The nucleus of an atom is composed of protons (positively charged) and neutrons (neutral). The electrons (negatively charged) revolve around the centre of the nucleus, forming the atom’s atomic structure.

The concept of atoms and the development of quantum mechanics can be traced back to Democritus, who was the first to suggest that matter was composed of atoms. John Dalton then furthered this idea in the 1800s by providing the first scientific theory of atomic structure. This knowledge has been essential in understanding the chemical reactions, bonds and physical properties of matter.

Quick Revision of Atomic Structure for JEE

Atomic Structure Quick Revision for JEE

Structure of Atom: Important Topics

Structure of Atom - Important Topics

Table of Contents

What is Atomic Structure?

Atomic Models

Dalton’s Atomic Theory

Thomson Atomic Model

Rutherford Atomic Theory

Subatomic Particles

Atomic Structure of Isotopes

Bohr’s Atomic Theory

Dual Nature of Matter

Frequently Asked Questions

The discovery of subatomic particles has been a major breakthrough in the fields of atomic structure and quantum mechanics, leading to the uncovering of other fundamental particles. This discovery has been the foundation for numerous other discoveries and inventions.

Atomic Structure refers to the structure of atoms, which are the basic units of matter. Atoms are composed of protons, neutrons, and electrons, which are held together by electromagnetic forces.

The atomic structure of an element refers to the constitution of its nucleus and the arrangement of the electrons around it. Primarily, the atomic structure of matter is made up of protons, electrons and neutrons.

The nucleus of the atom is composed of protons and neutrons, while the electrons belonging to the atom surround it. The atomic number of an element is determined by the total number of protons in its nucleus.

Neutral atoms have an equal number of protons and electrons. However, when atoms gain or lose electrons, they become charged and are then known as ions, which can increase their stability.

Atoms of different elements have different atomic structures because they contain different numbers of protons and electrons. This is the reason for the unique characteristics of different elements. Learn more about protons and electrons.

Atomic Models

In the 18th and 19th centuries, many scientists contributed to the development of the modern atomic model by attempting to explain the structure of the atom with the help of atomic models. These models had their own advantages and disadvantages, and the most significant contributions to the field were made by John Dalton, J.J. Thomson, Ernest Rutherford, and Niels Bohr. Their ideas on the structure of the atom are discussed in this subsection.

Dalton’s Atomic Theory

John Dalton, an English chemist, proposed that all matter is composed of indivisible and indestructible atoms. He also postulated that all atoms of a single element are identical, but atoms of different elements differ in size and mass.

According to John Dalton’s atomic theory, chemical reactions involve a rearrangement of atoms to form products. His postulates state that atoms are the smallest particles responsible for chemical reactions to take place.

These are the postulates of his theory:

Every matter is composed of atoms.

Atoms are indivisible.

Specific elements consist of only one type of atom.

The mass of each atom is constant and varies depending on the element.

Atoms rearrange themselves during a chemical reaction.

Atoms can neither be created nor destroyed, however they can be transformed from one form to another.

John Dalton’s atomic theory successfully explained the Laws of chemical reactions, including the Law of Conservation of Mass, Law of Constant Properties, Law of Multiple Proportions, and Law of Reciprocal Proportions.

Disadvantages of Dalton’s Atomic Theory

The theory failed to account for the presence of isotopes.

Nothing about the structure of atom was adequately explained.

The scientists later discovered particles inside the atom, proving that atoms are divisible.

The discovery of subatomic particles inside atoms led to a better understanding of chemical species. The discovery of various subatomic particles is as follows:

Thomson Atomic Model

The English chemist Sir Joseph John Thomson proposed his model of atomic structure in the early 1900s.

He was later awarded the Nobel Prize for the discovery of “electrons”, which was based on the cathode ray experiment. The construction and working of the experiment is as follows:

Cathode Ray Experiment

It has a tube made of glass which has two openings, one for the vacuum pump and the other for the inlet through which a gas is pumped in.

The role of the vacuum pump is to maintain a partial vacuum inside the glass chamber. A high voltage power supply connected to electrodes (i.e. cathode and anode) [1] is fitted inside the glass tube.

Conclusions:

The ‘Fluorescent spots’ on the ZnS screen used confirmed that when a high voltage power supply was switched on, there were rays emerging from the cathode towards the anode. These rays were called “Cathode Rays”.

In the presence of an external electric field, the cathode rays are deflected towards the positive electrode; however, when there is no electric field, the cathode rays travel in a straight line.

When rotor Blades are placed in the path of the cathode rays, they appear to rotate. This demonstrates that the cathode rays consist of particles with a certain mass, implying that they possess some energy.

Thompson concluded that cathode rays are made of negatively charged particles called “electrons” based on all the evidence he had.

Thomson found that the charge to mass ratio (e/m) of electrons was 17588 x 10^11 e/bg when electric and magnetic fields were applied to cathode rays (electrons).

