Structure Of Atom Class 11 Notes Chemistry Chapter 2 - CBSE

Chapter : 2

What Are Structure Of Atom ?

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    Discovery Of Sub-atomic Particle Electron

    An insight into the structure of atom was obtained from the experiments on electrical discharge through gases. In mid 1850s, Faraday began to study electrical discharge in partially evacuated tubes, known as cathode ray discharge tubes.

    Properties Of Cathode Rays

    • The cathode rays start from cathode and move towards the anode.
    • These rays themselves are not visible but their behaviour can be observed with the help of certain kind of
      materials (fluorescent or Phosphorescent) which glow when hit by them.
    • Cathode rays travel in straight line in the absence of electrical or magnetic field.
    • In the presence of electrical or magnetic field, the behaviour of cathode rays are similar to that expected from negatively charged particles, suggesting that the cathode rays consist of negatively charged particles, called electrons.
    • The characteristics of cathode rays (electrons) do not depend upon the material of electrodes and the nature
      of the gas present in the cathode ray tube.

    Charge On The Electron

    R.A. Millikan devised a method known as oil drop experiment to determine the charge on the electrons.

    Charge of an electron (e) = -1.6022 × 10-19 C

    $$\text{The mass of electron (m}_e) =\frac{e}{\frac{e}{m_{e}}}\\=\frac{1.6022 × 10^{-19}\text{C}}{1.758820 × 10^{11}\text{C}\space \text{kg}^{\normalsize-1}}$$

    = 9.1094 × 10-31 kg

    Discovery Of Protons And Neutrons

    Electrical discharge carried out in the cathode ray tube led to the discovery of canal rays carrying positively charged particles.

    Characteristics Of Positively Charged Particles

    • The mass of positively charged particles depends upon the nature of gas present in the cathode ray tube.
    • The charge to mass ratio of the particles depends on the gas from which these originate.
    • Some of the positively charged particles carry a multiple of the fundamental unit of electrical charge.
    • The behaviour of these particles in the magnetic or electrical field is opposite to that observed for electron or cathode rays.


    The smallest and lightest positive ion was obtained from hydrogen and was called proton.

    Mass of proton = 1.676 x 10-27 kg

    Charge on a proton = (+) 1.602 x 10-19 C


    It is a neutral particle. It was discovered by Chadwick (1932). By the bombardment of thin sheets of Beryllium with fast moving a-particles he observed that highly penetrating rays consist of neutral particles which were named neutrons.

    Properties of Fundamental Particles

    Name Symbol Absolute charge / C Relative Charge Mass / kg Mass / u Approx. mass / u
    Electron e -1.602176 × 10-19 -1 9.109382 × 10-31 0.00054 0
    Proton p +1.602176 × 10-19 +1 1.6726216 × 10-27 1.00727 1
    Neutron n 0 0 1.674927 × 10-27 1.00867 1

    Thomson Model Of Atom

    • J. J. Thomson proposed that an atom may be regarded as a sphere of radius approximately 10–10m carrying
      positive charge due to protons and in which negatively charged electrons are embedded.
    • In this model, the atom is visualized as a pudding or cake of positive charge with electrons embedded into it.
    • The mass of atom is considered to be evenly spread over the atom according to this model.

    Drawback of Thomson Model of Atom

    This model was able to explain the overall neutrality of the atom, it could not satisfactorily, explain the results of scattering experiments carried out by Rutherford in 1911.

    Rutherford’s Nuclear Model Of An Atom

    • The positive charge and most of the mass of the atom was densely concentrated in an extremely small region. This very small portion of the atom was called nucleus by Rutherford.
    • The nucleus is surrounded by electrons that move around the nucleus with a very high speed in circular paths called orbits.
    • Electrons and nucleus are held together by electrostatic forces of attraction.

    Atomic Number

    The number of protons present in the nucleus is equal to atomic number (Z).

    Atomic number (Z) = number of protons in the nucleus of an atom = number of electrons in a neutral atom.

    Mass Number

    The number of protons and neutrons present in the nucleus are collectively known as nucleons. The total number of nucleons is termed as mass number (A) of the atom.

    Mass Number (A) = Number of protons (p) + Number of neutrons (n).

    Drawbacks Of Rutherford Model

    • When a body is moving in an orbit, it achieves acceleration. Thus, an electron moving around nucleus in an
      orbit is under acceleration. But according to Maxwell’s electromagnetic theory, charged particles when accelerated must emit electromagnetic radiations. Therefore, an electron should spiral into nucleus within 10-8s. But actually this does not happen. Thus, Rutherford’s model cannot explain the stability of atom if the motion of electrons is described on the basis of classical mechanics and electromagnetic theory.
    • It does not give any idea about distribution of electrons around the nucleus and about their energies.


