Class 9 Science NCERT Notes – Chapter 4: The Structure of the Atom (PDF, MindMap, Q&A, Quizzes)

Chapter 4 (Chemistry): The Structure of the Atom – Class 9 NCERT Science Detailed Study Notes.

1. The Dawn of Subatomic Particles

At the end of the 19th century, the prevailing scientific view, based on Dalton’s atomic theory, was that atoms were indivisible and indestructible. However, a series of experiments concerning static electricity and electrical conduction began to challenge this idea, providing the first indications that atoms were composed of smaller constituents.

1.1 Discovery of the Electron

• Scientist: J.J. Thomson is credited with the identification of the electron.
• Significance: By 1900, it was established that the atom was not indivisible but contained at least one subatomic particle, the electron.
• Properties:
  • Represented as e⁻.
  • Charge is considered minus one (-1).
  • Mass is considered negligible. It is approximately 1/2000 times the mass of a hydrogen atom.
  • Electrons can be removed from atoms more easily than protons.

1.2 Discovery of the Proton

• Scientist: In 1886, E. Goldstein discovered new radiations in a gas discharge which he called canal rays.
• Canal Rays: These were positively charged radiations that ultimately led to the discovery of the proton.
• Properties:
  • Represented as p⁺.
  • Charge is equal in magnitude but opposite in sign to that of the electron, considered plus one (+1).
  • Mass is approximately 2000 times that of an electron and is taken as one unit.
  • Protons appeared to be located in the interior of the atom.

1.3 Discovery of the Neutron

• Scientist: J. Chadwick discovered the neutron in 1932.
• Properties:
  • Represented as ‘n’.
  • Has no charge (neutral).
  • Mass is nearly equal to that of a proton.
  • Present in the nucleus of all atoms, with the exception of the most common form of hydrogen (protium).

2. Models of Atomic Structure

The discovery of protons and electrons created a new challenge: determining how these particles were arranged within the atom. This led to the proposal of several atomic models.

2.1 Thomson’s Model of an Atom (Christmas Pudding Model)

• Proposed by: J.J. Thomson.
• Analogy: Likened to a Christmas pudding or a watermelon.
• Postulates:
  1. An atom consists of a positively charged sphere.
  2. Electrons (like currants in a pudding or seeds in a watermelon) are embedded within this positive sphere.
  3. The negative and positive charges are equal in magnitude, making the atom as a whole electrically neutral.
• Limitation: While it successfully explained the electrical neutrality of atoms, it could not explain the results of later experiments, such as Rutherford’s scattering experiment.

2.2 Rutherford’s Nuclear Model of an Atom

• Background: Ernest Rutherford designed an experiment to understand the arrangement of electrons within an atom.
• The Gold Foil Experiment:
  • Setup: Fast-moving, doubly-charged helium ions, known as alpha (α)-particles, were directed at a very thin gold foil (about 1000 atoms thick).
  • Expectation: Since α-particles (mass of 4 u) were much heavier than protons, Rutherford expected only small deflections.
  • Observations (Unexpected Results):
    1. Most α-particles passed straight through the foil undeflected.
    2. A small fraction of α-particles were deflected by small angles.
    3. Surprisingly, about 1 in 12,000 particles appeared to rebound (deflected by 180°).
• Conclusions from the Experiment:
  1. Most of the space inside an atom is empty.
  2. The positive charge of the atom occupies a very small volume.
  3. All the positive charge and nearly all the mass of the atom are concentrated in a very small, dense region within the atom.
• Features of the Nuclear Model:
  1. There is a positively charged centre in an atom called the nucleus. Nearly all the mass of an atom resides here.
  2. Electrons revolve around the nucleus in circular paths (orbits).
  3. The size of the nucleus is very small compared to the size of the atom (radius of the nucleus is about 10⁵ times less than the atom’s radius).
• Major Drawback: According to classical physics, a charged particle (like an electron) in a circular orbit would undergo acceleration and radiate energy. This would cause the revolving electron to lose energy, spiral inwards, and eventually fall into the nucleus, making the atom highly unstable. However, atoms are known to be stable.

