Class 9 › Science › Chapter 8
Journey Inside the Atom
Evolution of atomic models (Dalton, Thomson, Rutherford, Bohr), subatomic particles, atomic and mass number, electronic configuration, valency, isotopes and isobars.
1. Roots of Atomic Theory
- Acharya Kanada (India): dividing matter (dravya) repeatedly reaches smallest indivisible particles, parmanus (Vaisesika Sutras). They are infinitely small, cannot be perceived by senses, and combine as dyads (2) and triads (3) to form all matter. The proportions of combination were not specified.
- Leucippus and Democritus (Greece): indivisible particles called atomos (Greek for 'indivisible').
- These ideas were imaginary / philosophical, not based on experiments.
- John Dalton (1808): first scientific atomic theory, based on experiments: all matter is made of indivisible atoms, the fundamental building blocks. It was the starting point for later models.
- Questions that followed: what are atoms made of, what do they look like, and what makes atoms of one element differ from another?
- Radioactivity (certain elements emit radiation) showed in the late 19th century that atoms must contain smaller particles.
2. Discovery of the Electron and Thomson's Model
J. J. Thomson's cathode ray experiment (1897)
A glass tube at very low gas pressure, two electrodes, high voltage: rays travel from cathode (−) to anode (+): cathode rays.
In electric and magnetic fields they behave as negatively charged particles with a mass much smaller than an atom: electrons.
The rays were independent of the cathode material and the gas, so electrons are a fundamental component of all atoms.
- Charge of an electron: $-1.602\times10^{-19}$ C, taken as −1 by convention.
- Thomson's model (plum pudding / watermelon): a sphere of positive charge with electrons embedded throughout it; charges balance so the atom is neutral.
- Thomson received the 1906 Nobel Prize in Physics; he headed the Cavendish Laboratory, Cambridge, and guided Rutherford.
3. Rutherford's Model and the Proton
Gold foil (α-ray scattering) experiment, 1911 (Geiger and Marsden, under Rutherford)
A narrow beam of α-particles (positively charged; nuclei of helium: 2 protons + 2 neutrons) hits a very thin gold foil.
Expected (Thomson): pass straight through or tiny deflection.
Observed: most passed undeflected; some deflected sharply (scattering); a few bounced back.
| Observation | Conclusion |
|---|---|
| Most α-particles pass straight through | Most of the atom is empty space |
| Some deflected at large angles; a few bounce back | Positive charge and most of the mass are concentrated in a tiny, dense nucleus |
- Rutherford's (planetary) model: a dense positively charged nucleus at the centre containing all positive charge and most of the mass; electrons revolve around it like planets around the Sun.
- Size: atom diameter $\approx 10^{-10}$ m; nucleus $\approx 10^{-15}$ m, about $10^5$ times smaller. (If an atom were a 100 m cricket ground, the nucleus would be a black pepper grain.)
- Limitation: an electron in circular motion is accelerating, so it should lose energy and spiral into the nucleus; the model could not explain stability of atoms.
- Proton: discovered and named by Rutherford; positive charge $+1$ (equal and opposite to the electron), much heavier than an electron. Number of protons = number of electrons in a neutral atom (eg: helium 2 and 2; sodium 11 and 11).
- Ernest Rutherford (New Zealand): 'Father of Nuclear Physics'; Nobel Prize in Chemistry 1908.
- Estimate: a 0.1 mm ($10^{-4}$ m) sheet of paper is about $10^{-4}/10^{-10} = 10^6$ atoms thick.
4. Bohr's Model of the Atom (1913)
Postulates of Bohr's model
1. Electrons move in fixed circular paths called shells / orbits / stationary states / energy levels.
2. Shells are named K, L, M, N, … ($n = 1, 2, 3, 4,\ldots$); K is nearest the nucleus.
3. In a shell an electron has definite energy and does not lose energy while moving.
4. Energy increases with distance from the nucleus (K lowest).
5. Electrons can jump between shells by absorbing or releasing a fixed amount of energy equal to the energy difference. They cannot exist in between shells.
