Chapter 4 — Structure of the Atom
Welcome to HSLC Guru! This study guide presents complete English-medium notes, summary, and question answers for Class 9 Science Chapter 4 — Structure of the Atom from the ASSEB (Assam State School Education Board) curriculum. The chapter takes you on a fascinating journey from cathode ray experiments to the modern picture of the atom, explaining how electrons, protons, and neutrons were discovered, how they are arranged inside an atom, and how this arrangement controls the chemical behaviour of every element. Topics covered include Thomson’s plum-pudding model, Rutherford’s nuclear atom, Bohr’s shell model, the 2n² rule, valency, atomic number, mass number, isotopes, and isobars. All content is prepared strictly according to ASSEB guidelines for Class 9 students.
Summary
For a long time the atom was thought to be the smallest, indivisible particle of matter. The first hint that the atom has internal structure came from experiments with discharge tubes. When a high voltage is applied across a gas at very low pressure, a stream of negatively charged particles travels from the cathode to the anode. These streams are called cathode rays. Sir J. J. Thomson studied cathode rays in 1897 and showed that they consist of tiny negatively charged particles, which he called electrons. The charge-to-mass ratio of the electron was found to be the same regardless of the gas used, proving that electrons are present in every atom. Based on this discovery, Thomson proposed the plum-pudding model: the atom is a sphere of positive charge in which electrons are embedded like plums in a pudding (or seeds in a watermelon), so that the atom as a whole is electrically neutral.
In 1886 E. Goldstein observed positively charged rays in a modified discharge tube; these were named canal rays or anode rays, leading to the discovery of the proton. Years later in 1932, James Chadwick discovered a third subatomic particle, the neutron, which carries no charge and has a mass nearly equal to that of the proton. Meanwhile, in 1911, Ernest Rutherford performed his famous α-particle (alpha-particle) scattering experiment. He bombarded a thin gold foil with fast-moving α-particles and observed that (i) most particles passed straight through, (ii) a few were deflected through small angles, and (iii) very few (about 1 in 12,000) bounced back. From these observations Rutherford concluded that the atom is mostly empty space, that the entire positive charge and almost all the mass are concentrated in a tiny central region called the nucleus, and that electrons revolve around the nucleus. This is the nuclear model of the atom. Its main drawback was that, according to classical physics, a revolving electron should continuously radiate energy, spiral into the nucleus, and the atom should collapse — which does not happen.
To overcome this difficulty, Niels Bohr in 1913 proposed a new model. According to Bohr, electrons revolve around the nucleus only in certain fixed circular paths called orbits or energy shells, named K, L, M, N… from the nucleus outward. While moving in these orbits, electrons do not radiate energy. Energy is absorbed or emitted only when an electron jumps from one orbit to another. The distribution of electrons in these shells follows the Bohr–Bury scheme: the maximum number of electrons in the n-th shell is given by the formula 2n² (so K = 2, L = 8, M = 18, N = 32). Two further rules state that the outermost shell of any atom can never hold more than 8 electrons, and the next-to-outermost shell cannot hold more than 18. The number of electrons in the outermost shell decides the valency of the element — for elements with 1, 2 or 3 outer electrons, valency equals the number of outer electrons; for elements with 5, 6, 7 outer electrons, valency = 8 − (number of outer electrons). Atoms having 8 electrons in the outermost shell (the noble gases) are chemically inert and have valency zero.
Two important whole numbers describe an atom. The atomic number (Z) is the number of protons in the nucleus, which also equals the number of electrons in a neutral atom. The mass number (A) is the total number of protons and neutrons (together called nucleons). An element is symbolised as ᴬZX. Atoms of the same element having the same atomic number but different mass numbers are called isotopes. Common examples are the three isotopes of hydrogen — protium ¹H, deuterium ²H and tritium ³H; the carbon isotopes ¹²C and ¹⁴C; and the chlorine isotopes ³⁵Cl and ³⁷Cl (which give chlorine its average atomic mass of 35.5 u). Isotopes have identical chemical properties but different physical properties. They have many useful applications: ¹⁴C is used in radio-carbon dating of fossils, ⁶⁰Co (cobalt-60) and ¹³¹I (iodine-131) are used in the treatment of cancer and thyroid disorders, ²³⁵U is used as a fuel in nuclear reactors, and ³¹P is used as a tracer in agriculture. Atoms of different elements that have the same mass number but different atomic numbers are called isobars — for example, ⁴⁰18Ar, ⁴⁰19K and ⁴⁰20Ca all have mass number 40.
