Structure Of Atom

Matter contains charged particles. The attractive force shown in static electricity and the conduction of electricity through substances gave evidence that atoms are not simple indivisible particles.

The study of atomic structure mainly includes:

• Electron
• Proton
• Neutron
• Thomson’s atomic model
• Rutherford’s atomic model
• Bohr’s atomic model
• Electronic configuration
• Atomic number and mass number
• Isotopes, isobars and isotones

Electrical Nature Of Particles In Matter

The electrical nature of matter became clear from experiments related to static electricity and discharge of electricity through gases.

These experiments showed that matter contains charged particles.

Discovery Of Electron

The discovery of the electron is also known as the discovery of cathode rays.

J.J. Thomson discovered cathode rays while studying the passage of electricity through gases at extremely low pressure.

He passed electricity at high voltage through a discharge tube containing gas at very low pressure.

During the experiment, a green fluorescence was seen at the other end of the discharge tube.

This fluorescence was produced by rays emitted from the cathode, or negative plate, towards the anode, or positive plate.

These rays were called cathode rays.

The conclusion was that cathode rays are streams of negatively charged particles.

These negatively charged particles are called electrons.

An electron is represented as:

• e
• e⁻

Discovery Of Proton

The discovery of the proton is linked with the discovery of canal rays or anode rays.

Since atoms contain negatively charged electrons, scientists understood that atoms must also contain positively charged particles because an atom is electrically neutral.

E. Goldstein modified the discharge tube and passed electric current through it.

He observed that positively charged rays were emitted from the anode.

These rays were called canal rays.

They were also called anode rays because they were emitted from the anode in gas discharge experiments using a perforated cathode.

When an electric field was applied, these rays were deflected towards the negatively charged plate.

The conclusion was that atoms contain positively charged particles along with electrons.

The lightest positively charged particle was named proton by Ernest Rutherford.

A proton is represented as:

• p
• p⁺

Discovery Of Neutron

Till 1930, it was believed that the mass of an atom was mainly due to protons because electrons have negligible mass.

In 1932, James Chadwick discovered another subatomic particle.

He observed that when beryllium was exposed to alpha particles, a different kind of particle was emitted.

These particles had nearly the same mass as protons but carried no electrical charge.

They were named neutrons.

Neutrons are present in the nucleus along with protons.

Neutrons are present in the nucleus of all atoms except hydrogen.

Protons and neutrons inside the nucleus are together called nucleons.

Subatomic Particles Of Atom

Particle	Charge	Position	Important Point
Electron	Negative	Outside nucleus	Negligible mass
Proton	Positive	Inside nucleus	Charge equal in magnitude but opposite to electron
Neutron	No charge	Inside nucleus	Mass nearly equal to proton

Structure Of Atom

After the discovery of electron, proton and neutron, scientists tried to explain how these particles are arranged inside an atom.

Three important atomic models are discussed in this chapter:

• Thomson’s Model
• Rutherford’s Model
• Bohr’s Model

Thomson’s Model Of Atom

J.J. Thomson was the first scientist to propose a model for the structure of an atom.

His model is popularly known as:

• Watermelon model
• Christmas pudding model
• Plum pudding model

Main Points Of Thomson’s Model

• An atom is a positively charged sphere.
• Electrons are embedded in this positively charged sphere.
• The negative charge of electrons and the positive charge of the sphere are equal in magnitude.
• Therefore, the atom is electrically neutral.
• The size of the atom was about 10⁻¹⁰ m.

Limitations Of Thomson’s Atomic Model

Thomson’s model had important limitations.

• It could not explain the distribution of electrons in an atom.
• If electrons were embedded in positive charge, opposite charges should cancel each other.
• It could not explain why different elements have different chemical properties.
• It could not explain the results of Rutherford’s alpha-particle scattering experiment.

Rutherford’s Model Of Atom

Ernest Rutherford and his co-workers made an important contribution to the understanding of atomic structure.

His experiment is known as the alpha-particle scattering experiment or gold foil experiment.

Rutherford is also known as the Father of Nuclear Physics.

Rutherford’s Alpha-Particle Scattering Experiment

In this experiment, Rutherford used a very thin sheet of gold foil, about 1000 atoms thick.

He bombarded it with a stream of alpha particles.

Alpha particles are positively charged helium ions, or helium nuclei.