6. Each shell holds only a limited number of electrons.
2. Shells are named K, L, M, N, … ($n = 1, 2, 3, 4,\ldots$); K is nearest the nucleus.
3. In a shell an electron has definite energy and does not lose energy while moving.
4. Energy increases with distance from the nucleus (K lowest).
5. Electrons can jump between shells by absorbing or releasing a fixed amount of energy equal to the energy difference. They cannot exist in between shells.
6. Each shell holds only a limited number of electrons.
- Stability: Bohr introduced stationary states (a postulate) in which the energy stays constant, so electrons do not spiral into the nucleus. This is why Bohr's model explained what Rutherford's could not.
- K, L, M, N naming: from Charles Barkla's X-ray experiments (first line called K, leaving room for earlier series).
- Niels Bohr: professor at Copenhagen; Nobel Prize 1922.
- Later: electrons do not follow fixed paths; the modern quantum mechanical model describes electron clouds, regions where electrons are most likely found.
| Model | Scientist (year) | Key idea |
|---|---|---|
| Indivisible atom | Dalton (1808) | Atoms are indivisible particles |
| Plum pudding | Thomson (1897 discovery) | Electrons embedded in a positive sphere |
| Nuclear / planetary | Rutherford (1911) | Dense nucleus; electrons orbit; mostly empty space |
| Energy levels | Bohr (1913) | Electrons in fixed shells; stationary states |
| Quantum mechanical | Modern | Electron clouds; regions of probability |
5. The Neutron and the Mass of an Atom
- Puzzle: hydrogen has 1 proton, helium 2 protons, but a helium atom is about 4 times (not 2 times) as heavy as hydrogen.
- James Chadwick (1932) discovered the neutron: mass nearly equal to a proton, no charge. Found in the nucleus of all atoms except hydrogen (ordinary hydrogen). Nobel Prize in Physics, 1935.
- The mass of an atom comes mainly from protons and neutrons in the nucleus; the electron's mass is negligible.
| Particle | Symbol | Relative charge | Location | Discovered by |
|---|---|---|---|---|
| Electron | e− | −1 | Shells around the nucleus | J. J. Thomson |
| Proton | p+ | +1 | Nucleus | Rutherford |
| Neutron | n0 | 0 | Nucleus | Chadwick |
- Neutron-to-proton ratio: light atoms have about equal numbers (carbon 6 and 6; oxygen 8 and 8); heavier ones have many more neutrons (iron 26 p, 30 n; uranium 92 p, 146 n).
- Why protons don't fly apart: neutrons keep protons farther apart and strengthen the nuclear force that binds the nucleus together.
- BARC (Mumbai) uses neutron scattering with reactors such as Dhruva to study materials, batteries and drug molecules. About 118 elements are known today (69 were known by 1869).
6. Symbols of Elements
- Dalton (1803): first pictorial symbols for elements. Berzelius (1813): symbols from Latin names (alphabetic). Today the IUPAC approves names and symbols.
- Rules: first letter capital, second letter (if any) small (Al, not AL; Co, not CO). Many are the first one or two letters (H, Al); some use another letter (Cl, Zn); some come from Latin, Greek or German names.
| 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 |
- Other origins: mercury Hg (Greek hydrargyros), tungsten W (German wolfram), iron Fe (Latin ferrum).
- Symbols are international, so scientists everywhere communicate clearly.
7. Atomic Number and Mass Number
Definitions
Atomic number $Z$ = number of protons in the nucleus (identifies the element; for a neutral atom, equals the number of electrons).
Mass number $A$ = protons + neutrons (the nucleons).
$$A = Z + n \qquad\Rightarrow\qquad n = A - Z$$
Notation: ${}^{A}_{Z}\text{X}$ (eg: ${}^{12}_{6}\text{C}$: $Z=6$, $A=12$, 6 neutrons).
Mass number $A$ = protons + neutrons (the nucleons).
| Element | Protons | Neutrons | Mass number |
|---|---|---|---|
| Hydrogen | 1 | 0 | 1 |
| Helium | 2 | 2 | 4 |
| Lithium | 3 | 4 | 7 |
- Atomic number decides identity and chemical behaviour; electrons are ignored in mass-number calculations.
