NCERT Class 9 Science Chapter 4: Structure of the Atom - A Detailed Guide

Introduction to Structure of the Atom

Welcome, students! In our previous chapter, 'Atoms and Molecules', we learned about Dalton's Atomic Theory, which proposed that the atom is the smallest indivisible particle of matter. For a long time, this was the accepted view. However, the world of science is ever-evolving. By the end of the 19th century, scientists began to conduct experiments that challenged this very idea. The discovery of subatomic particles revealed that the atom itself has an internal structure. This raised new, exciting questions: If atoms are made of smaller particles, how are these particles arranged? What holds them together? Why are atoms of one element different from atoms of another?

Chapter 4 of your NCERT Class 9 Science textbook, 'Structure of the Atom', takes us on a fascinating journey to answer these very questions. We will explore the groundbreaking experiments and brilliant models proposed by scientists like J.J. Thomson, Ernest Rutherford, and Niels Bohr. This chapter is fundamental to understanding chemistry, as the arrangement of particles within an atom dictates its properties and how it interacts with other atoms. Let's delve deep into the heart of matter and uncover the secrets of the atom's structure.

Charged Particles in Matter: The Discovery of Subatomic Particles

The first clue that atoms were not indivisible came from experiments with static electricity and conduction of electricity through gases. These experiments suggested that atoms contained charged particles. The two primary particles discovered were the electron and the proton.

The Electron (e⁻)

In 1897, the English physicist J.J. Thomson was studying the properties of cathode rays. He conducted experiments using a cathode ray tube, which is a glass tube with most of the air pumped out. When a high voltage was applied across two electrodes inside the tube, a stream of particles was observed moving from the negative electrode (cathode) to the positive electrode (anode). These were called cathode rays.

Thomson discovered the following about these rays:

• They travel in straight lines.
• They consist of negatively charged particles. He showed this by applying an electric field; the rays bent towards the positive plate.
• The properties of these particles were the same, regardless of the gas inside the tube or the metal used for the cathode.

This led Thomson to a monumental conclusion: these negatively charged particles, which he called 'corpuscles' (later named 'electrons'), must be a fundamental constituent of all atoms. He determined that electrons were incredibly small, with a mass about 1/2000th that of a hydrogen atom, and carried a negative charge. For his work, he was awarded the Nobel Prize in Physics in 1906.

The Proton (p⁺)

Long before the electron was identified, in 1886, E. Goldstein conducted experiments with a modified cathode ray tube. He used a perforated cathode (a cathode with holes in it). He observed a new type of ray travelling in the opposite direction to the cathode rays, passing through the holes in the cathode. These rays were called 'canal rays'.

Properties of canal rays were found to be:

• They consisted of positively charged particles.
• The mass and charge of these particles depended on the gas taken in the tube.
• The lightest and simplest positive ion was obtained from hydrogen gas, and this was named the proton.

A proton was found to have a charge equal in magnitude but opposite in sign to that of an electron. Its mass was approximately 2000 times that of an electron. Since an atom is electrically neutral, it was concluded that it must contain an equal number of positively charged protons and negatively charged electrons.

The Structure of an Atom: Early Models

With the discovery of the electron and the proton, the question was no longer *if* an atom had parts, but *how* these parts were arranged. This led to the development of several atomic models, each attempting to explain the atom's structure and neutrality.

Thomson’s Model of an Atom (1904)

J.J. Thomson, the discoverer of the electron, was the first to propose a detailed model of the atom. His model is often referred to as the 'Plum Pudding Model' or the 'Watermelon Model'.

According to Thomson's model:

• An atom consists of a sphere of positive charge.
• The negatively charged electrons are embedded within this positive sphere, much like plums in a pudding or seeds in a watermelon.
• The total negative charge of the electrons is equal to the total positive charge of the sphere, making the atom electrically neutral as a whole.

While Thomson's model successfully explained the electrical neutrality of atoms, it was a purely hypothetical model based on the available data. It could not explain the results of later experiments, which led to its rejection.

Rutherford’s Model of an Atom (1911)

Ernest Rutherford, a student of J.J. Thomson, designed a brilliant experiment that would completely change the way we view the atom. This is famously known as the alpha-particle scattering experiment or the gold foil experiment.

