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Atomic structure is the structure of an atom that consists of a nucleus at the center containing neutrons and protons, while electrons are revolving around the nucleus. As atoms are made up of a very tiny, positively charged nucleus that is surrounded by a cloud of negatively charged electrons.

The earliest concept of atoms was given by Indian philosopher Maharshi Kanad who proposed that matter is made up of very small indestructible particles called ‘Parmanu’. A Greek philosopher named Democritus also initially claimed that matter is formed of atoms, and is credited with developing the concepts of atomic structure and quantum mechanics. Later in the 1800s, John Dalton a British Scientist put out the first atomic structure scientific theory.

What is Atomic Structure?

The composition of an element’s nucleus and how its electrons are arranged around it are referred to as the element’s atomic structure. Protons, electrons, and neutrons comprise the majority of the atomic structure of matter.

The atom’s nucleus, which is made up of protons and neutrons, is surrounded by the atom’s own electrons. The total number of protons in an element’s nucleus is expressed by the element’s atomic number. Protons and electrons are equal in number in neutral atoms. But atoms can receive or lose electrons to make them more stable, and the resulting charged object is known as an ion. Because different elements’ atoms contain varying numbers of protons and electrons, their atomic structures are also different. This explains why different elements have unique properties.

Atomic Structure

The atomic model which we study today was not given at one time. Several attempts were made by scientists and later improved leading to the current atomic model. Let’s learn about different atomic models which led to the evolution of the present model.

Atomic Models

Many scientists used atomic models to understand the structure of the atom in the early centuries. Each of these models had advantages and disadvantages of its own and played a significant role in the development of the modern atomic model. Scientists like John Dalton, J.J. Thomson, Ernest Rutherford, and Niels Bohr made the most noteworthy contributions to science.

This section of the article discusses the following theories regarding atomic structure:

  • Dalton’s Atomic Theory
  • Thomson’s Atomic Model
  • Rutherford’s Atomic Model
  • Bohr’s Atomic Model
  • Quantum Mechanical Model

Dalton’s Atomic Theory

John Dalton, a British Chemist proposed that every matter is made up of atoms. These atoms are indivisible and indestructible i.e. they can’t be broken down into smaller particles. He also suggested that all atoms of a particular matter are the same, but atoms of different elements differ in size and mass. This means atoms of each element are unique.

According to Dalton’s atomic theory, Chemical reactions occur at atomic level and involve the rearrangements of atoms in order to form the products. According to the postulates proposed in his theory, atomic structure is made up of atoms and they are the smallest particles responsible for chemical reactions to occur.

Postulates of Dalton’s  Atomic Theory

  • Every matter that exists is made of atoms.
  • Atoms are indivisible.
  • A particular element has only one type of atom in it.
  • Atoms of different elements differ in size and mass.
  • An atom has a constant mass that varies for every element.
  • During a chemical reaction, atoms undergo rearrangement.
  • Atoms can neither be created nor destroyed but can only be transformed from one form to another.

Dalton’s atomic theory was able to explain the Laws of chemical reactions successfully, named the Law of conservation of mass, Law of constant properties, Law of multiple proportions, and Law of reciprocal proportions.

Demerits of Dalton’s Atomic Theory

  • This theory was not able to explain the existence of isotopes and isobars.
  • No appropriate explanation was provided regarding the structure of atoms.
  • Later the atoms were found to be divisible, and Dalton’s claim of atoms being indivisible was proved to be wrong.

The discovery of constituting particles of atoms led to a better understanding of chemicals, these constituting particles are called subatomic particles.

Thomson’s Atomic Model

Sir Joseph John Thomson was also an English chemist famous for his discovery of electrons known as Thomson’s Atomic Model, for which he also got the Nobel Prize. He conducted a cathode ray experiment to invent electrons. He proposed that atoms are like a sphere of positive charge with negative charge embedded in them. He named this atomic model as Plum Pudding Model.

Cathode Ray Experiment

In this experiment, a glass tube with two openings is taken. One opening is for the vacuum pump and the other is for intake through which the gas to be filled in the tube is pushed in. Using the vacuum pump a partial vacuum pump is maintained inside the glass chamber. In simple words, a cathode and anode are placed inside the glass tube. The anode is perforated and a photosensitive foil made up of Zinc Sulfide is placed behind it. When high voltage is applied a ray originates from the cathode and moves towards the anode making a spot on Zinc Sulfide foil.

Thomson's Atomic Model

Observations of Cathode Ray Experiment

The following observations were made when the current was allowed to flow between the cathode and anode.