From Mullikin’s oil drop experiment, the charge of an electron was determined to be 1.6 × 10^-16 C and its mass was calculated to be 9.1093 × 10^-31 kg.

Conclusions:

Thomson described the atomic structure as a positively charged sphere with negatively charged electrons embedded within it, based on conclusions from his cathode ray experiment.

The “plum pudding model” of the atom is so-called because it can be likened to a plum pudding dish, with the pudding representing the positively charged atom and the plum pieces representing the electrons.

Thomson’s atomic structure described atoms as having equal positive and negative charges, i.e. being electrically neutral.

Limitations of Thomson’s Atomic Structure:

Thomson’s atomic model does not provide an explanation for the stability of an atom. Additionally, other subatomic particles discovered after Thomson’s model was proposed, could not be incorporated into his atomic model.

Rutherford Atomic Theory

Rutherford’s atomic model was based on the Alpha ray scattering experiment, and it was modified with the discovery of another subatomic particle called the “Nucleus”.

Alpha Ray Scattering Experiment

Construction:

A gold foil of 1000 atoms thick is very thin.

Alpha rays (doubly charged Helium He2+) were made to bombard the gold foil.

The gold foil is placed behind the ZnS screen.

Conclusions:

Most of the rays passed through the gold foil, causing scintillations (bright spots) to appear on the ZnS screen.

A few rays were reflected after hitting the gold foil.

One in 1000 rays got reflected by an angle of 180° (retraced its path) after hitting the gold foil.

Conclusions:

Rutherford concluded that most of the space inside the atom is empty, since most of the rays passed through.

Few rays were reflected due to the repulsion between the positive charge in the atom and another positive charge.

A 1/1000th of rays were strongly deflected due to a very strong positive charge in the center of the atom. He referred to this strong positive charge as the “nucleus”.

He said that most of the charge and mass of the atom is located in the Nucleus.

Rutherford’s Structure of the Atom

Rutherford proposed his own atomic structure, based on the observations and conclusions made above, as follows.

The nucleus is located at the center of an atom, where most of its charge and mass are concentrated.

Atomic structure is generally spherical.

Electrons orbit the nucleus in a similar fashion to the way planets orbit the sun, in a circular path.

Limitations of the Rutherford Atomic Model

If electrons revolve around the nucleus, they must expend energy to counteract the strong force of attraction from the nucleus. As they expend more energy, they eventually lose all of their energy and fall into the nucleus, which does not explain the stability of the atom.

If electrons continuously revolve around the ’nucleus’, a continuous spectrum is expected; however, a line spectrum is observed in reality.

Atomic Structure - Rutherford’s Model and J.J Thomson Model

Subatomic Particles

Protons

Protons are positively charged subatomic particles. The charge of a proton is 1e, which is equivalent to approximately 1.602 x 10^-19 Coulombs.

The mass of a proton is approximately 1.672 x 10^-24

Protons are over 1800 times heavier than electrons.

The atomic number of an element is always equal to the total number of protons in the atoms of that element.

Neutrons

The mass of a neutron is almost the same as that of a proton, approximately 1.674×10^-24 grams.

Neutrons are electrically neutral particles and have no charge.

Different isotopes of an element have the same number of protons, but the number of neutrons present in their respective nuclei can vary.

Electrons

The charge of an electron is approximately -1.602 × 10^-19, which is equal to -1e.

The mass of an electron is approximately 9.1 x 10^-31.

Electrons are not taken into account when determining the mass of an atom due to their extremely small mass.

Atomic Structure of Isotopes

Nucleons, consisting of protons and neutrons, are the components of the nucleus of an atom. Each element has a unique atomic number which describes the number of protons in it. However, the total number of nucleons in the atomic structure of an element can vary.

Isotopes of an element are variants of elements that have a different nucleon number (also known as the mass number). This means that the isotopes of an element have the same number of protons, but differ in the number of neutrons.

The atomic structures of the three naturally occurring hydrogen isotopes, protium, deuterium and tritium, can be described using their respective chemical symbols, atomic numbers and mass numbers. For example, protium has a chemical symbol of H, an atomic number of 1 and a mass number of 1; deuterium has a chemical symbol of D, an atomic number of 1 and a mass number of 2; and tritium has a chemical symbol of T, an atomic number of 1 and a mass number of 3. This article provides more information about these isotopes.

Isotopes of an element vary in their stability and half-lives, however they usually have the same chemical behavior due to their similar electronic structures.

![Atomic Structure of Isotopes]()

Atomic Structures of Some Elements

The structure of an element’s atom can be represented by the number of protons, electrons, and neutrons it contains. Below are a few examples of atomic structures:

Hydrogen

The most abundant isotope of hydrogen on the planet Earth is protium. Its atomic number is 1 and its mass number is 1.

