Chapter 8: Journey Inside the Atom
Verified step-by-step solutions for your class notes.
1. Choose the correct options and explain the reason for the correct and incorrect options in the context of Ernest Rutherford’s gold foil experiment:
- (i) The experiment clearly showed the existence of neutrons in the nucleus.
- (ii) The results disproved the plum pudding model and led to the idea of a nucleus at the centre of the atom.
- (iii) The large deflection of a few alpha particles indicated that most of the mass of the atom and positive charge are packed into a tiny centre.
- (iv) The way alpha particles were deflected showed that electrons move around the nucleus.
- (i) Incorrect: The experiment only detected a dense positive core (nucleus); neutrons were discovered much later by James Chadwick in 1932.
- (ii) Correct: Bouncing back of alpha particles showed positive charge isn't evenly spread, completely disproving Thomson’s model and establishing the central nucleus.
- (iii) Correct: Large-angle deflections and direct rebounds proved that almost all mass and positive charge are concentrated in a miniscule space.
- (iv) Incorrect: The deflection of positive alpha particles was caused by the positive nucleus repelling them; it did not map or show electronic motion directly.
2. Which of the following statements are correct or incorrect according to the Bohr’s atomic model? Give a reason for each statement.
- (i) Electrons lose energy while moving in fixed orbits and slowly fall into the nucleus.
- (ii) Electrons can exist anywhere around the nucleus with no fixed energy.
- (iii) Electrons revolve around the nucleus in orbits of fixed energy without losing energy.
- (iv) Electrons can be found between energy levels as they move around the nucleus.
- (i) Incorrect: Bohr's model states that while moving in a fixed allowed shell, an electron is in a stationary state and does not radiate or lose energy.
- (ii) Incorrect: Electrons cannot exist anywhere; they are strictly restricted to follow fixed specific circular paths called orbits or shells.
- (iii) Correct: This is a core postulate of Bohr’s model; stationary orbits keep the electron energy constant, ensuring atomic stability.
- (iv) Incorrect: Electrons can only revolve within the defined allowed shells and are strictly forbidden from existing in the spaces between them.
3. The composition of the nuclei of three atomic species X, Y, and Z are given as follows:
| X | Y | Z | |
|---|---|---|---|
| Number of protons | 18 | 17 | 17 |
| Number of neutrons | 19 | 18 | 20 |
Explain the relation between the following: (i) Y and Z, (ii) Z and X.
- (i) Relation between Y and Z (Isotopes): Both Y and Z have the same number of protons (17) but a different number of neutrons (Y has 18, Z has 20). Therefore, they are isotopes of the same element.
- (ii) Relation between Z and X (Isobars): Z and X have different atomic numbers (protons) but share the same mass number of 37 (Z: 17 + 20 = 37; X: 18 + 19 = 37). Therefore, they are isobars.
4. What conclusion did Rutherford draw about the position and characteristics of the atom’s positively charged part based on the few alpha particles that bounced back or were deflected at large angles in the gold foil experiment?
- Nucleus Characteristics: Rutherford concluded that the positive charge and nearly all the mass of an atom are not spread out uniformly, but instead concentrated in an extremely dense, compact central core called the nucleus. The rest of the atom is almost completely empty space.
5. Explain and arrange the following statements in the correct chronological order to show how atomic models have evolved over time.
- (i) Bohr’s model proposed that electrons move in fixed orbits around the nucleus, each with a definite energy.
- (ii) Thomson’s model depicted the atom as a ‘plum pudding’ with electrons embedded in a sphere of positive charge.
- (iii) Rutherford’s model proposed that atoms have a dense central nucleus.
- (iv) Dalton’s model described atoms as indivisible particles.
- Chronological Order: The correct scientific evolution timeline is (iv) → (ii) → (iii) → (i). Dalton started with indivisible particles (1808), Thomson added subatomic embedded charges (1897), Rutherford discovered the dense inner nucleus (1911), and Bohr established stable energy levels (1913).
6. Electrons move around the nucleus in orbits. Why do they not fly away from the atom? Explain what keeps them attracted to the nucleus.
- Electrostatic Attraction: Electrons carry a negative electrical charge, while the central nucleus contains protons carrying a positive electrical charge. The strong electrostatic force of attraction between these opposite charges keeps the electrons bound to the atom and prevents them from flying away.
7. Assertion (A): The discovery of subatomic particles helped in understanding the atomic structure.
Reason (R): The number of electrons is equal to the number of protons in an atom.
Choose the correct option: (i) Both A and R are true, and R is the correct explanation of A. (ii) Both A and R are true, but R is not the correct explanation of A. (iii) A is true, but R is false. (iv) A is false, but R is true.
- Correct Option: (ii) Both A and R are true, but R is not the correct explanation of A. The discovery of subatomic components (protons, electrons, neutrons) mapped out the structural framework of an atom, but the equality of electrons and protons simply explains why an atom maintains an overall neutral electric charge.
8. Magnesium is essential for many biological processes, including muscle contraction. For an atom of magnesium with a mass number of 24 and atomic number 12, determine the number of (i) protons, (ii) neutrons, (iii) electrons, and also illustrate the arrangement of electrons in a magnesium atom.
- (i) Protons: Equal to the atomic number, which gives 12 protons.