2.3 Bohr’s Model of an Atom

• Objective: To overcome the stability objection raised against Rutherford’s model.
• Proposed by: Niels Bohr.
• Postulates:
  1. Only certain special orbits, known as discrete orbits, are allowed for electrons inside the atom.
  2. While revolving in these discrete orbits, electrons do not radiate energy.
• Energy Levels (Shells):
  • These discrete orbits are also called energy levels or shells.
  • They are represented by the letters K, L, M, N,… or by the numbers n = 1, 2, 3, 4,….

3. Electron Distribution and Valency

3.1 Bohr-Bury Rules for Electron Distribution

The distribution of electrons into different shells follows a set of rules:

1. Maximum Electrons per Shell: The maximum number of electrons that can be present in a shell is given by the formula 2n², where ‘n’ is the orbit number or energy level index.
   • K-shell (n=1): 2(1)² = 2 electrons
   • L-shell (n=2): 2(2)² = 8 electrons
   • M-shell (n=3): 2(3)² = 18 electrons
   • N-shell (n=4): 2(4)² = 32 electrons
2. Outermost Shell Capacity: The maximum number of electrons that can be accommodated in the outermost orbit is 8. This is known as the octet rule.
3. Step-wise Filling: Electrons do not occupy a new shell until the inner shells are completely filled.

3.2 Valency

• Valence Electrons: The electrons present in the outermost shell of an atom.
• Octet: An outermost shell containing eight electrons, which is a very stable configuration. Atoms with a complete octet (like Neon and Argon) show little chemical activity and have a valency of zero.
• Combining Capacity (Valency): An atom’s tendency to react is driven by its attempt to achieve a fully-filled outermost shell (an octet). This is done by gaining, losing, or sharing electrons. Valency is the number of electrons an atom gains, loses, or shares to achieve this stable configuration.
  • Losing Electrons: If an atom has 1, 2, or 3 valence electrons (e.g., Sodium-1, Magnesium-2, Aluminium-3), it is easier to lose them. Their valencies are 1, 2, and 3 respectively.
  • Gaining Electrons: If an atom’s outermost shell is nearly full (e.g., Fluorine-7, Oxygen-6, Chlorine-7), it is easier to gain electrons. Valency is calculated as (8 – number of valence electrons). For Fluorine, valency is 8 – 7 = 1.

4. Atomic Number, Mass Number, and Related Concepts

4.1 Atomic Number (Z)

• Definition: The total number of protons present in the nucleus of an atom.
• Significance: All atoms of a given element have the same atomic number. It is the defining characteristic of an element.
• Example: For Carbon, Z = 6, meaning every carbon atom has 6 protons.

4.2 Mass Number (A)

• Nucleons: Protons and neutrons are collectively called nucleons as they reside in the nucleus.
• Definition: The sum of the total number of protons and neutrons in the nucleus of an atom.
• Significance: Represents the approximate mass of the atom, as the mass of electrons is negligible.
• Example: Carbon with 6 protons and 6 neutrons has a mass number of 12.
• Notation: ᴬZ X, where X is the element symbol, A is the mass number, and Z is the atomic number. E.g., ¹⁴₇N.

4.3 Isotopes

• Definition: Atoms of the same element that have the same atomic number (same number of protons) but different mass numbers (different number of neutrons).
• Properties: Isotopes of an element have similar chemical properties but different physical properties.
• Examples:
  • Hydrogen: Protium (¹H), Deuterium (²H), Tritium (³H). All have 1 proton; they have 0, 1, and 2 neutrons respectively.
  • Carbon: ¹²C and ¹⁴C.
  • Chlorine: ³⁵Cl and ³⁷Cl.
• Average Atomic Mass: For elements that exist as a mixture of isotopes, the atomic mass is the weighted average of the masses of its naturally occurring isotopes. For example, chlorine occurs as ³⁵Cl (75%) and ³⁷Cl (25%), giving it an average atomic mass of 35.5 u.
• Applications of Isotopes:
  • Uranium: An isotope is used as fuel in nuclear reactors.
  • Cobalt: An isotope is used in the treatment of cancer.
  • Iodine: An isotope is used in the treatment of goitre.

