Structure of Atom Class 9 Notes Advanced Science Chapter 6

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Structure of Atom Notes Class 9 Advanced Science

Class 9 Structure of Atom Notes

Discovery of Subatomic Particles

In 1803, Dalton proposed that atoms are the smallest indivisible particles of matter. However, this idea could not explain the results of several experiments. For example, it was observed that substances like glass or ebonite, when rubbed with silk or fur, acquire an electric charge. This and many other experiments on electrical discharge through gases showed that atoms are not indivisible. They are made up of smaller particles called subatomic particles.

J. J. Thomson, in 1897, discovered the electron as a constituent of the atom and confirmed that the atom is not the smallest particle of matter. He proposed the so-called plum-pudding model of the atom, in which electrons are embedded in a sphere of positive charge. This model was later shown to be incorrect by Rutherford, who demonstrated that it could not explain the results of the gold foil experiment. Rutherford then proposed a model in which electrons revolve around a small, positively charged nucleus. However, this model could not explain the stability of the atom. Thereafter, another model was proposed by Niels Bohr.

This section discusses the discovery of subatomic particles, the electron, proton and neutron, which helped in understanding the structure of the atom.

Basic principle: like charges repel each other, while unlike charges attract each other.

Discovery of the Electron
In the late nineteenth century, many scientists, including Michael Faraday, William Crookes and others, studied electrical discharge in partially evacuated tubes known as cathode ray discharge tubes.

J. J. Thomson conducted the cathode ray experiment to study the nature of matter, which led to the discovery of the electron.
Apparatus Used

• A glass discharge tube fitted with two metal electrodes:
• Cathode (negative electrode)
• Anode (positive electrode)
• The tube is connected to a high-voltage source
• Air inside the tube is removed to create a low-pressure (near-vacuum) condition.
  

Procedure:

1. A high voltage is applied across the electrodes.
2. Rays originate from the cathode and travel towards the anode. (Cathode Rays)
3. These rays produce a greenish glow on the glass wall of the tube.
4. The presence of these rays can be detected by allowing them to pass through a hole in the anode and strike a screen coated with a special material placed behind it. A bright spot is observed on the screen, indicating that the rays travel in straight lines.
5. The rays are then tested under different conditions to determine their nature:
   • By placing objects in their path, a shadow is formed on the glass screen. It shows that cathode rays travel in straight lines.
   • By applying an electric field, cathode rays are deflected towards the positive plate. It shows that cathode rays are negatively charged particles called electrons.
   • By applying a magnetic field, cathode rays are deflected from their path. It shows that cathode rays consist of charged particles (moving charges are affected by magnetic fields).
   • By using different gases (such as hydrogen, nitrogen, neon, etc.) and different electrode materials, it was found that the properties of cathode rays remained unchanged. This showed that electrons are present in all atoms.

The main characteristics of cathode rays are as follows:

• They originate from the cathode and move towards the anode.
• They are not visible themselves but produce a bright spot when they strike certain materials.
• They travel in straight lines in the absence of external fields.
• They are deflected by electric and magnetic fields in such a way that indicates them to be negatively charged.
• Their properties are independent of the nature of the gas or the electrode material.
  Thus, electrons are a fundamental constituent of all atoms.

Discovery of Protons
After the discovery of the electron, it was realised that since electrons are negatively charged, atoms must also contain a positive charge to maintain electrical neutrality.

Eugen Goldstein, in 1886, performed experiments using a discharge tube similar to that used for cathode rays, but with a cathode having holes in it (perforated cathode). When high voltage was applied, a faint glow was observed behind the cathode. The rays responsible for this glow passed through the holes (or canals) in the cathode and were therefore called canal rays.

Further studies showed that these rays were deflected towards the negatively charged plate in electric and magnetic fields, indicating that they consist of positively charged particles.

However, it is important to note that canal rays are not made up of a single type of particle. They consist of positively charged ions of the gas present in the tube. Therefore, their properties depend on the nature of the gas used.

When hydrogen gas was used in the discharge tube, the positively charged particles obtained were the lightest known and were identified as hydrogen ions (H+). These particles were later recognised as protons. The proton was finally established as a fundamental particle by Rutherford in 1919.

The main characteristics of canal rays are the following:

• They are positively charged.
• Their behaviour in electric and magnetic fields is opposite to that of electrons.
• Their properties depend on the nature of the gas present.
• The lightest positive particle was obtained from hydrogen and is called the proton.

Discovery of Neutrons
Once electrons and protons were known, it appeared that the structure of the atom was understood. However, another problem arose when atomic masses were measured. The mass of atoms was found to be greater than the sum of the masses of their protons and electrons. For example, helium contains two protons, yet its mass is about four times that of hydrogen. This suggested the possibility of another particle contributing to the mass of the atom.

It was proposed that there must be a neutral particle present in the atom. This particle was discovered by James Chadwick in 1932.

Experiment Setup

• A thin sheet of beryllium metal was bombarded with alpha particles (from a radioactive source).
• A block of paraffin wax (rich in hydrogen atoms) was placed in front of the beryllium.
• A detector was used to observe emitted particles.