Structure of Hydrogen Atom: This implies that it contains one proton, one electron, and no neutrons (total number of neutrons = 0, since the atomic number of Hydrogen is 1).

Carbon

Carbon has two stable isotopes - 12C and 13C. Of these, 12C has an abundance of 98.9% and contains 6 protons, 6 electrons, and 6 neutrons.

Structure of Carbon Atom: Carbon is an element with four electrons in its outermost shell (valence shell). This tetravalency allows it to form a multitude of chemical bonds with other elements. Learn more about Carbon.

Oxygen

The three stable isotopes of oxygen are 18O, 17O, and 16O, with oxygen-16 being the most abundant.

Structure of Oxygen Atom: Oxygen atom has an atomic number of 8 and a mass number of 16, meaning it consists of 8 protons and 8 neutrons. Of the 8 electrons, 6 are located in the valence shell. More information on Oxygen atoms can be found here.

Bohr’s Atomic Theory

In 1915, Neils Bohr proposed his model of the atom which is based on Planck’s theory of quantization. This model is the most widely used model to describe the atomic structure of an element.

Postulates:

The electrons inside atoms are placed in discrete orbits referred to as “stationary orbits”.

The quantum numbers of these shells can be used to represent their energy levels.

Electrons can absorb energy to jump to higher energy levels, and lose or emit energy to move to lower energy levels.

As long as an electron remains in its stationary state, there will be no absorption or emission of energy.

Electrons revolve around the nucleus in these stationary orbits only.

The energy of the stationary orbits is quantized.

Limitations of Bohr’s Atomic Theory:

Bohr’s atomic structure only applies to single-electron species like H, He+ , Li++ , Be+++ , etc.

When the emission spectrum of hydrogen was observed under a more accurate spectrometer, it was found that each line spectrum was composed of numerous smaller discrete lines.

Neither Stark effect nor Zeeman effect could be explained using Bohr’s theory.

Heisenberg’s Uncertainty Principle: Werner Heisenberg proposed that it is impossible to measure both the position and momentum of a particle with absolute accuracy and precision at the same time. There will always be some level of uncertainty associated with the measurement.

Advantage: Position and momentum are two conjugate quantities that were accurately measured by Bohr (theoretically).

Stark Effect: Phenomenon of electrons being deflected in the presence of an electric field.

The Zeeman Effect: A phenomenon in which electrons are deflected in the presence of a magnetic field.

Dual Nature of Matter

Thomas Young proved that electrons, which were previously thought to be particles, also have wave nature through the double slit experiment, which provided evidence of the photoelectric effect.

De-Broglie concluded that, since nature is symmetrical, light and any other matter wave should also be symmetrical.

Quantum Numbers

Principal Quantum Number (n): It denotes the orbital or shell number of an electron.

Azimuthal Quantum Numbers (*l*): It denotes the orbital (sub-orbit) of the electron.

Magnetic Quantum Number: It denotes the number of energy levels in each orbit.

Spin Quantum Number(s): It denotes the direction of spin, with S = -½ corresponding to an anticlockwise spin and S = ½ corresponding to a clockwise spin.

Electronic Configuration of an Atom

The electrons have to be filled in the s, p, d, f orbitals in accordance with the following rule.

1. Aufbau’s Principle: The filling of electrons should take place in ascending order of orbital energy.

Higher energy orbitals should be filled last and lower energy levels first.

The energy of orbital α(p + l) when two orbitals have the same (n + l) value is E α n.

Ascending order of energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, 8s, 5g, 6f, 7d, 8p

2. Pauli’s Exclusion Principle: No two electrons can have the same set of four quantum numbers, and if two electrons must occupy the same energy state, they must have opposite spins.

3. Hund’s Rule of Maximum Multiplicity: When filling degenerate (same energy) orbitals, all degenerate orbitals must be filled singly before any pairing occurs.

Atomic Structure: Solved Problems and Solutions

Important Questions about Atomic Structure

Revision of the Structure of Atom Class 11

Top 12 Most Important JEE Main Questions on the Structure of Atom

Frequently Asked Questions On Atomic Structure

What are Subatomic Particles?

Subatomic particles are the particles that make up an atom. Typically, these particles are protons, electrons, and neutrons.

What are the differences in atomic structures between isotopes?

The number of neutrons present in the nucleus of an atom is determined by its nucleon number, and varies from atom to atom.

What are the Limitations of Bohr’s Atomic Model?

The Bohr model of the atom offered insufficient predictions for larger atoms, and could not explain the Zeeman effect. It was only successful in explaining the hydrogen spectrum.

What is the process for determining the total number of neutrons in the nucleus of a given isotope?

The mass number of an isotope is given by the sum of the total number of protons and neutrons in it. The atomic number describes the total number of protons in the nucleus. Therefore, the number of neutrons can be determined by subtracting the atomic number from the mass number.