- All atoms are electrically neutral because protons = electrons.
8. Distribution of Electrons in Shells
Bohr-Bury rules
1. Maximum electrons in a shell $= 2n^2$: K ($n=1$) $=2$; L ($n=2$) $=8$; M ($n=3$) $=18$.
2. The outermost shell can hold at most 8 electrons (the first shell: 2).
3. Shells fill in order: K, then L, then M, … (next shell only after the previous is complete).
2. The outermost shell can hold at most 8 electrons (the first shell: 2).
3. Shells fill in order: K, then L, then M, … (next shell only after the previous is complete).
Bohr model of sodium: nucleus (11 protons, 12 neutrons) and electron shells 2, 8, 1.
Electronic configuration: the distribution of electrons among the shells, written as K, L, M (eg: sodium 2, 8, 1).
| Element | Symbol | Z | Neutrons | Configuration (K, L, M) | Valency |
|---|---|---|---|---|---|
| Hydrogen | H | 1 | 0 | 1 | 1 |
| Helium | He | 2 | 2 | 2 | 0 |
| Lithium | Li | 3 | 4 | 2, 1 | 1 |
| Beryllium | Be | 4 | 5 | 2, 2 | 2 |
| Boron | B | 5 | 6 | 2, 3 | 3 |
| Carbon | C | 6 | 6 | 2, 4 | 4 |
| Nitrogen | N | 7 | 7 | 2, 5 | 3 |
| Oxygen | O | 8 | 8 | 2, 6 | 2 |
| Fluorine | F | 9 | 10 | 2, 7 | 1 |
| Neon | Ne | 10 | 10 | 2, 8 | 0 |
| Sodium | Na | 11 | 12 | 2, 8, 1 | 1 |
| Magnesium | Mg | 12 | 12 | 2, 8, 2 | 2 |
| Aluminium | Al | 13 | 14 | 2, 8, 3 | 3 |
| Silicon | Si | 14 | 14 | 2, 8, 4 | 4 |
| Phosphorus | P | 15 | 16 | 2, 8, 5 | 3 |
| Sulfur | S | 16 | 16 | 2, 8, 6 | 2 |
| Chlorine | Cl | 17 | 18 | 2, 8, 7 | 1 |
| Argon | Ar | 18 | 22 | 2, 8, 8 | 0 |
- The valency column is derived from each configuration (the textbook leaves it for you to add).
9. Valency: Combining Capacity
- Combining capacity is expressed in terms of hydrogen or chlorine atoms (both have capacity 1). Eg: H2O: oxygen 2; NH3: nitrogen 3; MgCl2: magnesium 2.
- Valence shell: the outermost shell containing electrons; its electrons are valence electrons.
- Octet: 8 electrons in the valence shell (2 for helium). Atoms with a complete octet are stable and largely unreactive; others lose, gain or share electrons to complete it.
- Valency: the number of electrons gained, lost or shared to complete the octet.
| Valence electrons | Tendency | Valency | Example |
|---|---|---|---|
| Fewer than 4 | Lose electrons | = number of valence electrons | Sodium 2, 8, 1: loses 1; valency 1 |
| More than 4 | Gain electrons | = 8 − valence electrons | Oxygen 2, 6: gains 2; valency 2 |
| Exactly 4 | Share electrons | 4 | Carbon 2, 4: valency 4 |
| Octet complete | Unreactive | 0 | Neon 2, 8; argon 2, 8, 8 |
10. Isotopes, Isobars and Average Atomic Mass
Isotopes
Same Z, different A
Atoms of the same element with the same atomic number but different mass numbers (different neutrons). Same chemical properties (same electronic configuration); different physical properties (boiling and melting points).
Isobars
Same A, different Z
Atoms of different elements with the same mass number but different atomic numbers. Eg: argon (Z = 18), potassium (19), calcium (20), all with $A = 40$.