Textbook Questions and Answers
1-Mark Questions
Q1. Who discovered the electron?
Answer: The electron was discovered by Sir J. J. Thomson in 1897 during his study of cathode rays in a discharge tube.
Q2. What is the charge and approximate mass of a proton?
Answer: A proton carries one unit of positive charge (+1) and has a mass of approximately 1 atomic mass unit (1 u), nearly 1836 times the mass of an electron.
Q3. Name the scientist who discovered the neutron.
Answer: The neutron was discovered by James Chadwick in 1932.
Q4. What is meant by atomic number?
Answer: The atomic number (Z) of an element is the number of protons present in the nucleus of one atom of that element. In a neutral atom, it is also equal to the number of electrons.
Q5. Define mass number.
Answer: The mass number (A) is the total number of protons and neutrons (collectively called nucleons) present in the nucleus of an atom.
Q6. Write the maximum number of electrons that can be accommodated in the L shell.
Answer: Using the formula 2n² with n = 2, the L shell can hold a maximum of 2 × (2)² = 8 electrons.
Q7. What are the three isotopes of hydrogen?
Answer: The three isotopes of hydrogen are protium (¹H), deuterium (²H) and tritium (³H).
Q8. What is the valency of an element having 7 electrons in its outermost shell?
Answer: Valency = 8 − 7 = 1. So the element shows a valency of 1 (for example, chlorine).
Q9. Why is the atom electrically neutral?
Answer: An atom is electrically neutral because the number of positively charged protons in its nucleus is exactly equal to the number of negatively charged electrons revolving around the nucleus, so the total positive and negative charges cancel out.
Q10. Define isobars with one example.
Answer: Atoms of different elements which have the same mass number but different atomic numbers are called isobars. Example: ⁴⁰18Ar and ⁴⁰20Ca.
2–3 Mark Questions
Q1. State the main postulates of Thomson’s model of the atom and mention its drawback.
Answer: According to Thomson’s model (1898), (i) an atom is a sphere of uniformly distributed positive charge; (ii) electrons are embedded in this sphere like seeds in a watermelon or plums in a pudding; (iii) the negative charge of the electrons exactly balances the positive charge, making the atom electrically neutral. The main drawback of this model was that it could not explain the results of Rutherford’s α-particle scattering experiment, especially the large-angle deflection and rebound of α-particles.
Q2. Write the observations and conclusions of Rutherford’s α-particle scattering experiment.
Answer: Observations: (i) Most α-particles passed straight through the gold foil. (ii) A small fraction was deflected through small angles. (iii) A very few (about 1 in 12,000) bounced straight back. Conclusions: (i) Most of the atom is empty space. (ii) The entire positive charge and nearly all the mass are concentrated in a very small central region — the nucleus. (iii) The size of the nucleus is extremely small compared to the size of the atom; electrons revolve around it.
Q3. State the main postulates of Bohr’s model of the atom.
Answer: Bohr’s postulates are: (i) Electrons revolve around the nucleus only in certain permitted circular paths called orbits or energy shells (K, L, M, N…). (ii) Each orbit corresponds to a definite amount of energy; hence orbits are also called stationary states. (iii) While revolving in these orbits, electrons do not radiate energy. (iv) Energy is absorbed or emitted only when an electron jumps from one orbit to another; the energy difference appears in the form of a photon.
Q4. State the Bohr–Bury rules for the distribution of electrons in shells.
Answer: (i) The maximum number of electrons in the n-th shell is given by the formula 2n² (so K = 2, L = 8, M = 18, N = 32). (ii) The outermost shell of any atom cannot contain more than 8 electrons. (iii) The penultimate (second-last) shell cannot contain more than 18 electrons. (iv) Electrons fill the shells in order of increasing energy; the next shell starts only after the previous one is suitably filled.