  • When the high voltage power is connected and switched ON, rays were transmitted from the cathode towards the anode. Fluorescent spots were observed on the ZnS screen and it confirmed the fact the rays were being transmitted. These rays were given the name Cathode Rays.
  • When an external electric field was projected on the tube, the rays got deviated toward the positive electrode. But in the absence of the electric field, the rays got back in a straight line.
  • But when rotor blades were fixed in the path of the cathode rays, the rays seemed to rotate. This proved that cathode rays were made of particles that had some mass in them.
  • Using all the evidence, Thomson reached the conclusion that cathode rays are composed of negatively charged particles called electrons.
  • By applying electric and magnetic fields on the cathode ray, the charge-to-mass ratio (e/m) was found. The e/m for electrons came out to be 17588 × 1011 e/bg

Discovery of Electron

After performing the Cathode Ray Experiment, JJ Thomson explained that the rays that were originating from Cathode and moving towards Anode consists of negatively charged particles called Electron. He further stated that the presence of these negatively charged particles is not limited to specific matters but will be present in every matter irrespective of mass and property. The discovery of the Electron was done in 1897.

Mullikin did an oil-drop experiment to find the charge of the electron using the e/m ratio. He found the charge of the electron = 1.6 x 10-16 C and the Mass of the Electron = 9.1093 × 10-31 kg. 

Plum Pudding Model

After Thomson discovered Electron he attempted to describe the structure of the atom. He postulated that an atom is a positively charged sphere in which negatively charged electrons were embedded. The popular name given to this model is the “plum pudding model” because it can be observed as a plum pudding dish where the positively charged atom signifies the pudding and the plum pieces stand for the electrons. Plum Pudding Model is also sometimes referred to as the Watermelon Model where the red edible part of a watermelon is a sphere of positive charge while the seeds of the watermelon are referred to as negatively charged electrons.

Drawbacks of Thomson’s Atomic Model

The main drawback of Thomson’s model is that this model is not clear about the stability of an atom. This model could not adjust to other subatomic particles discovered in the future.

Rutherford Atomic Model

Rutherford who was a student of J. J. Thomson discovered Nucleus which contained protons and neutrons inside it. This discovery made huge changes to the atomic structure. The observations made by Thomson in his experiment were used by Rutherford to propose his theory for atomic structure through an experiment called Rutherford’s Alpha Ray Scattering Experiment.

Alpha Ray Scattering Experiment

Rutherford used the radioactivity phenomenon in conducting his experiment. He used the radioactive material radium bromide (RaBr). RaBr emits α particles which is a form of radiation. A thin gold metal sheet was put up in the setup. Then the alpha ‘α’ particles were bombarded on this sheet. The α particle has a charge of +2. To observe the deflection of the particles a screen of Zinc Sulfide (ZnS) was used and placed behind the Gold foil. Rutherford further developed a detector in order to count the number of radioactive particles. Initially, he recorded the count rate of RaBr as he kept a count of α particles emitted per minute.

Rutherford Atomic Model

Observation of Alpha Ray Scattering Experiment

Following observations were made by Rutherford and conclusions were drawn:

  • Most of the α particles passed through thin sheets. This means most of the atom’s space is empty.
  • Another observation made was that some of the α  particles deflected a bit in every direction. This leads to the conclusion that the positive charge is not distributed uniformly throughout the atom.
  • Very few α particles get deflected back along the path that they were traveling on. This happened because of charges repelling each other. Seeing this Rutherford concluded that the positive charge in an atom exists in a very small volume.
  • Not only the positively charged particles but a lot of mass is also concentrated in a very small volume. Rutherford named this region as Nucleus.
  • Rutherford also came up with the argument that electrons are present in orbits around the orbits, much like the planets in the solar system. Electrons are negatively charged and they revolve around the nucleus.
  • The electrons and nucleus are held by the electrostatic force of attraction because they are negatively and positively charged respectively.

Conclusion of Rutherford’s model

Drawing conclusions from all the above observations, Rutherford proposed his Atomic structure which had the following properties –

  • The nucleus lies at the center of the atom, and the maximum of the charge and mass is concentrated there only.
  • Atoms are spherical in nature.
  • Electrons revolve around the nucleus in a circular orbit.

Discovery of Nucleus

In Gold Foil Alpha Particle Experiment, Rutherford observed that most of the spaces inside an atom are vacant and there is a small dense region located at the center inside the atom. He termed this region as Nucleus and said that this Nucleus is positively charged and most of the masses of the atom are concentrated in Nucleus only.

Limitations of Rutherford Atomic Model

Just like other atomic models, Rutherford’s model also had many shortcomings.