The Alpha-Particle Scattering Experiment

Setup:

• Source of Alpha Particles: Rutherford used a radioactive source (like radium) that emitted fast-moving, positively charged alpha (α) particles. An alpha particle is essentially the nucleus of a helium atom (He²⁺), containing two protons and two neutrons.
• Target: A very thin sheet of gold foil was chosen. Gold is highly malleable, and the foil was only about 1000 atoms thick, ensuring that the alpha particles would likely interact with only one atom at a time.
• Detector: A circular screen coated with zinc sulphide was placed around the foil. When an alpha particle struck this screen, it would produce a tiny flash of light (scintillation), allowing the scientists to observe its path.

Observations of the Experiment

The results were completely unexpected based on Thomson's model. If the atom were a uniform sphere of positive charge, the alpha particles should have passed through with only minor deflections. Instead, Rutherford observed:

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. Very few alpha particles (about 1 in 12,000) appeared to bounce back, being deflected by nearly 180 degrees.

Rutherford famously remarked, "It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you."

Conclusions from the Experiment

Based on these startling observations, Rutherford drew the following conclusions about the structure of an atom:

• Since most particles passed through undeflected, he concluded that most of the space inside an atom is empty.
• Since some positively charged alpha particles were deflected, they must have been repelled by a strong positive charge within the atom. This led him to conclude that the entire positive charge of the atom is concentrated in a very small, dense region at the center. He called this region the nucleus.
• Since very few particles bounced back, the nucleus must be incredibly small and dense, occupying a tiny fraction of the atom's total volume. The radius of the nucleus was calculated to be about 10⁵ times smaller than the radius of the atom.

Rutherford's Nuclear Model of the Atom

Based on these conclusions, Rutherford proposed his nuclear model:

1. There is a positively charged center in an atom called the nucleus. Nearly all the mass of an atom resides in the nucleus.
2. The electrons revolve around the nucleus in well-defined circular paths or orbits.
3. The size of the nucleus is very small compared to the size of the atom.

Drawbacks of Rutherford's Model

Despite its successes, Rutherford's model had a major flaw. According to Maxwell's classical theory of electromagnetism, any charged particle moving in a circular path (i.e., undergoing acceleration) must radiate energy continuously. Therefore, an electron revolving around the nucleus should continuously lose energy and follow a spiral path, eventually falling into the nucleus. If this were true, atoms would be highly unstable. However, we know that atoms are quite stable. Rutherford's model could not explain the stability of the atom.

Bohr’s Model of an Atom (1913)

To overcome the objections raised against Rutherford's model, a Danish physicist named Niels Bohr proposed a new model of the atom in 1913. He incorporated ideas from the emerging quantum theory to explain the stability of atoms.

Postulates of Bohr's Model

Bohr's model was based on the following key postulates:

1. Electrons revolve around the nucleus only in certain special orbits known as discrete orbits or energy levels.
2. While revolving in these discrete orbits, the electrons do not radiate energy. This explains the stability of the atom.
3. These orbits or shells are associated with a fixed amount of energy. They are represented by the letters K, L, M, N... or the numbers n = 1, 2, 3, 4..., starting from the shell closest to the nucleus.
4. An electron can jump from a lower energy level to a higher one by absorbing a specific amount of energy, or from a higher level to a lower one by emitting energy.

Imagine these energy levels like rungs on a ladder. An electron can be on one rung or another, but not in between. It can move up or down the ladder, but it must land on a specific rung.

Discovery of the Neutron

The models so far accounted for protons and electrons. However, there was a discrepancy. The mass of a helium atom, for example, was known to be four times that of a hydrogen atom. But it only had two protons. This suggested the presence of another particle in the nucleus that had mass but no charge.

In 1932, James Chadwick, a colleague of Rutherford, discovered this missing particle. He bombarded a thin sheet of beryllium with alpha particles and observed the emission of highly penetrating radiation consisting of neutral particles. He named these particles neutrons (n⁰).

• A neutron has no charge.
• Its mass is almost equal to that of a proton.
• Neutrons are present in the nucleus of all atoms, except for ordinary hydrogen (Protium), which has only one proton and no neutrons.

The total mass of an atom is therefore the sum of the masses of protons and neutrons present in its nucleus (the mass of electrons is negligible).

How are Electrons Distributed in Different Orbits (Shells)?

The arrangement of electrons in the various energy levels or shells is known as the atom's electronic configuration. This distribution follows a set of rules proposed by Bohr and Bury.