Q5. Write the electronic configuration and valency of (a) sodium (Z = 11) and (b) chlorine (Z = 17).
Answer: (a) Sodium (Na, Z = 11): K = 2, L = 8, M = 1. As there is 1 electron in the outermost shell, valency = 1. (b) Chlorine (Cl, Z = 17): K = 2, L = 8, M = 7. As there are 7 electrons in the outermost shell, valency = 8 − 7 = 1.
Q6. Distinguish between isotopes and isobars.
Answer: Isotopes are atoms of the same element having the same atomic number but different mass numbers (e.g., ¹²C and ¹⁴C). They have identical chemical properties but different physical properties. Isobars are atoms of different elements having the same mass number but different atomic numbers (e.g., ⁴⁰18Ar and ⁴⁰20Ca). They have different chemical properties as they are different elements.
5–6 Mark Questions
Q1. Describe Rutherford’s α-particle scattering experiment with the help of a labelled diagram. State the observations made and the conclusions drawn from this experiment. Mention one limitation of this model.
Answer: Experimental set-up: Rutherford directed a narrow beam of high-energy α-particles (helium nuclei) from a radioactive source onto a very thin gold foil (about 1000 atoms thick). A movable fluorescent screen of zinc sulphide was placed around the foil to detect the scattered particles. Whenever an α-particle struck the screen it produced a tiny flash of light.
Observations: (i) The majority of α-particles passed straight through the foil without any deflection. (ii) A small fraction of α-particles were deflected through small angles. (iii) A very few — about 1 in every 12,000 — were deflected through large angles, and some even bounced straight back along the path of incidence.
Conclusions: (i) Since most α-particles passed undeflected, the atom is largely empty space. (ii) Since only a few were deflected, the positive charge of the atom is concentrated in a very small region called the nucleus. (iii) Since some α-particles bounced back, the nucleus is very dense and contains almost the entire mass of the atom. (iv) The electrons revolve around the nucleus in circular paths.
Limitation: According to classical electromagnetic theory, a charged particle moving in a circular path should continuously emit energy, spiral into the nucleus, and the atom should collapse. Rutherford’s model could not explain the stability of the atom.
Q2. Explain Bohr’s model of the atom. How does it overcome the drawbacks of Rutherford’s model? Draw the orbital diagram for an atom of magnesium (Z = 12).
Answer: Bohr’s model (1913) modified Rutherford’s nuclear model with the following postulates: (i) Electrons revolve around the nucleus only in certain definite circular paths called orbits or stationary states, designated K, L, M, N… as we move outward from the nucleus. (ii) Each orbit is associated with a fixed amount of energy; hence the orbits are also called energy levels. (iii) As long as the electron stays in a particular orbit, it neither absorbs nor emits energy. This explains the stability of the atom. (iv) Energy is absorbed when the electron jumps to a higher orbit and emitted (as light of a definite frequency) when it falls to a lower orbit. The energy of the photon equals the difference in the energies of the two orbits.
By assuming that electrons in stationary orbits do not radiate energy, Bohr removed the major drawback of Rutherford’s model — the predicted collapse of the atom. The model also satisfactorily explained the line spectrum of the hydrogen atom.
Magnesium (Z = 12): Electronic configuration K = 2, L = 8, M = 2. Orbital diagram: nucleus (12 p, 12 n) at the centre; first ring (K) — 2 electrons; second ring (L) — 8 electrons; third ring (M) — 2 electrons. Valency = 2.
Q3. What are isotopes? Give two examples each of isotopes of hydrogen, carbon and chlorine. Mention any four important applications of isotopes.
Answer: Isotopes are atoms of the same element which have the same atomic number but different mass numbers. They occur because the nuclei contain the same number of protons but different numbers of neutrons. Isotopes have identical chemical properties but slightly different physical properties such as density and rate of diffusion.