They are represented as:

• He²⁺

Important features of alpha particles:

• They carry two units of positive charge.
• Their mass is four times that of hydrogen atom.
• Their mass is 4 u.
• They have considerable energy.

The alpha particles were emitted from radioactive elements such as radium.

A radioactive substance was placed in a lead block with a narrow slit.

This allowed only a narrow beam of alpha particles to pass through.

The beam was directed at the thin gold foil.

A movable screen coated with zinc sulphide was placed around the gold foil to detect deflection.

When alpha particles struck the zinc sulphide screen, flashes of light called scintillations were produced.

Observations Of Rutherford’s Experiment

Rutherford observed that:

• Most alpha particles passed straight through the gold foil without any deflection.
• Some alpha particles were deflected through small angles.
• Very few alpha particles, about 1 in 12,000, showed large deflection or came back in the same direction.

Conclusions Of Rutherford’s Experiment

From these observations, Rutherford concluded:

• Most of the space inside an atom is empty.
• The positive charge and most of the mass of the atom are concentrated in a very small region.
• This small central region was named the nucleus.
• Large deflection of alpha particles happened due to direct collision with the positively charged nucleus.
• The nucleus is about 10⁵ times smaller than the atom.

Important values:

• Radius of atom = 10⁻¹⁰ m
• Radius of nucleus = 10⁻¹⁵ m

Thus, the atom is mostly hollow, with a small nucleus at the centre.

Rutherford’s Atomic Model

According to Rutherford’s model:

• Atom has a positively charged nucleus.
• Almost all the mass of the atom is concentrated in the nucleus.
• The nucleus is very small compared to the size of the atom.
• The nucleus is surrounded by negatively charged electrons.
• Number of electrons = number of protons in a neutral atom.
• Positive charge on the nucleus differs for different atoms.
• Electrons revolve around the nucleus at very high speed.
• Electrons revolve like planets around the Sun.
• Therefore, electrons are also called planetary electrons.

Drawbacks Of Rutherford’s Model

Rutherford’s model could not explain the stability of atoms.

According to classical physics, any charged particle in motion should radiate energy.

Since electrons are charged particles and revolve around the nucleus, they should continuously lose energy.

If electrons lose energy, they should eventually fall into the nucleus.

This would make the atom unstable.

But atoms are stable.

Therefore, Rutherford’s model failed to explain atomic stability.

Bohr’s Model Of Atom

Niels Bohr modified Rutherford’s model and explained the stability of atoms.

Main Postulates Of Bohr’s Atomic Model

• Electrons possess a fixed amount of energy.
• This energy allows them to revolve around the nucleus.
• The nucleus is situated at the centre of the atom.
• Electrons revolve in fixed circular paths called orbits or shells.
• These orbits have fixed energy and are called energy levels.
• Electrons do not radiate energy while revolving in stationary orbits.
• Smaller orbits have lower energy.
• Energy increases as we move away from the nucleus.
• Orbits are numbered as 1, 2, 3, 4 and so on.
• They are also named as K, L, M, N and so on.
• When an electron moves to a higher energy level, it absorbs energy.
• When an electron moves to a lower energy level, it emits energy.

Distribution Of Electrons In Orbits

The maximum number of electrons present in a shell is given by the formula:

$2n^2$2n2

Here, n is the number of the shell.

Shell	Orbit Number	Maximum Electrons
K shell	n = 1	2 × 1² = 2
L shell	n = 2	2 × 2² = 8
M shell	n = 3	2 × 3² = 18
N shell	n = 4	2 × 4² = 32

Rules For Filling Electrons

• The maximum number of electrons in the outermost orbit is 8.
• Shells are filled in a stepwise manner.
• Electrons are not accommodated in a shell unless the inner shells are filled first.

This stepwise filling is linked with Aufbau’s Principle.

Octet Rule

The octet rule states that the maximum number of electrons in the outermost shell of a chemically stable and electrically neutral atom is 8.

Exception:

• Hydrogen
• Helium

Hydrogen and helium can have only 2 electrons in their outermost shell.

This is called duplet.

Valency

Valency is the combining capacity of an element.

It is also defined as the number of electrons lost, gained or shared by an atom during chemical combination.

Valence Shell And Valence Electrons

The outermost shell or orbit of an atom is called the valence shell.

The electrons present in the outermost shell are called valence electrons.