  • Since electrons revolve in a circular orbit around the nucleus in an atom it is an accelerated motion. As per Electromagnetic Theory when a charged particle is in accelerated motion it loses energy. Hence, electrons will spend a lot of energy and eventually, they will lose the entire energy and the atom will collapse. This raises serious questions about the stability of the atom.
  • Rutherford didn’t say anything about the position of electrons whether all electrons be in the same or different orbits and the reason behind it.
  • If the electrons are revolving continuously around the nucleus, then the spectrum that they emit should be a continuous spectrum, but what we observe is a line spectrum.

Bohr’s Atomic Model

Neils Bohr, a student of Rutherford proposed his model in 1915 to address the limitation of Rutherford’s Atomic Model. It is the most widely used atomic model and is based on Planck’s theory of quantization. It explains that electrons always move in fixed orbitals only, and they are not present everywhere in the atom. Bohr also explained that each orbit has a fixed energy level. An orbit is also called an Energy Shell. Rutherford only explained the nucleus of the atom while Bohr made changes to that model and added electrons and energy levels.

Bohr’s Atomic Theory

As per Bohr’s model, inside an atom, there is a small nucleus that is positively charged and is surrounded by negative electrons which move around in orbits which has specific energy level. To revolve in a particular orbit, electrons must possess energy equal to the energy level of the shell. Bohr found out that the larger the distance of an electron from the nucleus, the larger its energy which means the orbits near the nucleus has smaller energy and the shell farthest from the nucleus has larger energy.

Postulates of Bohr’s Atomic Theory

  • Inside atoms, electrons are present in discrete orbits called “stationary orbits”.
  • Quantum numbers are used to represent the energy levels of these shells.
  • Electrons can go to higher levels by absorbing energy and move to lower energy levels by losing or emitting some energy.
  • When an electron stays in its own orbit, no absorption or emission of energy takes place.
  • Electrons revolve in these stationary orbits only.
  • The energy of the stationary orbits is quantized.

Limitations of Bohr’s Atomic Theory

  • It works only for single-electron species such as H, He+, Li2+, and Be3+
  • When a more accurate spectrometer was used to observe the emission spectrum of hydrogen, each line spectrum was seen to be a combination of multiple smaller discrete lines.
  • Bohr’s theory was unable to explain Stark and Zeeman’s effects.

Quantum Mechanical Model of Atom

Quantum Mechanics is the branch of physics that deals with the motion and kinematics of microscopic objects. Since atoms are of below microscopic size and the limitations of Bohr’s Atomic Model motivated the scientists to give a more general and accurate atomic model based on Quantum Theory. The Quantum Mechanical Model of the Atom basically uses two following theories to explain the structure of the atom:

  • Dual Behaviour of Matter
  • Heisenberg Uncertainty Principle

Dual Behaviour of Matter

Dual Behaviour of Matter was proposed by French physicist de-Broglie. He stated that every matter irrespective of its size exhibits both wave-like properties and particle-like properties. He meant to say that just like a photon has both wavelength and momentum similarly an electron will have both wavelength(λ) and momentum(p).  He called these waves Matter Waves. The relation between wavelength and momentum is given by

 λ = h/p 

where, 

  • λ is Wavelength
  • p is Momentum
  • h is Planck’s Constant

Heisenberg Uncertainty Principle

Heisenberg’s Uncertainty Principle states that when a microscopic particle is in motion it is impossible to find the exact position and momentum of the particle simultaneously. He meant that at a time we can find either position or momentum i.e. if the exact position is known then momentum is uncertain and vice-versa. It is represented as 

 Δx.Δp ≥ h/4π

where,

  • Δx is Uncertainty in Position
  • Δp is Uncertainty in Momentum
  • h is Planck’s Constant

From the formula, it means that if Δ for the position is very small i.e. if the position is known exactly then Δp will be very large hence, physically we will have a blurred image of the measurement. Hence, it talks about probability which is the basis of the Quantum Mechanical Model of Atom.

Although the above two concepts are important for understanding of Quantum Mechanical Model of Atom, it is equally important to know the Schrodinger Wave Equation which was the most fundamental equation of Quantum Mechanics related to the energy of the system.

Schrodinger Wave Equation

Schrodinger Wave Equation gives the equation for the total energy of the system (an atom or a molecule) whose energy doesn’t change with time i.e. there is no loss or gain of energy. Mathematically, Schrodinger Wave Equation is represented as 

Hψ = Eψ

where 

  • H is Hamiltonian Operator in Mathematics
  • E is the Total Energy of the System
  • ψ is a Wave function

The solution of the Schrodinger Wave Equation gives the value of E and ψ.