    Atoms with identical atomic number but different atomic mass number are known as isotopes.

    Name Symbol Atomic Number Mass Number Number of Protons Number of Neutrons
    Protium or hydrogen 1H1 1 1 1 0
    Deuterium or heavy hydrogen 1H2 1 2 1 1
    Tritium 1H3 1 3 1 2

    Developments Leading To The Bohr’s Model Of Atom

    The two developments that played a major role in the formulation of Bohr’s model of atom are :

    • Dual character of the electromagnetic radiation which means that radiations possess both wave like and particle like properties.
    • Experimental results regarding atomic spectra which can be explained only by assuming quantised electronic energy levels in atoms.

     Electromagnetic Wave Theory

    This theory was put forward by James Clark Maxwell in 1864. The main points of this theory are as follows:

    • The energy is emitted from any source (like the heated rod or the filament of a bulb through which electric current is passed) continuously in the form of radiations and is called the radiant energy.
    • The radiations consist of electric and magnetic fields oscillating perpendicular to each other and both perpendicular to the direction of propagation of the radiation.
    • The radiations possess wave character and travel with the velocity of light 3 x 108 m/sec.
    • These waves do not require any material medium for propagation. For example, rays from the sun reach us through space which is a non-material medium.

    Characteristics Of A Wave


    It is defined as the distance between any two consecutive crests or troughs. It is represented by (λ) and its S.I. unit is metre.

    1A° = 10-10m


    Frequency of a wave is defined as the number of waves passing through a point in one second. It is represented by u (nu) and is expressed in Hertz (Hz).

    1 Hz = 1 cycle/sec.


    Velocity of a wave is defined as the linear distance travelled by the wave in one second. It is represented by c and is expressed in cm/sec or m/sec.


    Amplitude of a wave is the height of the crest or the depth of the through. It is represented by V and is expressed in the units of length.


    It is defined as the number of wavelengths per unit length. Its units are reciprocal of wavelength uni t, i.e., m–1

    Electromagnetic Spectrum

    When electromagnetic radiations are arranged in order of their increasing wavelengths or decreasing frequencies, the complete spectrum obtained is called electromagnetic spectrum.

    Limitations of Electromagnetic Wave Theory

    Electromagnetic wave theory was successful in explaining properties of light such as interference, diffraction etc; but it could not explain the following:

    • The phenomenon of black body radiation.
    • The photoelectric effect.
    • The variation of heat capacity of solids as a function of temperature.
    • The line spectra of atoms with reference to hydrogen.

    Photoelectric Effect

    Hertz, in 1887, discovered that when a beam of light of certain frequency strikes the surface of some metals, electrons are emitted or ejected from the metal surface. The phenomenon is called photoelectric effect.

    Observations in Photoelectric Effect

    • Only photons of light of certain minimum frequency called threshold frequency (v0) can cause the photoelectric effect. The value of v0 is different for different metals.
    • The kinetic energy of the electrons which are emitted is directly proportional to the frequency of the striking photons and is quite independent of their intensity.
    • The number of electrons that are ejected per second from the metal surface depends upon the intensity of the striking photons or radiations and not upon their frequency

    Black Body Radiation

    The ideal body, which emits and absorbs all frequencies is called a black body and the radiation emitted by such a body is called black body radiation. The exact frequency distribution of the emitted radiation from a
    black body depends only on its temperature.

    Planck’s Quantum Theory

    To explain the phenomenon of ‘Black body radiation’ and photoelectric effect, Max Planck in 1900, put forward a theory known as Planck’s Quantum Theory. This theory was further extended by Einstein in 1905.

    The main points of this theory was as follows :

    • The radiant energy emitted or absorbed in the form of small packets of energy. Each such packets of energy is called a quantum.
    • The energy of each quantum is directly proportional to the frequency of the radiation i.e. E µ n or E = hn Where, h is a proportionality constant, called Planck’s constant. Its value is equal to 6.626 x 10-34 Jsec.

    Dual Behaviour Of Electromagnetic Radiation

    Electromagnetic radiations have dual nature, i.e., wave nature as well as particle nature. Whenever radiation interacts with matter, it displays particle like properties in contrast to the wave like properties (interference
    and diffraction) which it exhibits when it propagates. Some microscopic particles, like electrons, also exhibit this wave-particle duality.


    When a ray of white light is passed through a prism the wave with shorter wavelength bends more than the one with a longer wavelength. Since ordinary white light consists of waves with all the wavelengths in the visible range, array of white light is spread out into a series of coloured bands called spectrum. The light of red colour which has longest wavelength is deviated the least while the violet light, which has shortest wavelength is deviated the most.