Examples: Hydrogen — protium ¹1H, deuterium ²1H, tritium ³1H. Carbon — ¹²6C and ¹⁴6C. Chlorine — ³⁵17Cl and ³⁷17Cl. (The natural mixture of the two chlorine isotopes in the ratio 3 : 1 gives chlorine its average atomic mass of 35.5 u.)
Applications: (i) Carbon-14 (¹⁴C) is used in the radio-carbon dating of fossils, archaeological remains and old wooden objects. (ii) Cobalt-60 (⁶⁰Co) is used in the treatment of cancer (radiotherapy). (iii) Iodine-131 (¹³¹I) is used in the diagnosis and treatment of diseases of the thyroid gland. (iv) Uranium-235 (²³⁵U) is used as nuclear fuel in atomic reactors to generate electricity. (Other uses: phosphorus-31 in agricultural research, sodium-24 to detect blockage in blood vessels.)
Q4. An element X has atomic number 17 and mass number 35. Find (a) number of protons, (b) number of neutrons, (c) number of electrons, (d) electronic configuration, (e) valency, and (f) name another element which is an isotope of X.
Answer: Given Z = 17 and A = 35. (a) Number of protons = Z = 17. (b) Number of neutrons = A − Z = 35 − 17 = 18. (c) Number of electrons = 17 (atom is neutral). (d) Distributing 17 electrons by 2n² and outer-shell rules: K = 2, L = 8, M = 7. (e) Outermost shell has 7 electrons, hence valency = 8 − 7 = 1. (f) The element is chlorine (Cl). Its other naturally occurring isotope is ³⁷17Cl, which has 17 protons and 20 neutrons.
Q5. Write the electronic configurations and valencies of the first 18 elements. Show your work in a neat table.
Answer: Using the 2n² rule and the outer-shell limit of 8, the first 18 elements are arranged as follows:
| Z | Element | Symbol | K | L | M | Valency |
|---|---|---|---|---|---|---|
| 1 | Hydrogen | H | 1 | — | — | 1 |
| 2 | Helium | He | 2 | — | — | 0 |
| 3 | Lithium | Li | 2 | 1 | — | 1 |
| 4 | Beryllium | Be | 2 | 2 | — | 2 |
| 5 | Boron | B | 2 | 3 | — | 3 |
| 6 | Carbon | C | 2 | 4 | — | 4 |
| 7 | Nitrogen | N | 2 | 5 | — | 3 |
| 8 | Oxygen | O | 2 | 6 | — | 2 |
| 9 | Fluorine | F | 2 | 7 | — | 1 |
| 10 | Neon | Ne | 2 | 8 | — | 0 |
| 11 | Sodium | Na | 2 | 8 | 1 | 1 |
| 12 | Magnesium | Mg | 2 | 8 | 2 | 2 |
| 13 | Aluminium | Al | 2 | 8 | 3 | 3 |
| 14 | Silicon | Si | 2 | 8 | 4 | 4 |
| 15 | Phosphorus | P | 2 | 8 | 5 | 3 |
| 16 | Sulphur | S | 2 | 8 | 6 | 2 |
| 17 | Chlorine | Cl | 2 | 8 | 7 | 1 |
| 18 | Argon | Ar | 2 | 8 | 8 | 0 |
Additional Practice — Multiple Choice Questions
Q1. The electron was discovered by —
(a) Rutherford (b) Goldstein (c) J. J. Thomson (d) Chadwick
Answer: (c) J. J. Thomson.
Q2. The neutral particle present in the nucleus is —
(a) electron (b) proton (c) neutron (d) positron
Answer: (c) neutron.
Q3. The maximum number of electrons in the M shell is —
(a) 2 (b) 8 (c) 18 (d) 32
Answer: (c) 18.
Q4. Rutherford used which particle in his scattering experiment?
(a) electron (b) α-particle (c) β-particle (d) γ-ray
Answer: (b) α-particle.
Q5. The atomic number of an element is equal to the number of —
(a) neutrons (b) protons (c) nucleons (d) ions
Answer: (b) protons.
Q6. Which of the following pairs are isotopes?