4.4 Isobars

• Definition: Atoms of different elements that have different atomic numbers but the same mass number.
• Example: Calcium (Atomic Number 20, Mass Number 40) and Argon (Atomic Number 18, Mass Number 40). Both have a mass number of 40 but are different elements.

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Q&A Section

Short-Answer Questions (Answer in 2-3 sentences)

1. What discovery provided the first indication that atoms are not indivisible?
2. Who discovered canal rays, and what did this discovery lead to?
3. Describe the properties of a proton in terms of its charge and mass relative to an electron.
4. What are the two main postulates of J.J. Thomson’s model of the atom?
5. What analogy did Thomson use to describe his atomic model?
6. What are alpha (α)-particles, and what were they used for in Rutherford’s experiment?
7. Why did Rutherford select a gold foil for his scattering experiment?
8. State the most surprising observation from Rutherford’s gold foil experiment.
9. Based on his experiment, what did Rutherford conclude about the distribution of mass and positive charge in an atom?
10. What is the major drawback of Rutherford’s nuclear model?
11. How did Neils Bohr’s model address the instability issue of Rutherford’s model?
12. Who discovered the neutron, and what are its key properties?
13. What is the formula used to determine the maximum number of electrons in a given shell?
14. According to the Bohr-Bury rules, what is the maximum number of electrons that can be accommodated in the outermost shell?
15. What are valence electrons?
16. How is the valency of an element like sodium (with 1 valence electron) determined?
17. How is the valency of an element like fluorine (with 7 valence electrons) determined?
18. Define Atomic Number (Z). What does it represent?
19. Define Mass Number (A). What particles contribute to it?
20. What is the collective name for protons and neutrons, and why are they called that?
21. Define isotopes, and provide one example with two specific isotopes.
22. Do isotopes of an element have similar or different chemical properties? Explain why.
23. State two applications of isotopes in medicine or energy.
24. Define isobars and provide a specific example of a pair of isobars.
25. An atom of sodium has an atomic number of 11 and a mass number of 23. How many protons, neutrons, and electrons does a neutral sodium atom have?

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Multiple-Choice Questions (MCQ)

1. Who is credited with the discovery of the electron? a) E. Goldstein b) J.J. Thomson c) J. Chadwick d) E. Rutherford
2. Canal rays, discovered by E. Goldstein, are streams of: a) Negatively charged particles b) Neutral particles c) Positively charged radiations d) Alpha-particles
3. According to Rutherford’s experiment, the size of the nucleus is approximately how many times smaller than the radius of the atom? a) 100 times b) 1,000 times c) 100,000 times (10⁵) d) 1,000,000 times (10⁶)
4. The mass of a proton is taken as one unit, and its charge is: a) Plus one b) Minus one c) Neutral d) Plus two
5. The major drawback of Rutherford’s model was its inability to explain: a) The existence of a nucleus b) The atom’s electrical neutrality c) The atom’s stability d) The results of the gold foil experiment
6. In Bohr’s model of the atom, electrons do not radiate energy while revolving in: a) Any circular path b) Discrete orbits c) The nucleus d) Elliptical orbits
7. The discovery of the neutron is credited to: a) Neils Bohr b) J.J. Thomson c) E. Rutherford d) J. Chadwick
8. What is the maximum number of electrons that the L-shell (n=2) can hold? a) 2 b) 8 c) 18 d) 32
9. The electrons present in the outermost shell of an atom are called: a) Nucleons b) Core electrons c) Valence electrons d) Octet electrons
10. An element with 8 electrons in its outermost shell is said to possess an octet and has a valency of: a) 1 b) 8 c) 2 d) 0
11. The atomic number (Z) of an element is defined by the number of: a) Neutrons b) Protons c) Electrons d) Nucleons
12. The mass number (A) of an atom is the sum of: a) Protons and electrons b) Neutrons and electrons c) Protons and neutrons d) Only protons
13. The species ¹²₆C and ¹⁴₆C are examples of: a) Isobars b) Isotopes c) Isomers d) Allotropes
14. Atoms of different elements with the same mass number but different atomic numbers are known as: a) Isotopes b) Allotropes c) Isotones d) Isobars
15. An isotope of which element is used in the treatment of goitre? a) Uranium b) Cobalt c) Iodine d) Carbon
16. What is the valency of magnesium (Atomic Number 12)? a) 1 b) 2 c) 3 d) 4
17. A neutral Helium atom (Atomic Mass 4 u, 2 protons) contains how many neutrons? a) 4 b) 0 c) 2 d) 6
18. In the notation ¹⁴₇N, the number 14 represents the: a) Atomic Number b) Number of neutrons c) Number of electrons d) Mass Number
19. In the Rutherford scattering experiment, the fact that most α-particles passed straight through the gold foil indicated that: a) The atom is electrically neutral b) The nucleus is positively charged c) Most of the space in an atom is empty d) Electrons are negatively charged
20. Which of the following is the correct electronic configuration for a neutral sodium atom (Z=11)? a) 2, 8 b) 8, 2, 1 c) 2, 1, 8 d) 2, 8, 1