- Hydrogen isotopes: ${}^{1}_{1}$H protium (about 99.98%), ${}^{2}_{1}$H deuterium (about 0.015%, 1 neutron), ${}^{3}_{1}$H tritium (traces, 2 neutrons); all have 1 proton and 1 electron.
- Carbon isotopes: ${}^{12}_{6}$C (most abundant), ${}^{13}_{6}$C, ${}^{14}_{6}$C: all 6 protons and 6 electrons.
- Dalton thought all atoms of an element have the same mass; isotopes showed this is not exactly so.
| Isotope | Use |
|---|---|
| ${}^{235}_{92}$U | Fuel in nuclear reactors (electricity) |
| ${}^{60}_{27}$Co | Radiation treatment of cancer |
| ${}^{131}_{53}$I | Treat goitre and thyroid cancer |
| ${}^{14}_{6}$C | Age of ancient fossils and artefacts (archaeology, geology) |
Average (weighted) atomic mass
$$\text{average atomic mass} = \sum (\text{isotope mass}\times\text{fractional abundance})$$
Chlorine: ${}^{35}$Cl : ${}^{37}$Cl $= 3:1$ (75% : 25%)
$$35\times\tfrac{75}{100} + 37\times\tfrac{25}{100} = 35.5\ \text{u}$$
A simple average $(35+37)/2 = 36$ ignores abundance and is wrong.
- Unit: unified atomic mass unit (u) (earlier 'amu') is used because a kilogram is far too large for atoms.
- 35.5 u does not mean any single atom has a fractional mass; in 1 million chlorine atoms, about 7.5 lakh are ${}^{35}$Cl and 2.5 lakh are ${}^{37}$Cl.
- Homi Jehangir Bhabha: father of the Indian nuclear programme; established TIFR and BARC.
- Atomic-scale imaging: Scanning Tunnelling Microscopes (surfaces) and Transmission Electron Microscopes (thin samples) show individual atoms.
11. Quick Revision
| To find | Use |
|---|---|
| Protons | = atomic number $Z$ |
| Electrons (neutral atom) | = $Z$ |
| Neutrons | $A - Z$ |
| Mass number | $Z$ + neutrons |
| Electrons in shell $n$ (max) | $2n^2$ |
| Valency | Valence electrons (if < 4) or $8 - $ valence electrons (if > 4); carbon 4; octet complete = 0 |
| Weighted average mass | $\sum$ (mass × abundance fraction) |
Remember
Order of discoveries
Electron (Thomson), nucleus and proton (Rutherford), neutron (Chadwick). Models: Dalton → Thomson → Rutherford → Bohr → quantum.
Remember
Isotope vs isobar
Isotopes: same Z (same element). Isobars: same A (different elements).
Remember
Why Bohr succeeded
Stationary states: electrons in fixed shells do not lose energy, so atoms are stable.
Common mistakes:
• Saying atoms are indivisible (they contain electrons, protons and neutrons).
• Saying electrons contribute to mass number, or neutrons are charged.
• Writing symbols with wrong case (CO for cobalt, AL for aluminium).
• Using a simple average instead of a weighted average for atomic mass.
• Putting more than 8 electrons in the outer shell, or filling M before L is complete.
• Saying isotopes differ in chemical properties (they differ in physical properties).
• Saying the nucleus is half the size of the atom (it is about $10^5$ times smaller).
• Giving a valency of 5 or 6 to P or S (count electrons needed to complete the octet).
• Saying atoms are indivisible (they contain electrons, protons and neutrons).
• Saying electrons contribute to mass number, or neutrons are charged.
• Writing symbols with wrong case (CO for cobalt, AL for aluminium).
• Using a simple average instead of a weighted average for atomic mass.
• Putting more than 8 electrons in the outer shell, or filling M before L is complete.
• Saying isotopes differ in chemical properties (they differ in physical properties).
• Saying the nucleus is half the size of the atom (it is about $10^5$ times smaller).
• Giving a valency of 5 or 6 to P or S (count electrons needed to complete the octet).