    Continuous Spectrum

    When a ray of white light is analysed by passing through a prism it is observed that it splits up into seven different wide bands of colours from violet to red (like rainbow). These colours are so continuous that each
    of them merges into the next. Hence, the spectrum is called continuous spectrum.

    Emission Spectra

    Emission Spectra occurs when the radiations emitted from a source are passed through a prism and then received on the photographic plate. Radiations can be emitted in a number of ways such as:

    • From sun or glowing electric bulb.
    • By passing electric discharge through a gas at low pressure.
    • By heating a substance to high temperature.

    Line Spectra

    When the vapours of some volatile substance are allowed to fall on the flame of a Bunsen burner and then analysed with the help of a spectroscope. Some specific coloured lines appear on the photographic plate which are different for different substances.

    Absorption Spectra

    When white light is passed through the vapours of a substance and the transmitted light is then allowed to strike a prism, dark lines appear in the continuous spectrum. The dark lines indicate that the radiations
    corresponding to them were absorbed by the substance from the white light. This spectrum is called absorption spectrum.

    Line Spectrum Of Hydrogen

    When electric discharge is passed through hydrogen gas enclosed in discharge tube under low pressure and the emitted light is analysed by a spectroscope, the spectrum consists of a large number of lines which are
    grouped into different series. The complete spectrum is known as hydrogen spectrum. Johannes Rydberg, on the basis of his experimental observations, noted that all series of lines in the hydrogen spectrum could be described by the following expression:

    $$\bar{v} = 109,677\bigg[\frac{1}{n_{1}^{2}}-\frac{1}{n_{1}^{2}}\bigg]\text{cm}^{\normalsize-1}$$

    where n1 = 1,2............

    n2 = n1 + 1,   n1 + 2..........

    The value 10 9,677 cm-1 is called the Rydberg constant for hydrogen.

    A simple theoretical equation, known as Rydberg formula, for the calculation of wavelengths and wave numbers of the spectral lines in different series of hydrogen spectrum can be given as,

    $$\frac{1}{\lambda}* = \bar{v} = R_{\text{H}}\bigg(\frac{1}{n_{1}^{2}}-\frac{1}{n_{2}^{2}}\bigg)$$

    This relation is valid for hydrogen atom only. For other species,

    $$\frac{1}{\lambda} =\bar{v} =\text{R×Z}^{2}\bigg(\frac{1}{n_{1}^{2}} -\frac{1}{n_{1}^{2}}\bigg)$$

    where Z is the atomic number of the species.

    Here RH = constant, called Rydberg constant for hydrogen and n1 , n2 are integers (n2 > n1)

    Bohr’s Model Of Atom

    Niels Bohr in 1913, proposed a new model of atom on the basis of Planck’s Quantum Theory. The main points of this model are as follows :

    • In an atom, the electrons revolve around the nucleus in certain definite circular paths called orbits.
    • Each orbit is associated with definite energy and therefore these are known as energy levels or energy shells.
      These are numbered as 1, 2, 3, 4……….. or K, L, M, N………..
    • Only those energy orbits are permitted for the electron in which angular momentum of the electron is a whole
      number multiple of h/2π

    Angular momentum of electron (mvr) = nh/2π (n = 1, 2, 3, 4 etc).

    m = mass of the electron.

    v = tangential velocity of the revolving electron.

    r = radius of the orbit.

    h = Planck’s constant.

    n is an integer.

    • As long as electron is present in a particular orbit, it neither absorbs nor loses energy and its energy,
      therefore, remains constant.
    • When energy is supplied to an electron, it absorbs energy only in fixed amounts as quanta and jumps to higher energy state known as excited state. The excited state is unstable, the electron may jump back to the lower energy state and in doing so, it emits the same amount of energy. (∆E = E2 – E1).

    Achievements of Bohr’s theory of Atom

    • It can explain the stability of atom.
    • It successfully explains the line specturm of hydrogen.
    • It explains the line ‘spectra of single electron ions like He+ and Li2+.

    Limitations of Bohr’s theory of Atom

    • The theory could not explain the atomic spectra of the atoms containing more than one electron or multielectron atoms.
    • Bohr’s theory failed to explain the fine structure of the spectral lines.
    • Bohr’s theory could not offer any satisfactory explanation of Zeeman effect and Stark effect.
    • Bohr’s theory failed to explain the ability of atoms to form molecule formed by chemical bonds.
    • It was not in accordance with the Heisenberg’s uncertainty principle.

    Dual Behaviour Of Matter (De Broglie Equation)

    $$\lambda =\frac{h}{\text{mv}} =\frac{h}{\text{p}}$$

    where m = mass of the particle

    v = velocity of particle

    p = momentum of the particle

    This relationship has been verified by an experiment.