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Answer Keys

Short-Answer Questions – Answer Key

1. The first indications that atoms are not indivisible came from studying static electricity and the conditions under which different substances conduct electricity. These phenomena suggested the existence of smaller, charged particles within the atom.
2. E. Goldstein discovered canal rays in 1886. This discovery of positively charged radiations in a gas discharge ultimately led to the identification of the proton, a fundamental subatomic particle.
3. A proton has a charge that is equal in magnitude but opposite in sign to an electron’s charge (+1). Its mass is approximately 2000 times that of an electron and is considered to be one unit.
4. Thomson’s model proposed that (i) an atom consists of a positively charged sphere with electrons embedded in it, and (ii) the negative and positive charges are equal in magnitude, making the atom electrically neutral as a whole.
5. Thomson used the analogies of a Christmas pudding, where electrons are like currants in a positive sphere, or a watermelon, where electrons are like seeds studded in the red, positively charged edible part.
6. Alpha (α)-particles are doubly-charged helium ions with a mass of 4 u. Rutherford used these fast-moving particles in his experiment by directing them at a thin gold foil to probe the structure of the atom.
7. Rutherford selected a gold foil because he wanted the layer to be as thin as possible. The gold foil used was only about 1000 atoms thick, which was crucial for observing the scattering effects on individual atoms.
8. The most surprising observation was that one out of every 12,000 alpha-particles appeared to rebound, being deflected by 180 degrees. Rutherford described this as being as incredible as firing a shell at tissue paper and having it bounce back.
9. Rutherford concluded that all the positive charge and nearly all the mass of an atom were concentrated in a very small, dense volume within the atom, which he called the nucleus. He also concluded that most of the atom is empty space.
10. The major drawback is that it could not explain the stability of the atom. A revolving electron in a circular orbit should accelerate, radiate energy, and spiral into the nucleus, causing the atom to collapse, which does not happen.
11. Neils Bohr proposed that electrons do not radiate energy while revolving in certain special orbits called discrete orbits or energy levels. This postulate directly countered the classical physics objection that caused Rutherford’s model to be seen as unstable.
12. J. Chadwick discovered the neutron in 1932. The neutron has no charge (it is neutral) and has a mass nearly equal to that of a proton.
13. The maximum number of electrons in a given shell is determined by the formula 2n², where ‘n’ represents the orbit number or energy level index (e.g., n=1 for K-shell, n=2 for L-shell).
14. According to the Bohr-Bury rules, the maximum number of electrons that can be accommodated in the outermost shell of an atom is 8. This is also known as the octet rule.
15. Valence electrons are the electrons that are present in the outermost shell of an atom. These electrons are the ones that participate in chemical reactions.
16. Sodium has one valence electron in its outermost shell. To achieve a stable octet, it is easier for it to lose this single electron than to gain seven. Therefore, its valency is one.
17. A fluorine atom has seven electrons in its outermost shell. It is easier for it to gain one electron to complete its octet than to lose all seven. Its valency is determined by subtracting the number of valence electrons from 8 (8 – 7 = 1), giving it a valency of one.
18. The Atomic Number (Z) is defined as the total number of protons present in the nucleus of an atom. It is the unique identifier for an element.
19. The Mass Number (A) is defined as the sum of the total number of protons and neutrons present in the nucleus of an atom. These particles, protons and neutrons, contribute practically all of the atom’s mass.
20. Protons and neutrons are collectively called nucleons. They are given this name because they are both located in the atom’s nucleus.
21. Isotopes are atoms of the same element that have the same atomic number but different mass numbers. An example is the isotopes of carbon: Carbon-12 (¹²₆C) and Carbon-14 (¹⁴₆C), which both have 6 protons but have 6 and 8 neutrons, respectively.
22. Isotopes of an element have similar chemical properties. This is because chemical properties are determined by the number and arrangement of electrons, particularly the valence electrons, which is the same for all isotopes of an element since they have the same number of protons.
23. In medicine, an isotope of cobalt is used in cancer treatment, and an isotope of iodine is used to treat goitre. In energy, an isotope of uranium is used as a fuel in nuclear reactors.
24. Isobars are atoms of different elements which have different atomic numbers but the same mass number. An example is Calcium (Atomic Number 20, Mass Number 40) and Argon (Atomic Number 18, Mass Number 40).
25. For a neutral sodium atom, the number of protons equals the atomic number (11), and the number of electrons equals the number of protons (11). The number of neutrons is the mass number minus the atomic number (23 – 11 = 12).

