CBSE Class 9 Science: Structure Of Atom

From Indivisible Atoms to Subatomic Particles

For a long time, scientists believed that atoms were the fundamental, indivisible particles of matter, as proposed by John Dalton's atomic theory in 1808. However, experimental evidence towards the end of the 19th century began to challenge this idea. The discovery of cathode rays by William Crookes and their investigation by J.J. Thomson revealed the existence of negatively charged particles, much smaller than an atom, which he named electrons. This discovery proved that atoms are indeed divisible and contain internal structures. Shortly after, the discovery of positively charged particles, called protons, led to further modifications in the understanding of atomic structure. Ernest Rutherford's famous alpha-particle scattering experiment provided conclusive evidence for a central, dense nucleus within the atom, containing all the positive charge and most of the atom's mass. Later, James Chadwick discovered neutrons, neutral particles also residing in the nucleus. These discoveries completely revolutionized our perception of the atom, transforming it from a simple, solid sphere into a complex system of subatomic particles.

Key Subatomic Particles and Atomic Terminology

Electron (e⁻)

A negatively charged subatomic particle found orbiting the nucleus of an atom. It has negligible mass compared to protons and neutrons.

Proton (p⁺)

A positively charged subatomic particle found in the nucleus of an atom. It has a mass approximately equal to that of a neutron.

Neutron (n⁰)

A neutral (no charge) subatomic particle found in the nucleus of an atom. It has a mass approximately equal to that of a proton.

Atomic Number (Z)

The number of protons present in the nucleus of an atom. It uniquely identifies an element and is equal to the number of electrons in a neutral atom.

Mass Number (A)

The total number of protons and neutrons (nucleons) in the nucleus of an atom. It approximates the atomic mass of the element.

Valency

The combining capacity of an element, determined by the number of electrons an atom can gain, lose, or share to achieve a stable outer electron shell (octet or duplet).

Journey Through Atomic Models

1. Thomson's Model of an Atom (1904) — J.J. Thomson proposed that an atom resembles a 'plum pudding' or 'watermelon'. He suggested that the atom is a positively charged sphere, and the electrons are embedded in it like plums in a pudding or seeds in a watermelon. This model explained the electrical neutrality of an atom but failed to explain later experimental results, particularly Rutherford's alpha-particle scattering experiment.
2. Rutherford's Model of an Atom (1911) — Ernest Rutherford conducted the famous alpha-particle scattering experiment, where he bombarded a thin gold foil with fast-moving alpha-particles. Most particles passed straight through, some were deflected by small angles, and a very few bounced back. Based on these observations, he concluded that: 1. Most of the atom is empty space. 2. There is a tiny, dense, positively charged nucleus at the center, containing almost all the mass. 3. Electrons revolve around this nucleus in well-defined orbits. This model, often called the 'nuclear model', explained the existence of a nucleus but had a major drawback: according to classical physics, orbiting electrons would lose energy and eventually spiral into the nucleus, making the atom unstable. However, atoms are known to be stable.
3. Bohr's Model of an Atom (1913) — Niels Bohr addressed the limitations of Rutherford's model. He proposed that electrons revolve around the nucleus in specific, discrete orbits or shells, called 'energy levels' or 'stationary orbits'. Electrons do not radiate energy while revolving in these orbits. Energy is only absorbed or emitted when an electron jumps from one energy level to another. The orbits are designated as K, L, M, N shells (n=1, 2, 3, 4...). This model successfully explained the stability of atoms and the discrete line spectra observed for elements, forming the basis of our current understanding of electron shells.