Postulates of Quantum Mechanical Model of Atom

Quantum Mechanical Model states the following about structure of the atom

  • The energy of electrons in atoms is quantized i.e. energy level of an electron is an integral multiple of the smallest energy quantity.
  • Quantized Energy levels exist due to Wave like Properties of electrons and their solution can be given by Schrodinger Wave Equation.
  • Since it is impossible to find the position and momentum of an electron simultaneously therefore we talk about the probability of different physical points related to the electron.
  • Atomic Orbital of an atom is represented by wave function ψ. Each orbital can be occupied by two electrons at maximum. When an electron occupies an orbital it is represented by ψ.
  • Quantum Model states that there is an electron cloud around the nucleus inside an atom. The probability to find an electron inside an atom is given by |ψ|2, which is called Probability Distribution Function.

Quantum Number

To describe the location of an electron inside an atom we need a set of 4 numbers known as Quantum Numbers. These Quantum Number helps in distinguishing different orbitals which contain electron. Orbitals that have a smaller number mean they are closer to the nucleus, have a smaller size, and have a larger probability of finding an electron. The four types of Quantum Numbers are listed below:

  • Principal Quantum Number
  • Azimuthal Quantum Number
  • Magnetic Quantum Number
  • Spin Quantum Number

Principal Quantum Number(n)

It is represented by ‘n’. It gives the idea of a shell in which an electron is present and also about its energy. A lower value of ‘n’ means the shell is closer to the nucleus and has lower energy. It is given by n = 1,2,3…

nShell
1K
2L
3M
4N

Azimuthal Quantum Number(l)

It is represented by ‘l’. It gives an idea of the subshell and 3D shape of the orbital. The subshells are given as s, p, d, and f. The value assigned to subshells are 0 = s, 1 = p, 2 = d, 3 = f. A shell contains 0 to n-1 subshell. For Example, the third shell i.e. n = 3 will have 0 to (3-1) subshells i.e. 0 to 2 subshells which include 0,1 and 2.

nlSubshell Notation
101s
202s
212p
303s
313p
323d

Magnetic Quantum Number(ml)

It is represented by ml. It gives an idea of the orientation of orbital in space with respect to coordinate axes. A subshell contains -l to l orbitals. For Example, subshell p contains -1 to 1 orbital i.e. -1, 0, 1, a total of three orbitals oriented along different coordinate axes and coordinate planes.

Spin Quantum Number(ms)

It is represented by ms. It gives an idea about the spinning or orientation of electrons. It takes the value of +1/2 or -1/2. If ms is +1/2 it means the electron is rotating clockwise and is represented as ⇡ while if ms is -1/2 it means the electron is rotating anticlockwise and is represented as ⇣.

As of now, we have learned all the atomic models, now we will look at the properties of all the subatomic particles.

Subatomic Particles

The subatomic particles are the particles that are present inside the atom, There are three subatomic particles that are,

  • Protons
  • Neutrons
  • Electrons

Protons

  • Protons have a positive charge. This charge is 1e, which is approximately 1.602 × 10-19
  • Mass of a proton is approximately 1.672 × 10-24
  • Protons are over 1800 times heavier than electrons.
  • Total number of protons in the atoms of an element and the atomic number of the element is always equal.

Neutrons

  • Mass of a neutron is almost similar to that of a proton i.e. 1.674 × 10-24
  • Neutrons are always electrically neutral particles and do not carry any charge.
  • Isotopes of an element have the same number of protons but a different number of protons in their respective nuclei.

Electrons

  • Charge of an electron is -1e, which is approximately -1.602 × 10-19
  • Mass of an electron is approximately 9.1 × 10-31.
  • Mass of an electron is almost negligible as compared to the mass of an atom, so an electron’s mass is ignored while calculating the mass of an atom.

Isotopes

Isotopes are the atoms of same elements that have the same atomic number but different mass numbers. Example C-12, C-13 and C-14. Here all are Carbon atoms and have the same atomic number i.e. 6 but different mass numbers. This difference in mass numbers can be understood from their atomic structure.