Multiple-Choice Questions (MCQ) – Answer Key

1. b) J.J. Thomson
2. c) Positively charged radiations
3. c) 100,000 times (10⁵)
4. a) Plus one
5. c) The atom’s stability
6. b) Discrete orbits
7. d) J. Chadwick
8. b) 8
9. c) Valence electrons
10. d) 0
11. b) Protons
12. c) Protons and neutrons
13. b) Isotopes
14. d) Isobars
15. c) Iodine
16. b) 2
17. c) 2
18. d) Mass Number
19. c) Most of the space in an atom is empty
20. d) 2, 8, 1

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Suggested Essay Questions

1. Compare and contrast the atomic models of J.J. Thomson, Ernest Rutherford, and Niels Bohr. Discuss the key experiments or reasoning that led to the development of each model and highlight the limitations that each successive model sought to address.

Answer: J.J. Thomson’s model, the first to propose subatomic structure, envisioned the atom as a “plum pudding” — a sphere of uniform positive charge with negatively charged electrons embedded within it. This model successfully explained the overall electrical neutrality of atoms. However, it was a static model with no concept of a nucleus.

Ernest Rutherford’s model was a direct result of his gold foil experiment. The surprising rebound of a few alpha-particles led him to conclude that the atom’s positive charge and mass were concentrated in a tiny, dense center called the nucleus, with electrons orbiting it like planets. This nuclear model replaced Thomson’s diffuse positive sphere but had a major flaw: classical physics predicted that an orbiting electron would radiate energy and spiral into the nucleus, making the atom unstable.

Niels Bohr addressed Rutherford’s stability problem by introducing quantum concepts. He postulated that electrons could only exist in specific “discrete orbits” or energy levels and that they would not radiate energy while in these stable orbits. This model successfully explained atomic stability and the line spectra of hydrogen. While a significant improvement, Bohr’s model worked best for simple atoms and was later superseded by more complex quantum mechanical models. Each model was a critical step, building upon the experimental evidence and theoretical shortcomings of its predecessor.

2. Describe Rutherford’s alpha-particle scattering experiment in detail. Include the experimental setup, the three main observations, and the three corresponding conclusions that led to the formulation of the nuclear model of the atom.

Answer: Rutherford’s experiment was designed to probe the internal structure of the atom. The setup involved a source of fast-moving alpha-particles (doubly-charged helium ions) aimed at an extremely thin sheet of gold foil. A circular screen was placed around the foil to detect the path of the alpha-particles after they interacted with the gold atoms.

The experiment yielded three main, largely unexpected observations:

1. Most of the alpha-particles passed straight through the gold foil without any deflection.
2. A small fraction of the alpha-particles were deflected from their original path by small angles.
3. A very small number of alpha-particles (about 1 in 12,000) bounced back, deflected by 180°.

From these observations, Rutherford drew three profound conclusions:

1. Since most particles passed through undeflected, most of the space inside an atom must be empty.
2. The small deflections indicated that the positive charge of the atom was concentrated in a very small volume, which repelled the positively charged alpha-particles.
3. The rare, large-angle deflections (rebounding) meant that the atom’s positive charge and almost all its mass were concentrated in a dense central core, later named the nucleus. This reasoning formed the basis of his revolutionary nuclear model of the atom.