Rules of the Bohr-Bury Scheme

1. The maximum number of electrons that can be accommodated in any given shell is determined by the formula 2n², where 'n' is the orbit number or energy level index (n=1 for K shell, n=2 for L shell, etc.).
   • K shell (n=1): Max electrons = 2(1)² = 2
   • L shell (n=2): Max electrons = 2(2)² = 8
   • M shell (n=3): Max electrons = 2(3)² = 18
   • N shell (n=4): Max electrons = 2(4)² = 32
2. The outermost shell of an atom cannot accommodate more than 8 electrons, even if it has the capacity to hold more according to the 2n² rule. This is often called the octet rule.
3. Electrons do not enter a new shell until the inner shells are completely filled. Shells are filled in a stepwise manner.

Electronic Configuration of First 18 Elements

Let's see how these rules apply with a table showing the electronic configuration of the first 18 elements.

Valency

The chemical properties of an element are determined by the number of electrons in its outermost shell. These are called valence electrons. Atoms react with each other to achieve a stable electronic configuration, which usually means having a completely filled outermost shell (typically 8 electrons, the octet rule).

Valency is defined as the combining capacity of an atom of an element. It is the number of electrons an atom loses, gains, or shares to achieve a stable configuration.

How to determine valency:

• If an atom has 1, 2, 3, or 4 valence electrons, it is easier for it to lose these electrons. Its valency is equal to the number of valence electrons. For example, Sodium (Na) has 1 valence electron, so its valency is 1. Magnesium (Mg) has 2, so its valency is 2.
• If an atom has 5, 6, or 7 valence electrons, it is easier for it to gain electrons to complete its octet. Its valency is calculated as (8 - number of valence electrons). For example, Chlorine (Cl) has 7 valence electrons, so its valency is 8 - 7 = 1. Oxygen (O) has 6, so its valency is 8 - 6 = 2.
• If an atom has a completely filled outermost shell (like Helium with 2 electrons or Neon and Argon with 8 electrons), it is chemically inert and does not need to lose, gain, or share electrons. Its valency is zero. These are the noble gases.

Atomic Number and Mass Number

These two numbers are crucial for identifying an atom and understanding its composition.

Atomic Number (Z)

The atomic number of an element is defined as the total number of protons present in the nucleus of its atom. It is denoted by the letter 'Z'.

• The atomic number is the unique identity of an element. All atoms of a particular element have the same atomic number. For example, all carbon atoms have Z = 6.
• In a neutral atom, the number of electrons is equal to the number of protons. Therefore, Z also indicates the number of electrons in a neutral atom.

Mass Number (A)

The mass number of an atom is defined as the sum of the total number of protons and neutrons present in the nucleus. It is denoted by the letter 'A'. The protons and neutrons are collectively known as nucleons.

Mass Number (A) = Number of protons (Z) + Number of neutrons (n)

For example, an atom of carbon has 6 protons and 6 neutrons. So, its mass number is A = 6 + 6 = 12. The standard notation for an atom is written as: ^A_ZX, where X is the symbol of the element. For carbon, this would be ¹²₆C.

Isotopes and Isobars

Dalton's theory stated that all atoms of an element are identical in mass and properties. However, this was later found to be incorrect. We now know about the existence of isotopes and isobars.

Isotopes

Isotopes are defined as atoms of the same element that have the same atomic number but different mass numbers.

This means isotopes of an element have the same number of protons (and electrons) but a different number of neutrons in their nuclei. Since they have the same number of electrons, their chemical properties are virtually identical. However, their physical properties, such as mass and density, are different.

Examples of Isotopes:

• Hydrogen: Has three isotopes - Protium (¹₁H, 1 proton, 0 neutrons), Deuterium (²₁H, 1 proton, 1 neutron), and Tritium (³₁H, 1 proton, 2 neutrons).
• Carbon: Has two common isotopes - Carbon-12 (¹²₆C, 6 protons, 6 neutrons) and Carbon-14 (¹⁴₆C, 6 protons, 8 neutrons).
• Chlorine: Exists as two isotopes in nature - Chlorine-35 (³⁵₁₇Cl) and Chlorine-37 (³⁷₁₇Cl) in a ratio of 3:1. This is why the atomic mass of chlorine is a fractional value (35.5 u), as it's the weighted average of the masses of its naturally occurring isotopes.

Applications of Isotopes

Isotopes have many important applications in various fields:

• An isotope of uranium (U-235) is used as fuel in nuclear reactors for generating electricity.
• An isotope of cobalt (Co-60) is used in the treatment of cancer.
• An isotope of iodine (I-131) is used in the treatment of goitre.
• Carbon-14 is used in carbon dating to determine the age of fossils and archaeological artifacts.