Atomic Structure of Isotopes

The isotopes of an atom have the same atomic number which means that the number of protons are same. Also, their chemical properties are the same because their electronic configuration is the same. The difference in mass number arises due to the difference in number of neutrons present inside the nucleus. Hence, the atomic structure of isotopes comprises of the same number of electrons and protons but different number of neutrons. We can understand this with the example of isotopes of hydrogen illustrated below:

To describe the structure of an isotope, the element’s symbol is used along with the atomic number and the mass number of the isotope. To give an example, Hydrogen has 3 isotopes named protium, deuterium, and tritium. The atomic configuration  of three isotopes of hydrogen is tabulated below:

Isotopes of HydrogenAtomic NumberMass NumberNo. of ElectronsNo. of ProtonsNo. of Neutrons
Protium11110
Deuterium12111
Tritium13112

 

The stability of isotopes is different. The half-lives are also different. But they generally have similar chemical behavior because they have the same electronic structures. The pictorial representation of isotopes of hydrogen can be seen below:

Atomic Structure of Isotopes

Electronic Configuration of Elements

The Electronic Configuration of Elements refers to the arrangement of electrons in different energy levels. The rule for the arrangement of electrons is governed by the following three laws:

  • Aufbau Principle
  • Hund’s Rule
  • Pauli Exclusion Principle

Aufbau Principle

Aufbau is a German word that means ‘to build’. In Chemistry, Aufbau Principle states that the electronic arrangement of an element is done by filling electrons in ascending order of energy of subshell. It means electrons first enter subshells of lower energy and then of higher energy levels. The energy of a subshell is determined by adding Principal Quantum Number and Azimuthal Quantum Number i.e. (n+l). If two subshells have the same (n+l) value then the subshell having a lower value of n is of lower energy. Hence, electrons enter in the order of 1s, 2s, 2p, 3s, 3p, 4s, and 3d…..

Hund’s Rule

Hund’s Rule states that electrons in the subshell in the manner that in the first attempt of filling the subshell is half-filled i.e. each orbital has one electron and then the pairing of electrons is done. This is because half-filled and full-filled orbitals are more stable than incompletely-filled orbitals.

Pauli Exclusion Principle

Pauli Exclusion Principle states that an orbital can have a maximum of two electrons with opposite spin. This is because if two electrons of the same spin are in an orbital then all four quantum numbers will be the same which is not possible as per the Quantum Mechanical Model of the atom.

Summary of Atomic Structure

  • Atom: The defining structure and basic units of matter of an element are called atoms. The term “atom” came from a Greek word that means indivisible because earlier atom was thought to be the smallest things in the universe that could not be divided
  • Atomic Structure: The structure of an atom comprising a nucleus, in which the protons and neutrons are present. The negatively charged particles called electrons revolve around the center of the nucleus.
  • Nucleus: A collection of particles called protons and neutrons is called Nucleus. Protons are positively charged and neutrons, are electrically neutral. Protons and neutrons are made up of particles called quarks. The chemical element of an atom is determined by the number of protons, or the atomic number, Z, of the nucleus.
  • Proton: Positively charged particles found within atomic nuclei are given the name Proton. Rutherford discovered the proton in his famous cathode ray experiment that was conducted between 1911 and 1919. Protons are about 99.86% as massive as neutrons. The number of protons in an atom is unique for each element
  • Electron: Electrons are very tiny compared to protons and neutrons, about 1800 times smaller than either a proton or a neutron. Electrons are just 0.054% as massive as neutrons. Electrons were discovered in 1897 by Joseph John (J.J.) Thomson, a British physicist. Electrons have a negative charge and are electrically attracted to the positively charged protons
  • Neutron: Rutherford theorized the neutron’s existence in 1920 and was later discovered by Chadwick in 1932. Neutrons were found during experiments where atoms were shot at a thin sheet of beryllium. Subatomic particles with no charge were released – and were named neutrons. Neutrons are uncharged particles found within all atomic nuclei
  • Isotopes: Members of the same family of an element that all have the same number of protons but different numbers of neutrons are named isotopes. The number of protons in a nucleus determines the element’s atomic number on the Periodic Table. All the isotopes have unique properties, just like all family members have their own qualities. 

Also, Read

FAQs on Atomic Structure

Q1: What is Atomic Structure?

Answer:

The arrangement of sub atomic particles electron, proton and neutron inside an atom is called atomic structure.

Q2: What is Thomson’s Atomic Model?

Answer:

According to Thomson’s Atomic Model, an atom is a sphere of positive charge with negative charge embedded in it.

Q3: Who discovered Nucleus?

Answer:

Nucleus of Atom by discovered by Ernst Rutherford in 1911 in his Alpha Particle Scattering Experiment.

Q4: What is Rutherford’s Atomic Model?

Answer:

According to Rutherford’s Atomic Model, proton and neutron are confined in a small region inside atom called nucleus and electron revolve around the nucleus.

Q5: What is Aufabu Principle?

Answer:

Aufbau Principle states that the electron enter in subshells of an atom in the order of increasing energy level.



Last Updated : 25 Feb, 2024
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