to reach more stable configurations.
10. Formation of Ions
An ion is a charged atom or group of atoms. Ions form when atoms gain or lose electrons.
A positively charged ion formed when an atom loses one or more electrons.
Example: Na → Na+ + e−
A negatively charged ion formed when an atom gains one or more electrons.
Example: Cl + e− → Cl−
Example: Sodium chloride
The resulting oppositely charged ions attract each other, forming an ionic compound such as sodium chloride.
11. Atomic Number and Mass Number
Atomic number (Z)
The atomic number is the number of protons in the nucleus of an atom.
For a neutral atom:
Mass number (A)
The mass number is the total number of protons and neutrons in the nucleus.
Example
For an atom with Z = 17 and A = 35:
12. Isotopes
Isotopes are atoms of the same element having the same atomic number but different mass numbers. They therefore have the same number of protons but different numbers of neutrons.
| Isotope | Protons | Neutrons | Mass number |
|---|---|---|---|
| Hydrogen-1 (¹H) | 1 | 0 | 1 |
| Hydrogen-2 (²H) | 1 | 1 | 2 |
| Hydrogen-3 (³H) | 1 | 2 | 3 |
| Carbon-12 | 6 | 6 | 12 |
| Carbon-14 | 6 | 8 | 14 |
Uses of isotopes
- Cobalt-60: used as a source of radiation in some medical and industrial applications.
- Iodine-131: used in certain medical diagnostic and treatment applications involving the thyroid.
- Carbon-14: useful for radiocarbon dating of once-living materials.
13. Why Understanding the Atom Matters
Electron behaviour helps explain electrical properties of materials.
Radioisotopes have diagnostic and therapeutic applications.
Atomic and nuclear science supports materials analysis and radiation-based technologies.
Radiation and isotope techniques can support research and some agricultural applications.
Electron arrangement helps explain bonding, valency and chemical reactions.
Atomic structure is foundational to modern physics, chemistry and materials science.
14. Quick Revision
15. Practice Questions
- What is an atom?
- Name the three main subatomic particles.
- State the charge and location of an electron, proton and neutron.
- Why is a neutral atom electrically neutral?
- What did J. J. Thomson’s cathode-ray experiments show?
- Describe Thomson’s model of the atom.
- Why was Thomson’s model rejected?
- Describe the main observations of Rutherford’s gold foil experiment.
- What did Rutherford conclude from the fact that most α-particles passed straight through?
- Why was the rare back-scattering of α-particles significant?
- State the main features of Rutherford’s nuclear model.
- Compare the approximate sizes of an atom and its nucleus.
- Why did Bohr propose specific electron shells?
- Write the maximum electron capacity of the K, L, M and N shells using 2n².
- What is electronic configuration?
- Write the electronic configurations of carbon, oxygen, sodium and chlorine.
- What are valence electrons?
- Define valency.
- Explain how sodium forms Na⁺.
- Explain how chlorine forms Cl⁻.
- What is a cation? Give an example.
- What is an anion? Give an example.
- Define atomic number and mass number.
- An atom has Z = 12 and A = 24. Find protons, electrons and neutrons.
- An atom has 17 protons and 18 neutrons. Find its atomic number and mass number.
- What are isotopes?
- Why do isotopes of the same element have the same atomic number?
- Give two applications of isotopes.
- Differentiate Rutherford’s and Bohr’s models.
- Explain how the study of atomic structure helps us understand chemistry and modern technology.
16. Interactive 30-Question Quiz
Choose one option for every question, then click Check Score.