- (ii) Neutrons: Mass number minus atomic number (24 − 12), which gives 12 neutrons.
- (iii) Electrons: Equal to the proton count in a neutral atom, giving 12 electrons.
- Electronic Arrangement: Following Bohr-Bury rules, the 12 electrons fill the energy levels shell-by-shell to give a configuration of K=2, L=8, M=2.
9. Find the following information for the elements shown in Fig. 8.17: (i) Name of the element, (ii) Symbol, (iii) Total number of electrons, (iv) Number of valence electrons, (v) Valency of the element, (vi) Number of protons, (vii) Atomic number.
- Atom (a) [Lithium]: (i) Lithium, (ii) Li, (iii) 3, (iv) 1, (v) 1, (vi) 3, (vii) Z = 3.
- Atom (b) [Carbon]: (i) Carbon, (ii) C, (iii) 6, (iv) 4, (v) 4, (vi) 6, (vii) Z = 6.
- Atom (c) [Oxygen]: (i) Oxygen, (ii) O, (iii) 8, (iv) 6, (v) 8 − 6 = 2, (vi) 8, (vii) Z = 8.
- Atom (d) [Neon]: (i) Neon, (ii) Ne, (iii) 10, (iv) 8, (v) 0 (stable octet), (vi) 10, (vii) Z = 10.
10. Both Rutherford’s and Bohr’s models have electrons orbiting the nucleus. Why did Rutherford’s model fail to explain atomic stability, while Bohr’s model succeeded?
- Atomic Stability: In Rutherford's planetary model, accelerating circular electrons must continuously radiate electromagnetic energy, causing them to spiral inward and collapse into the nucleus. Bohr bypassed this limitation by postulating discrete stationary states (orbits) where mechanical rules keep electron energy fixed and radiationless during orbital motion.
11. An atom 70X has 31 electrons. How many neutrons are there in its nucleus?
- Neutron Count: In a neutral atom, electrons equal protons, so protons = 31. Using Mass Number = Protons + Neutrons (70 = 31 + n), the total number of neutrons is 70 − 31 = 39 neutrons.
12. An atom has 79 protons and a mass number of 197. Calculate (i) the number of neutrons, and (ii) the number of electrons.
- (i) Number of Neutrons: Mass number minus protons (197 − 79) gives 118 neutrons.
- (ii) Number of Electrons: Equal to the number of protons in a neutral atomic state, giving 79 electrons.
13. Complete the Table 8.5:
| Atomic No. | Mass No. | Neutrons | Protons | Electrons | Element Name |
|---|---|---|---|---|---|
| 5 | – | 6 | – | – | – |
| – | 14 | – | – | 7 | Nitrogen |
| – | 24 | – | 12 | – | – |
| 15 | – | 16 | – | – | – |
| – | 1 | 0 | – | – | – |
- Row 1 [Boron]: Mass No. = 5 + 6 = 11; Protons = 5; Electrons = 5; Name = Boron.
- Row 2 [Nitrogen]: Atomic No. = 7; Neutrons = 14 − 7 = 7; Protons = 7.
- Row 3 [Magnesium]: Atomic No. = 12; Neutrons = 24 − 12 = 12; Electrons = 12; Name = Magnesium.
- Row 4 [Phosphorus]: Mass No. = 15 + 16 = 31; Protons = 15; Electrons = 15; Name = Phosphorus.
- Row 5 [Hydrogen]: Atomic No. = 1 − 0 = 1; Protons = 1; Electrons = 1; Name = Hydrogen.
14. Element X has a mass number of 35 and contains 18 neutrons. Answer the following:
- (i) How many electrons and protons does element X have?
- (ii) What is its atomic number?
- (iii) Identify the element X.
- (iv) Write its electronic configuration.
- (v) How many valence electrons does it have?
- (vi) What will be the mass number if two neutrons are added to its nucleus?
- (vii) What will be the relation of X with the new atom?
- (i) Electrons & Protons: Protons = 35 − 18 = 17 protons; Electrons = 17 electrons.
- (ii) Atomic Number: Equal to the proton count, giving Z = 17.
- (iii) Identify Element X: Atomic number 17 corresponds to Chlorine (Cl).
- (iv) Electronic Configuration: Distributing 17 electrons yields 2, 8, 7.
- (v) Valence Electrons: Outer shell contains 7 valence electrons.
- (vi) New Mass Number: Initial mass 35 + 2 added neutrons = 37.
- (vii) Relation: They share the same atomic number (17) but have different mass numbers (35 and 37), making them isotopes.
15. In an atom, there are 12 protons and 12 neutrons in the nucleus. Now, imagine that all the electrons are replaced with some hypothetical particles that have the same charge as electrons but are 500 times heavier. What effect will this replacement have on the atom’s: (i) Atomic number, (ii) Atomic mass, (iii) Mass number, (iv) Overall charge?
- (i) Atomic Number: No change; it is strictly defined by the 12 protons inside the nucleus.
- (ii) Atomic Mass: Increases significantly; the mass of 12 electrons becomes 500 times heavier, making a measurable subatomic mass contribution.
- (iii) Mass Number: No change; mass number counts only the total integer tally of nuclear nucleons (12p + 12n = 24).
- (iv) Overall Charge: No change; the electrical charge of the replacement particle remains identical to a standard electron (−1).