A major contribution to the discovery of subatomic particles is related to discharge tubes. Faraday used Cathode-ray tubes with a long glass body filled with gas with two electrodes at both ends made of metals. Under low-pressure and high voltages, it was observed that a stream of particles moves from the negative electrode to the positive one. These are the cathode rays. Experiments done by J.J. Thomson using the cathode-ray tubes concluded that the rays are made up of negatively charged particles called electrons.

R.A. Millikan had found through his famous oil-drop experiment the charge on an electron to be1.67*10^-19C. Using that value, Thomson calculated the mass to charge ratio of electrons.

Where e = charge on an electron

m_e= 9.1094 * 10^-31kg = mass of an electron

Ernest Rutherford modified the cathode ray tube for experiments. He found the positively charged canal rays and discovered the proton. Chadwick discovered neutrons using the bombardment of beryllium using particles.

To explain the properties and stability of an atom after discovering its constituents, many models came up. Thomson’s model of the atom suggested that electrons were embedded in a positively charged sphere like a watermelon. This is the plum pudding model of the atom.

The well-known particle scattering experiment by Rutherford suggested that an atom had a positively charged sphere at its center and electrons revolving around it in orbits. The center is the point of concentration of the mass of the atom. The particles are held by strong electrostatic forces. This model had certain limitations like failure to explain the stability of atoms, distribution of electrons, and the energy of electrons. 

The number of protons in an atom is the atomic number. The number of nucleons (protons+neutrons) is the mass number. Isotopes were elements having the same atomic number and different mass numbers. If the mass numbers are constant, they are isobars, and those having an equal number of neutrons are isotones.

Maxwell’s study suggested electromagnetic radiation contained alternating oscillating electric and magnetic fields. Later, the electromagnetic spectrum was structured with visible light in it. 

The wave nature of EM radiation was able to explain concepts such as interference and diffraction. But, the failure to explain the blackbody radiation photoelectric effect, the spectrum of atoms, and the difference in heat capacity of solids put forward the particle nature of the radiations. 

Einstein suggested that the light falling on certain materials leads to the ejection from the surface, called the photoelectric effect. He added that these radiations are small packets of energy called quanta. Building upon this concept, De Broglie suggested the dual nature of particles and gave the equation:

where =wavelength of the radiation

H = Planck’s constant (6.626 * 10^-34Js)
P = momentum

The range of lights arranged in the order of their wavelengths is the spectra. When atoms emit a range of energies, it is called emission spectra. These could be observed as lines or bands. The line spectrum of hydrogen suggested that the electromagnetic radiation is indeed quantized. The hydrogen spectrum is made up of Lyman, Balmer, Paschen, Brackett, and Pfund series in the order of wavelengths. 

Series	Spectral Region
Lyman	Ultraviolet
Balmer	Visible
Paschen	Infrared
Brackett	Infrared
Pfund	Infrared

Niels Bohr had developed a model to describe the structure of atoms using the quantization of energy. The major postulates of this model are:

1. The electrons in hydrogen atoms move in circular paths around the positively charged nucleus. These are called orbits. They have fixed energy, hence are termed stationary states.

2. Energy in a respective orbit is constant. The energy is conserved. Only transition is allowed between the allowed energy states. This happens when the hydrogen atom absorbs or releases energy. 

3. The frequency of the absorbed or emitted radiation depends on the energy difference.
   

4. Angular momentum in each orbit is quantized in terms of the modified Planck’s constant.
   
   Where n = represents the orbital number

Important concepts of Bohr’s theory

3. The stationary states are represented using n; it is called the principal quantum number. The n can have values from 1,2,3...n. As the number increases, so does the size of the atom.

4. The radius of each orbit is represented as r_n= r₀n²
   wherer ₀ is a constant having the value 52.9 pm

5. The energy of the stationary state is given as
   
   Where R is the Rydberg constant 2.18 * 10^-18J

4. Bohr’s theory can be applied to one-electron systems of ions like He+, Li2+, etc.

5. The velocities of electrons can also be calculated in this model, besides the fact that the path of motion of electrons is not defined.
   The hydrogen spectrum is explained by the transition of electrons in hydrogen atoms in the respective orbits.

Demerits of the model:

• Failure to explain fine structure spectra
• Failure to explain chemical bonding
• Failure to explain atoms of high atomic number

De Broglie explained that matter possesses both wave and particle nature. This is the dual nature of matter.

Heisenberg stated that we cannot exactly measure the position and velocity of an electron simultaneously. It is valid only for microscopic systems. Both these ideas gave rise to the quantum mechanical model of the atom. It emphasized the three-dimensional region called orbitals around the nucleus, where we could predict the probability of finding an electron. There are s, p, d, and f orbitals, and each of them has a varied shape.

The Quantum mechanical model of the atom tried to overcome the limitations of Bohr’s model. In the model, Schrödinger's equation is the basic concept. It gave rise to the mathematical function called wave function to describe the nature of electrons. The wave function is a formulation having no physical significance. But the square of the wave function gives the probability of finding an electron. If it is normalized and the value is 1, it represents that there is a high chance of finding an electron in the region, and if it is zero, it indicates that there is a node and we cannot find a particle at that position. Using Schrödinger's equation, a hydrogen model was studied, and the quantum numbers emerged.

Quantum Number	What it represents	Additional information
Principal Quantum number, n	the shells K, L, M, N, etc.	The number of electrons in each shell is n2
Azimuthal quantum number, l	the orbitals s,p,d,f	The range of values starts from 0,1 to n-1
Magnetic quantum number, ml	Orientation of the orbital	Have 2l+1 values
Spin magnetic number, s	Spin value	s = \( \pm \)12

1. What are the main topics discussed in this chapter?  

Bohr’s model, Quantum mechanical model and quantum numbers, Heisenberg’s uncertainty principle, Photoelectric effect, the Dual nature of matter, orbit, and orbitals are the important topics discussed in the chapter.

2. Explain the photoelectric effect in simple words.

When light having a suitable frequency hits some metal surfaces, electrons are emitted from it. This emission is directly dependent on the frequency of the incident radiation and as the intensity increases, more electrons eject out and move with a certain kinetic energy. Also, the incident frequency must be higher than the minimum frequency required for ejection, called threshold frequency.

3. What is the importance of NCERT Class 11 Chemistry Chapter 2?

In Class 11 Structure Of Atom is an important topic related to atomic physics and nuclear physics. It’s a lengthy chapter, but the facts and concepts will prepare the student for higher studies.

4. Is the Structure of Atom an important chapter in chemistry?

This chapter is important from the perspective of competitive exams as well as for science graduates. There are an ample number of problems and concepts that the student needs to work out. So, students must emphasize their study on the important topics.

5. Does this site contain the Structure of Atom Class 11 NCERT solutions? 

Yes, on this website you could find answers to all the questions in the text in a very detailed manner. Also, this helps to guide the student to attempt and practice more questions in the scope of this chapter. Practising problems is the key to scoring well in competitive examinations. Step-by-step solutions are available here so that the student could follow these contents very effectively.