Observations

• A new type of radiation was produced from beryllium.
• This radiation was uncharged (not deflected by electric or magnetic fields).
• When it struck paraffin wax, it knocked out protons (hydrogen nuclei) at high speed.

Explanation

• As the radiation could knock out heavy protons, it must consist of particles with mass similar to protons.
• As it showed no charge, it could not be positive or negative.

Conclusion

• The radiation consisted of neutral particles, called neutrons.
• Neutrons have no charge, and their mass is nearly equal to that of protons.

Neutrons are present in the nuclei of almost all atoms. The most common isotope of hydrogen that does not contain a neutron is Protium (₁H¹), but its heavier isotopes, Deuterium (₁H²) and Tritium (₁H³), contain neutrons. Thus, the presence of neutrons explains the mass of atoms. For example, helium contains two protons and two neutrons, which accounts for its mass being approximately four times that of hydrogen. Chadwick was awarded the Nobel Prize in Physics in 1935 for the discovery of the subatomic particle neutron. From these discoveries, it became clear that atoms are composed of three subatomic particles:

1. Electrons (negative charge)
2. Protons (positive charge)
3. Neutrons (no charge)

These particles together determine the structure and properties of atoms.

Spectrum, Line Spectrum of Hydrogen, Limitations of Rutherford Model, Bohr’s Model

Spectrum
A spectrum is the band of colours or wavelengths of light obtained when light is dispersed (split), usually by a prism.

There are two types of spectrum:
1. Continuous spectrum: A continuous spectrum is one in which all wavelengths of radiation are so intermixed that there is no line of separation between two colours. For example – Sunlight, rainbow (a continuous spread of many colours from violet to red (VIBGYOR), etc.

2. Line spectrum: A line spectrum is a spectrum that consists of distinct, separate lines of specific wavelengths (or colours), with dark spaces in between, instead of a continuous range of colours.

For example, a sodium vapour lamp and a mercury vapour lamp.

A line spectrum means only specific wavelengths of radiation are emitted, not all wavelengths. It is important to know that a line spectrum is characteristic of the element that produces it. This fact is used to identify elements present in stars by studying their spectra.

Line Spectrum of Hydrogen
When radiation emitted by hydrogen gas at low pressure is passed through a discharge tube and examined with a spectroscope (which uses a prism), the resulting line spectrum is called the atomic spectrum of hydrogen.

The Spectral Lines for Atomic Hydrogen

Series	ni	nf
Lyman	1	2, 3,…
Balrner	2	3, 4,…
Paschen	3	4, 5,…
Brackett	4	5, 6,…
Pfund	5	6, 7,…

Rydberg gave a general expression that applies to all series in the hydrogen spectrum for expressing the energies of the distinct spectral lines.
This expression is called the Rydberg formula.
∆E = R_EZ²\(\left(\frac{1}{n_i^2}-\frac{1}{n_f^2}\right)\)
Where R_E = 1.097 × 10^7, m^-1 is the Rydberg constant expressed in terms of energy, Z is the atomic number n_i and n_f and m are the initial and final energy levels, respectively.

These empirical formulas worked very well but could not be explained until Bohr’s model came.

Limitation of Rutherford Model of Atom

1. Rutherford could not explain stability, as the electron continuously loses energy when it moves around the nucleus.
2. As the electron in the atom is allowed to have continuous energies, the emitted radiation is expected to give a continuous set of radiation. However, we observe a line spectrum. Therefore, we say that Rutherford’s model fails to explain the existence of the line spectrum of hydrogen.

Bohr’s model
In 1913, Niels Bohr, a student of Rutherford, proposed his model for an atom. He combined Rutherford’s nuclear model with the new quantum idea introduced by Max Planck.

1. Bohr explained the stability of an atom.
2. Electrons revolve around the nucleus in certain selected circular paths, called ‘orbits’, without emitting any energy.
3. The electrons revolve only in those orbits that have a fixed value of energy.
4. As long as the electron remains in a particular orbit, it neither loses nor gains energy. The radiation is emitted or absorbed only when an electron jumps from one allowed orbit to another.

When energy is supplied to an electron, it may jump instantaneously from a lower energy level (say K) to a higher energy level (say L, M, N, etc.) by absorbing one or more quanta of energy.

Achievement of Bohr’s Model- Bohr’s model was able to explain the line spectrum of hydrogen. When an electron jumps from an orbit of higher energy to that of a lower energy, it releases energy in the form of radiation. The amount of energy released depends on the difference in the energies of the two levels. As the orbits of only certain energies exist, only fixed-quantity energy differences are possible. Therefore, a well- defined line spectrum is obtained.

Limitations of Bohr’s model

1. It can only explain the line spectrum of hydrogen adequately. It fails to explain why single spectral lines are composed of several closely-spaced fine lines, as revealed by high- resolution spectroscopy.
2. It can only work for (at least) one-electron atoms and cannot explain multi-line spectra with each colour. It fails to explain spectra for multi-electron atoms.
3. The theory does not take into account the wave properties of electrons.
4. The theory does not account for the splitting of spectral lines in the presence of a magnetic field (Zeeman effect) or an electric field (Stark effect).