3. Explain the concept of valency and its relationship to the Bohr-Bury rules for electron distribution. Use the examples of Sodium (Na, Z=11), Chlorine (Cl, Z=17), and Neon (Ne, Z=10) to illustrate your explanation.

Answer: Valency is the combining capacity of an atom, determined by its tendency to achieve a stable electron configuration, typically a full outermost shell of 8 electrons (an octet). The Bohr-Bury rules govern how electrons fill shells (2, 8, 18…) and state that the outermost shell can hold a maximum of 8 electrons. An atom’s valency is the number of electrons it must gain, lose, or share to attain this stable octet.

• Sodium (Na, Z=11): Its electron distribution is 2, 8, 1. It has one valence electron. To achieve an octet, it is far easier to lose this one electron than to gain seven. By losing one electron, its valency becomes 1.
• Chlorine (Cl, Z=17): Its electron distribution is 2, 8, 7. It has seven valence electrons. It is much easier to gain one electron to complete its octet than to lose all seven. Therefore, its valency is 1 (determined by 8 – 7).
• Neon (Ne, Z=10): Its electron distribution is 2, 8. Its outermost shell is already completely filled with 8 electrons, possessing a stable octet. It has no tendency to gain, lose, or share electrons, so its combining capacity, or valency, is zero, making it an inert element.

4. Define and differentiate between isotopes and isobars. Provide specific, named examples for each, and explain why isotopes of an element share chemical properties while isobars do not.

Answer: Isotopes are atoms of the same element that have the same atomic number (Z) but different mass numbers (A). This means they have the same number of protons but a different number of neutrons. For example, the isotopes of hydrogen are Protium (¹H), Deuterium (²H), and Tritium (³H); all have one proton, but 0, 1, and 2 neutrons, respectively.

Isobars, on the other hand, are atoms of different elements that have different atomic numbers but the same mass number. This means they have a different number of protons but the same total number of nucleons (protons + neutrons). An example is Argon-40 (¹⁸Ar) and Calcium-40 (²⁰Ca). Argon has 18 protons and 22 neutrons, while Calcium has 20 protons and 20 neutrons, but both have a mass number of 40.

The key difference in their properties stems from their electronic structure. Isotopes of an element share chemical properties because chemical behaviour is determined by the electron configuration, which in turn is dictated by the number of protons (atomic number). Since all isotopes of an element have the same number of protons and thus the same number of electrons, they react chemically in the same way. Isobars, being different elements with different atomic numbers, have different numbers of electrons and thus different electron configurations, leading to entirely different chemical properties.

5. How is the mass of an atom determined? Explain the concepts of Atomic Number (Z) and Mass Number (A), and describe how to calculate the average atomic mass of an element that exists in nature as a mixture of isotopes, using Chlorine (³⁵Cl at 75% and ³⁷Cl at 25%) as an example.

Answer: The mass of an atom is practically due entirely to the protons and neutrons in its nucleus, as the mass of electrons is negligible. The Atomic Number (Z) represents the number of protons and defines the element. The Mass Number (A) represents the total number of protons and neutrons (nucleons) in the nucleus.

For an element with no isotopes, the atomic mass would simply be the sum of the masses of its protons and neutrons. However, many elements exist as a mixture of isotopes. In this case, the atomic mass listed on the periodic table is the weighted average mass of all naturally occurring atoms of that element. To calculate this average, one must know the mass and natural abundance (percentage) of each isotope.

Using chlorine as an example: It exists as ³⁵Cl (75% abundance) and ³⁷Cl (25% abundance). The average atomic mass is calculated as follows: (Mass of Isotope 1 × Abundance of Isotope 1) + (Mass of Isotope 2 × Abundance of Isotope 2) = (35 u × 75/100) + (37 u × 25/100) = (35 × 0.75) + (37 × 0.25) = 26.25 + 9.25 = 35.5 u. This means that in any natural sample of chlorine, the average mass of an atom is 35.5 u.