Isobars

Isobars are defined as atoms of different elements that have different atomic numbers but the same mass number.

This means isobars have different numbers of protons (they are different elements) but the same total number of nucleons (protons + neutrons). Since their electron configurations are different, their chemical properties are completely different.

Example of Isobars:

• Calcium (Ca) and Argon (Ar): Both have a mass number of 40. Argon has 18 protons and 22 neutrons (¹⁸⁺²²=40), while Calcium has 20 protons and 20 neutrons (²⁰⁺²⁰=40). So, ⁴⁰₁₈Ar and ⁴⁰₂₀Ca are isobars.

Important Questions and Answers

Question 1: Compare the properties of electrons, protons, and neutrons.

Answer:

Question 2: What are the limitations of J.J. Thomson’s model of the atom?

Answer: The main limitations of J.J. Thomson’s 'Plum Pudding' model were:

• It failed to explain the results of the alpha-particle scattering experiment conducted by Ernest Rutherford. The model could not explain why most alpha particles passed straight through the foil while a few were deflected at large angles.
• It did not provide any experimental evidence and was purely speculative.
• It could not explain the stability of the atom or the arrangement of electrons and the positively charged sphere.
• There was no mention of a nucleus in his model.

Question 3: If K and L shells of an atom are full, then what would be the total number of electrons in the atom?

Answer: The distribution of electrons in shells follows the 2n² rule, where 'n' is the shell number.

• For the K shell (n=1), the maximum number of electrons is 2(1)² = 2.
• For the L shell (n=2), the maximum number of electrons is 2(2)² = 8.

If both the K and L shells are full, the total number of electrons in the atom would be the sum of electrons in both shells: 2 (in K shell) + 8 (in L shell) = 10 electrons. The element with 10 electrons is Neon (Ne).

Question 4: Explain with examples (i) Atomic number, (ii) Mass number, (iii) Isotopes and (iv) Isobars. Give any two uses of isotopes.

Answer:

(i) Atomic Number (Z): The atomic number is the number of protons in the nucleus of an atom. It uniquely identifies an element. For example, the atomic number of Sodium (Na) is 11, which means every sodium atom has 11 protons.

(ii) Mass Number (A): The mass number is the total number of protons and neutrons in the nucleus of an atom. For example, an atom of Aluminium (Al) has 13 protons and 14 neutrons, so its mass number is 13 + 14 = 27.

(iii) Isotopes: These are atoms of the same element having the same atomic number but different mass numbers. This means they have a different number of neutrons. For example, Carbon has isotopes ¹²C (6 protons, 6 neutrons) and ¹⁴C (6 protons, 8 neutrons).

(iv) Isobars: These are atoms of different elements having different atomic numbers but the same mass number. For example, Argon (⁴⁰₁₈Ar) and Calcium (⁴⁰₂₀Ca) are isobars. Both have a mass number of 40 but different atomic numbers (18 and 20).

Two uses of isotopes are:

1. An isotope of Cobalt (Co-60) is used in radiotherapy for cancer treatment.
2. An isotope of Uranium (U-235) is used as a fuel in nuclear reactors to generate power.

Chapter Summary

Here are the key takeaways from our exploration of the 'Structure of the Atom':

• Atoms are not indivisible; they are made up of subatomic particles: electrons (negatively charged), protons (positively charged), and neutrons (neutral).
• J.J. Thomson's model proposed a positively charged sphere with electrons embedded in it, but it failed to explain experimental results.
• Rutherford's alpha-particle scattering experiment led to the discovery of the atomic nucleus – a small, dense, positively charged center where most of the atom's mass is concentrated.
• Rutherford's model described electrons orbiting the nucleus but couldn't explain the atom's stability.
• Niels Bohr's model refined this by proposing that electrons move in discrete energy levels or shells (K, L, M, N) without radiating energy.
• The maximum number of electrons in a shell is given by 2n², and the outermost shell can hold a maximum of 8 electrons.
• Valency is the combining capacity of an atom, determined by the number of valence electrons.
• Atomic Number (Z) is the number of protons and defines the element.
• Mass Number (A) is the sum of protons and neutrons.
• Isotopes are atoms of the same element with different numbers of neutrons (e.g., ¹²C and ¹⁴C).
• Isobars are atoms of different elements with the same mass number (e.g., ⁴⁰Ar and ⁴⁰Ca).