The number of valence electrons varies from 1 to 8.

Valence electrons determine the valency of an element.

Atomic Number

The atomic number is the number of protons present in the nucleus of an atom.

It is represented by Z.

Formula:

$Z = \text{Number of protons}$Z=Number of protons

For a neutral atom:

Number of protons = Number of electrons

So:

$Z = p^{+} = e^{-}$Z=p+=e−

Atomic Mass Number

The atomic mass number is the sum of the number of protons and neutrons present in the nucleus.

It is represented by A.

Formula:

$A = p + n$A=p+n

The atomic number, mass number and symbol of an element are written together as:

${}^{A}_{Z}X$ZA​X

Here:

• A = Atomic mass number
• Z = Atomic number
• X = Symbol of element

Charged Species

For charged species, atomic number remains equal to the number of protons, but it is not equal to the number of electrons.

Positive Ion Or Cation

A positively charged ion forms when an atom loses electrons.

For a cation:

• Number of protons > Number of electrons

Example:

Sodium has atomic number 11.

Neutral sodium atom:

• Protons = 11
• Electrons = 11

Sodium ion Na⁺:

• Protons = 11
• Electrons = 10

Negative Ion Or Anion

A negatively charged ion forms when an atom gains electrons.

For an anion:

• Number of protons < Number of electrons

Example:

Chlorine has atomic number 17.

Neutral chlorine atom:

• Protons = 17
• Electrons = 17

Chloride ion Cl⁻:

• Protons = 17
• Electrons = 18

Isotopes

Isotopes are atoms of the same element having the same atomic number but different mass numbers.

They differ in the number of neutrons.

Example:

• Carbon-12
• Carbon-13

Isotopes have:

• Same atomic number
• Same number of protons
• Different mass numbers
• Different number of neutrons
• Similar chemical properties
• Different physical properties

Isotopes may be stable or unstable.

Unstable isotopes that emit radiations are called radioactive isotopes.

Applications Of Isotopes In Daily Life

Radioactive isotopes are used in medical imaging and diagnosis.

Isotope	Use
Carbon-11	Brain scan to trace glucose metabolism
Fluorine-18	Brain scan to trace glucose metabolism
Phosphorus-32	Detection of eye tumours
Chromium-51	Diagnosis of anaemia and imaging of spleen and gastrointestinal tract
Iron-59	Bone marrow function and diagnosis of anaemia
Gallium-67	Whole-body scan for tumours
Selenium-75	Pancreas scan
Krypton-81m	Lung ventilation scan
Strontium-81	Bone disease scan
Technetium-99m	Brain, liver, kidney, bone scans and diagnosis of heart disease
Iodine-131	Diagnosis of thyroid malfunction
Mercury-197	Brain scan
Thallium-201	Heart scan and exercise stress test

For isotopic elements, average atomic mass is calculated because isotopes occur in fixed proportions in nature.

Isobars

Isobars are atoms of different elements having different atomic numbers but the same mass number.

Isobars	Protons	Neutrons	Mass Number
Chlorine-37	17	20	37
Argon-37	18	19	37
Cerium-76	32	44	76
Selenium-76	34	42	76
Iron-58	26	32	58
Nickel-58	27	31	58

Isotones

Isotones are atoms having the same number of neutrons but different numbers of protons.

In isotones:

• Number of neutrons is same.
• Number of protons is different.
• Number of electrons is different.

Examples:

• ³₁H and ⁴₂He
• ¹⁴₆C, ¹⁵₇N and ¹⁶₈O

Important Conceptual Points

• Atom is divisible and contains charged particles.
• Electron was discovered by J.J. Thomson.
• Proton is linked with canal rays and was named by Ernest Rutherford.
• Neutron was discovered by James Chadwick in 1932.
• Electron has negative charge.
• Proton has positive charge.
• Neutron has no charge.
• Alpha particles are helium nuclei.
• Nucleus contains protons and neutrons.
• Almost the whole mass of atom is concentrated in the nucleus.
• Rutherford discovered the nucleus through gold foil experiment.
• Bohr’s model explains stable orbits of electrons.
• Maximum number of electrons in a shell is 2n².
• Atomic number equals number of protons.
• Mass number equals protons plus neutrons.
• Isotopes have same atomic number but different mass number.
• Isobars have same mass number but different atomic number.
• Isotones have same number of neutrons.

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