6. Trace the historical development of the understanding of atomic structure from Dalton’s initial theory to the discovery of the three main subatomic particles. Who were the key scientists involved and what were their primary contributions?

Answer: The modern understanding of atomic structure evolved through several key discoveries. Initially, Dalton’s atomic theory suggested atoms were indivisible, fundamental building blocks of matter. This view held until the late 19th century. The first subatomic particle was hinted at by E. Goldstein in 1886, who discovered “canal rays,” which were positively charged radiations. This discovery paved the way for the later identification of the proton (p⁺), the positively charged particle in the atom’s interior. Credit for the discovery of the electron (e⁻) goes to J.J. Thomson. Around 1900, his work established that atoms contained these small, negatively charged particles, definitively proving that atoms were divisible. This led to his “plum pudding” model. The final key component, the neutron (n), was discovered much later, in 1932, by J. Chadwick. He identified a particle in the nucleus with a mass nearly equal to a proton but with no electrical charge. This discovery completed the picture of the three primary subatomic particles: the proton, neutron, and electron.

7. Discuss the rules for the distribution of electrons in different orbits as proposed by Bohr and Bury. Apply these rules to determine the electron configuration for Aluminium (Z=13) and Argon (Z=18).

Answer: The distribution of electrons into different orbits or shells, as suggested by Bohr and Bury, follows three main rules. First, the maximum number of electrons that can occupy a shell is given by the formula 2n², where ‘n’ is the shell number (1 for K, 2 for L, etc.). This gives capacities of 2, 8, 18, and 32 for the K, L, M, and N shells, respectively. Second, the outermost shell of an atom cannot accommodate more than 8 electrons. Third, shells are filled in a stepwise manner, meaning a new shell does not begin to fill until the inner shells are completely filled.

• Aluminium (Al, Z=13): It has 13 electrons. Following the rules, the first shell (K, n=1) takes a maximum of 2 electrons. The second shell (L, n=2) takes a maximum of 8 electrons. So far, 2 + 8 = 10 electrons are placed. The remaining 3 electrons go into the third shell (M, n=3). Thus, the electronic configuration for Aluminium is 2, 8, 3.
• Argon (Ar, Z=18): It has 18 electrons. The K-shell is filled with 2 electrons. The L-shell is filled with 8 electrons. This accounts for 10 electrons. The remaining 8 electrons go into the M-shell. This fills the outermost shell to its capacity of 8, making Argon stable. Thus, the electronic configuration for Argon is 2, 8, 8.

8. Explain what would happen if Rutherford’s alpha-particle scattering experiment were carried out using a foil of a metal other than gold. Consider a much lighter metal like lithium and a much heavier metal like lead. How might the observations change?

Answer: The fundamental observations of Rutherford’s experiment would remain the same regardless of the metal foil used, but the quantitative results would differ. The general pattern of most alpha-particles passing straight through, some deflecting slightly, and a very few deflecting by large angles would persist, as all atoms are mostly empty space with a dense, positive nucleus.

• Using a Lighter Metal (e.g., Lithium): A lithium nucleus is much smaller and has a much lower positive charge (+3) than a gold nucleus (+79). The electrostatic repulsion between the positive alpha-particle and the lithium nucleus would be significantly weaker. Consequently, the alpha-particles would experience smaller deflections on average. The number of particles deflected by large angles would be substantially lower than with gold.
• Using a Heavier Metal (e.g., Lead): A lead nucleus is heavier and has a greater positive charge (+82) than gold. The repulsive force would be stronger. This would result in slightly greater average deflections and a marginally higher number of particles scattering at large angles compared to the gold foil experiment. In essence, the heaviness and high positive charge of the nucleus are key to causing the dramatic large-angle scattering.

9. The average atomic mass of a sample of element X is 16.2 u. The sample contains two isotopes, ¹⁶₈X and ¹⁸₈X. Calculate the percentage of each isotope in the sample.

Answer: To calculate the percentage of each isotope, we can use an algebraic approach. Let the percentage of the ¹⁶X isotope be ‘p’. Then, the percentage of the ¹⁸X isotope must be (100 – p). The average atomic mass is the weighted average of the isotopic masses.