(a) ⁴⁰18Ar and ⁴⁰20Ca (b) ¹²6C and ¹⁴6C (c) ²³11Na and ²⁴12Mg (d) None
Answer: (b) ¹²6C and ¹⁴6C.
Q7. The valency of an element with electronic configuration 2, 8, 6 is —
(a) 6 (b) 4 (c) 2 (d) 0
Answer: (c) 2 (since 8 − 6 = 2).
Q8. Which isotope is used in radio-carbon dating?
(a) ¹²C (b) ¹⁴C (c) ¹³¹I (d) ⁶⁰Co
Answer: (b) ¹⁴C.
Q9. The model that compared the atom with a watermelon was given by —
(a) Bohr (b) Rutherford (c) Thomson (d) Dalton
Answer: (c) Thomson.
Q10. The mass number of an atom is the sum of —
(a) protons and electrons (b) electrons and neutrons (c) protons and neutrons (d) all subatomic particles
Answer: (c) protons and neutrons.
Fill in the Blanks
Q1. The maximum number of electrons in the n-th shell is given by the formula __________.
Answer: 2n².
Q2. The outermost shell of any atom can hold a maximum of __________ electrons.
Answer: 8.
Q3. The positively charged rays observed by Goldstein are called __________ rays.
Answer: canal (or anode) rays.
Q4. Atoms of different elements having the same mass number are called __________.
Answer: isobars.
Q5. The radioactive isotope used to treat thyroid disorders is __________.
Answer: iodine-131 (¹³¹I).
True or False
Q1. Most of the space inside an atom is empty.
Answer: True.
Q2. Electrons revolve around the nucleus in fixed circular orbits according to Bohr.
Answer: True.
Q3. Isotopes of an element have different chemical properties.
Answer: False. Isotopes have identical chemical properties; only their physical properties differ.
Q4. The atomic number of chlorine is 35.
Answer: False. The atomic number of chlorine is 17; 35 is its mass number.
Q5. The neutron has a mass nearly equal to that of an electron.
Answer: False. The neutron has a mass nearly equal to that of a proton (about 1 u), which is about 1836 times the mass of an electron.
Glossary
| Term | Meaning |
|---|---|
| Cathode rays | Streams of negatively charged particles (electrons) emitted from the cathode in a discharge tube at low pressure. |
| Canal rays | Positively charged rays observed by Goldstein, which led to the discovery of the proton. |
| Electron | A negatively charged subatomic particle (charge −1, mass ≈ 1/1836 u) revolving around the nucleus. |
| Proton | A positively charged subatomic particle (charge +1, mass ≈ 1 u) present inside the nucleus. |
| Neutron | A neutral subatomic particle (charge 0, mass ≈ 1 u) present inside the nucleus; discovered by Chadwick. |
| Nucleus | The tiny, dense, positively charged centre of an atom containing protons and neutrons. |
| Atomic number (Z) | Number of protons in the nucleus of an atom; equals the number of electrons in a neutral atom. |
| Mass number (A) | Total number of protons and neutrons in the nucleus of an atom. |
| Energy shell / orbit | A fixed circular path (K, L, M, N…) in which an electron revolves around the nucleus, having a definite energy. |
| 2n² rule | The maximum number of electrons that can occupy the n-th shell of an atom. |
| Valency | The combining capacity of an element, decided by the number of electrons in its outermost shell. |
| Isotopes | Atoms of the same element having the same atomic number but different mass numbers (e.g., ¹²C, ¹⁴C). |
| Isobars | Atoms of different elements having the same mass number but different atomic numbers (e.g., ⁴⁰Ar, ⁴⁰Ca). |
| α-particle | A helium nucleus (2 protons + 2 neutrons) used by Rutherford as a projectile in his scattering experiment. |
| Plum-pudding model | Thomson’s model in which electrons are embedded in a sphere of positive charge. |
| Nuclear model | Rutherford’s model in which the entire positive charge and mass are concentrated in a small central nucleus. |
| Bohr’s model | Atomic model in which electrons revolve in fixed energy shells without radiating energy. |