Introduction to the Topic

Welcome to NCERT Explained! In this post, we dive into Class IX Science, Chapter 4 - Structure of the Atom. Previously, we learned about atoms and molecules as the fundamental building blocks of matter. But what is an atom itself made of? Is it truly indivisible, as John Dalton once suggested? Through fascinating experiments and scientific discoveries, we realized that atoms have internal structures composed of even smaller particles called subatomic particles. Let us explore how scientists unraveled the mystery of the atom.

Key Concepts Explained

To understand how an atom is structured, we must look at the historical discoveries of subatomic particles and the different models proposed by scientists over time.

1. Charged Particles in Matter

The concept of divisibility in matter started gaining traction through static electricity experiments. When you rub two objects together, they acquire electric charge. This proved the presence of charged particles inside the atom. The two primary subatomic particles discovered initially are:

  • Protons: Positively charged particles discovered by E. Goldstein using a modified cathode ray tube.
  • Electrons: Negatively charged particles discovered by J.J. Thomson in 1897.
  • Later, James Chadwick discovered a neutral particle called the Neutron, which has no charge but has mass nearly equal to that of a proton.

2. Thomson's Model of the Atom

J.J. Thomson was the first to propose a model for the atomic structure. He envisioned the atom as a positively charged sphere in which electrons were embedded, much like seeds in a watermelon or plums in a Christmas pudding. According to Thomson's model:

  • An atom consists of a positively charged sphere, and the electrons are embedded in it.
  • The negative and positive charges are equal in magnitude, making the atom electrically neutral as a whole.

However, this model could not explain the results of later scattering experiments conducted by other scientists.

3. Rutherford's Model of the Atom

Ernest Rutherford wanted to know how electrons are arranged within an atom. He designed an experiment known as the Alpha ($\alpha$) particle scattering experiment. Fast-moving alpha particles (doubly-charged helium ions) were made to fall on a very thin gold foil.

  • Most of the $\alpha$-particles passed straight through the gold foil without any deflection.
  • A small fraction of the $\alpha$-particles was deflected by small angles.
  • Surprisingly, out of every 12,000 particles, one appeared to bounce back completely!

Based on these observations, Rutherford concluded that:

  • Most space inside the atom is empty because most $\alpha$-particles passed through.
  • A very small mass and positive charge are concentrated at the center of the atom, called the nucleus.
  • Electrons revolve around the nucleus in circular paths.

However, a major drawback of Rutherford's model was that any charged particle in accelerated motion would radiate energy, causing the revolving electrons to spiral into the nucleus, making the atom unstable. But atoms are stable!

4. Bohr's Model of the Atom

To overcome the objections raised against Rutherford's model, Neils Bohr in 1913 proposed a revised model:

  • Only certain special orbits known as discrete orbits of electrons are allowed inside the atom.
  • While revolving in these discrete orbits, the electrons do not radiate energy.
  • These orbits or shells are represented by the letters $K, L, M, N...$ or numbers $n = 1, 2, 3, 4...$.

5. Distribution of Electrons in Orbits (Bohr-Bury Scheme)

The distribution of electrons into different orbits of an atom is governed by rules known as the Bohr-Bury scheme:

  • The maximum number of electrons present in a shell is given by the formula $2n^2$, where 'n' is the orbit number or energy level index. The maximum capacity for $K$ shell ($n=1$) is $2(1)^2 = 2$, for $L$ shell ($n=2$) is $2(2)^2 = 8$, and so on.
  • The outermost shell can accommodate a maximum of 8 electrons.
  • Shells are filled in a step-by-step manner, meaning inner shells are filled before outer shells begin filling.

6. Valency

The electrons present in the outermost shell of an atom are known as valence electrons. Atoms try to achieve a stable configuration of 8 electrons in their outermost shell (an octet). The combining capacity of an atom, or its tendency to share, gain, or lose electrons to complete its octet, is called its valency.

7. Atomic Number and Mass Number

Every element is uniquely identified by its fundamental properties:

  • Atomic Number ($Z$): Defined as the total number of protons present in the nucleus of an atom. Since atoms are neutral, it is also equal to the number of electrons. $Z = \text{Number of Protons}$
  • Mass Number ($A$): Defined as the sum of the total number of protons and neutrons in the nucleus of an atom. These particles are collectively called nucleons. $A = \text{Number of Protons} + \text{Number of Neutrons}$

8. Isotopes and Isobars

Sometimes, atoms of the same element can have differences:

  • Isotopes: Atoms of the same element having the same atomic number but different mass numbers. Example: Hydrogen has three isotopes—Protium ($^1_1H$), Deuterium ($^2_1H$), and Tritium ($^3_1H$). Isotopes have identical chemical properties due to the same number of electrons, but different physical properties.
  • Isobars: Atoms of different elements with different atomic numbers but the same mass number. Example: Calcium (Atomic number 20, Mass number 40) and Argon (Atomic number 18, Mass number 40).

Summary & Key Takeaways

Here are the crucial points to remember from this chapter:

  • An atom is composed of negatively charged electrons, positively charged protons, and neutral neutrons. Protons and neutrons reside in the dense nucleus, while electrons revolve around it.
  • J.J. Thomson proposed the plum pudding model, which was later refined by Rutherford's nuclear model through his alpha particle scattering experiment.
  • Bohr improved upon Rutherford's model by introducing discrete stationary orbits where electrons do not radiate energy.
  • The maximum number of electrons in any shell is determined by the formula $2n^2$.
  • Atomic number ($Z$) equals the number of protons, and Mass number ($A$) is the sum of protons and neutrons.
  • Isotopes are atoms of the same element with different mass numbers, while isobars are atoms of different elements with the same mass number.