The formula for average atomic mass is: Average Mass = [(Mass of Isotope 1) × (% Abundance 1) + (Mass of Isotope 2) × (% Abundance 2)] / 100

Substituting the given values: 16.2 = [16 × p + 18 × (100 – p)] / 100

Now, we solve for ‘p’: 16.2 × 100 = 16p + 1800 – 18p 1620 = 1800 – 2p 2p = 1800 – 1620 2p = 180 p = 90

Therefore, the percentage of the ¹⁶₈X isotope is 90%. The percentage of the ¹⁸₈X isotope is 100 – 90 = 10%. The sample contains 90% of ¹⁶₈X and 10% of ¹⁸₈X.

10. A sodium ion, Na⁺, has completely filled K and L shells. Explain this statement based on the atomic structure of a neutral sodium atom (Z=11).

Answer: A neutral sodium atom has an atomic number (Z) of 11, which means it has 11 protons in its nucleus and 11 electrons orbiting it. According to the Bohr-Bury rules, these 11 electrons are distributed in the shells as follows: 2 electrons in the K-shell (its maximum), 8 electrons in the L-shell (its maximum), and the remaining 1 electron in the M-shell. The electronic configuration is 2, 8, 1.

The M-shell is the outermost shell and contains only one valence electron. To achieve a stable electron configuration (a full outer shell), the sodium atom has a strong tendency to lose this single valence electron. When it loses one electron, it becomes a positively charged ion, denoted as Na⁺.

After losing one electron, the Na⁺ ion has only 10 electrons left. The electron configuration is now 2, 8. The K-shell is full with 2 electrons, and what was previously the inner L-shell is now the outermost shell, which is completely filled with 8 electrons. This configuration is very stable, which is why the statement that the Na⁺ ion has completely filled K and L shells is correct.

Glossary of Key Terms

Term	Definition
Alpha (α)-particle	A doubly-charged helium ion (He²⁺) with a mass of 4 u. These fast-moving particles were used by Rutherford in his gold foil scattering experiment.
Atom	The fundamental building block of matter. It consists of a nucleus containing protons and neutrons, with electrons revolving around the nucleus in discrete orbits.
Atomic Number (Z)	The total number of protons present in the nucleus of an atom. It is the defining characteristic of an element.
Canal Rays	Positively charged radiations discovered by E. Goldstein in a gas discharge tube, which led to the discovery of the proton.
Discrete Orbits	Certain special orbits or shells within an atom where electrons can revolve without radiating energy, as proposed by Neils Bohr. Also known as energy levels.
Electron (e⁻)	A negatively charged subatomic particle with a negligible mass (approx. 1/2000 of a proton’s mass) that revolves around the nucleus of an atom.
Energy Levels	The specific orbits or shells (designated K, L, M, N,…) in which electrons are distributed around the nucleus. Each level corresponds to a fixed energy state.
Isobars	Atoms of different elements having different atomic numbers but the same mass number. For example, Argon-40 and Calcium-40.
Isotopes	Atoms of the same element having the same atomic number but different mass numbers. They have the same number of protons but different numbers of neutrons.
Mass Number (A)	The sum of the total number of protons and neutrons present in the nucleus of an atom.
Neutron (n)	A subatomic particle with no electrical charge (neutral) and a mass nearly equal to that of a proton. It is found in the nucleus of all atoms except protium (¹H).
Nucleons	The particles present in the nucleus of an atom, i.e., protons and neutrons.
Nucleus	The very small, dense, positively charged center of an atom, discovered by Rutherford. It contains protons and neutrons and constitutes nearly all the mass of the atom.
Octet	A stable state where the outermost shell of an atom is completely filled with eight electrons. Atoms tend to react to achieve this configuration.
Proton (p⁺)	A positively charged subatomic particle found in the nucleus of an atom. Its mass is taken as one unit.
Shells	Another term for the discrete orbits or energy levels (K, L, M, N,…) in which electrons revolve around the nucleus.
Valence Electrons	The electrons present in the outermost shell of an atom. These electrons determine the chemical properties and combining capacity of the element.
Valency	The combining capacity of an atom of an element. It is the number of electrons an atom gains, loses, or shares to achieve a stable electron configuration (an octet).