Electronic Configuration and Valency

• Example 1: Hydrogen (H) Atomic Number (Z) = 1. Number of electrons = 1. Electronic Configuration: K-1 (1 electron in the first shell). Valency: 1 (It can share 1 electron to achieve stability).
• Example 2: Carbon (C) Atomic Number (Z) = 6. Number of electrons = 6. Electronic Configuration: K-2, L-4 (2 electrons in K shell, 4 in L shell). Valency: 4 (It needs 4 electrons to complete its L shell, usually by sharing).
• Example 3: Oxygen (O) Atomic Number (Z) = 8. Number of electrons = 8. Electronic Configuration: K-2, L-6 (2 electrons in K shell, 6 in L shell). Valency: 2 (It needs 2 electrons to complete its L shell, usually by gaining or sharing).
• Example 4: Sodium (Na) Atomic Number (Z) = 11. Number of electrons = 11. Electronic Configuration: K-2, L-8, M-1 (2 in K, 8 in L, 1 in M). Valency: 1 (It tends to lose its 1 outermost electron to achieve a stable octet in the L shell).
• Example 5: Chlorine (Cl) Atomic Number (Z) = 17. Number of electrons = 17. Electronic Configuration: K-2, L-8, M-7 (2 in K, 8 in L, 7 in M). Valency: 1 (It tends to gain 1 electron to achieve a stable octet in the M shell).

Mastering Isotopes and Isobars for Exams

Students often confuse isotopes and isobars. Remember: Isotopes are atoms of the same element (same atomic number, 'Z', meaning same number of protons) but have different mass numbers ('A', meaning different number of neutrons). Examples: Carbon-12 and Carbon-14. They have identical chemical properties. Conversely, Isobars are atoms of different elements (different atomic number, 'Z') that have the same mass number ('A', meaning the sum of protons and neutrons is the same). Examples: Argon-40 and Calcium-40. They have different chemical properties. A common mistake is to mix up the 'same Z, different A' and 'different Z, same A' definitions. Always clarify these two concepts to avoid losing marks. Also, understanding the rules for filling electrons in shells (Bohr-Bury rules, 2n² rule) is critical for valency calculations.

Practice Questions with Solutions

• Q: An atom has 7 protons, 7 electrons, and 8 neutrons. What are its atomic number and mass number? Also, identify the element. A: Step 1: Atomic Number (Z) is the number of protons. Here, Z = 7. Step 2: Mass Number (A) is the sum of protons and neutrons. Here, A = 7 + 8 = 15. Step 3: The element with atomic number 7 is Nitrogen (N). Final answer: Atomic Number = 7, Mass Number = 15. The element is Nitrogen.
• Q: Write the electronic configuration and valency for an element with atomic number 12. A: Step 1: The element has an atomic number of 12, so it has 12 electrons in a neutral atom. Step 2: Fill the electron shells according to Bohr-Bury rules: K shell: 2 electrons L shell: 8 electrons M shell: 2 remaining electrons Electronic configuration: K-2, L-8, M-2. Step 3: To achieve stability (an octet in its outermost shell), the atom will tend to lose the 2 electrons in its M shell. Final answer: Electronic configuration is 2, 8, 2. Valency is 2.
• Q: State two postulates of Bohr's model of an atom. A: Step 1: Recall the key features Bohr proposed to overcome Rutherford's limitations. Step 2: Postulate 1: Electrons revolve around the nucleus in specific, discrete orbits called 'energy shells' or 'stationary orbits'. Step 3: Postulate 2: While revolving in these discrete orbits, electrons do not radiate energy. Energy is only absorbed or emitted when an electron jumps from one energy level to another. Final answer: Two postulates are (1) electrons occupy discrete orbits without radiating energy, and (2) energy is only exchanged when electrons transition between these orbits.
• Q: What are isotopes? Give an example. A: Step 1: Define isotopes based on their atomic number and mass number. Step 2: Isotopes are atoms of the same element (i.e., they have the same atomic number or number of protons) but have different mass numbers (i.e., different numbers of neutrons). Step 3: Provide a common example. For instance, Carbon has two main isotopes: Carbon-12 (6 protons, 6 neutrons) and Carbon-14 (6 protons, 8 neutrons). Final answer: Isotopes are atoms of the same element with the same atomic number but different mass numbers. Example: Carbon-12 and Carbon-14.
• Q: If an atom loses an electron, what type of ion is formed and why? A: Step 1: Consider the initial state of a neutral atom (equal protons and electrons). Step 2: If an atom loses a negatively charged electron, it will have more positively charged protons than negatively charged electrons. Step 3: This imbalance results in a net positive charge on the atom. Final answer: A positively charged ion, called a cation, is formed because the atom now has more protons than electrons.

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