    Heisenberg’s Uncertainty Principle

    It states that, “It is impossible to determine simultaneously, the exact position and exact momentum (or velocity) of an electron”.

    Mathematically, it can be given as,

    $$\Delta x×\Delta \rho_{x}\geq\frac{h}{4\pi}\\\text{or}\space\Delta x ×\Delta(mV_{x})\geq\frac{h}{4\pi}\\\text{or}\space\Delta x ×\Delta V_{x}\geq\frac{h}{4\pi m}$$

    where ∆x is the uncertainity in position and ∆Px (or ∆Vx) is the uncertainty in momentum (or velocity) of the particle and h is Planc’s constant.

    Significance of Uncertainty Principle

    • It rules out existence of definite paths or trajectories of electrons and other similar particles.
    • The effect of Heisenberg’s uncertainty principle is significant only for microscopic objects and is negligible
      for macroscopic objects.

    Important Features Of Quantum Mechanical Model Of Atom

    • The energy of electrons in atom is quantized i.e., can only have certain values.
    • The existence of quantized electronic energy level is a direct result of the wave like properties of electrons.
    • Both the exact position and exact velocity of an electron in an atom cannot be determined simultaneously.
    • An atomic orbital has wave function φ. There are many orbitals in an atom. Electron occupy an atomic orbital
      which has definite energy. An orbital cannot have more than two electrons. The orbitals are filled in increasing order of energy. All the information about the electron in an atom is stored in orbital wave function Φ.
    • The probability of finding electron at a point within an atom is proportional to square of orbital wave function i.e., |φ2| at that point. It is known as probability density and is always positive.

    Quantum Numbers

    These are used to get complete information about electron i.e., its location, energy, spin etc. There are four quantum numbers :

    • n - principal quantum number: describes the energy level
    • ℓ - azimuthal or angular momentum quantum number: describes the subshell
    • mℓ or m - magnetic quantum number: describes the orbital of the subshell
    • ms or s - spin quantum number: describes the spin

    Quantum Number Values

    According to the Pauli exclusion principle, no two electrons in an atom can have the same set of quantum numbers. Each quantum number is represented by either a half-integer or integer value.

    • The principal quantum number is an integer that is the number of the electron's shell. The value is 1 or higher
      (never 0 or negative).
    • The angular momentum quantum number is an integer that is the value of the electron's orbital (for example,
      s=0, p=1). ℓ is greater than or equal to zero and less than or equal to n - 1.
    • The magnetic quantum number is the orientation of the orbital with integer values ranging from -ℓ to ℓ. So, for the p orbital, where ℓ=1, m could have values of -1, 0, 1.
    • The spin quantum number is a half-integer value that is either -1/2 (called "spin down") or 1/2 (called "spin up").

    Shape Of s-orbital

    • The s-orbital boundary surface diagram resembles a sphere with the nucleus at its center, which can be shown in two dimensions as a circle.
    • s-orbitals are spherically symmetric, which means that the probability of finding an electron at a given distance is the same in all directions.
    • The size of the s-orbital is likewise shown to increase as the value of the primary quantum number (n)
      increases; hence, 4s > 3s > 2s > 1s.

    Shape Of p-orbital

    • The p-orbitals are formed like dumbbells.
    • The p-orbital node is located at the nucleus’s center.
    • Due to the presence of three orbitals, the p orbital can occupy a maximum of six electrons.
    • Each p-orbital is made up of two parts known as lobes that are located on either side of the plane that runs
      across the nucleus.
    • Each p-orbital has parts known as lobes on either side of the plane that runs across the nucleus. At the plane
      where the two lobes intersect, the likelihood of finding an electron is nil.
    • The three orbitals are known as degenerate orbitals because they have the same size, shape, and energy.
    • The sole difference between the orbitals is the orientation of the lobes because the lobes are orientated along the x, y, or z-axis, they are given the names 2px, 2py, and 2pz. The formula n –2 is used to calculate the number of nodes.
    • Similarly to s-orbitals, the size and energy of p-orbitals rise as the primary quantum number increases (4p > 3p > 2p).

    Shape Of d-orbital

    • For d-orbitals, the magnetic orbital quantum number is given as (-2,-1,0, 1,2). As a result, we can claim there
      are five d-orbitals.
    • These orbitals are denoted by the symbols dxy, dyz, dxz, dx2 -y2, and dz2.
    • The forms of the first four d orbitals are similar to each other, which differs from the dz2 orbital, but the energy of all five d-orbitals is the same.

    AUFBAU Principle

    According to the Aufbau principle, “electrons are filled in an orbital in the increasing order of orbital energy level“. In other words, electrons are filled in the orbital in such a way that the orbital with the lowest energy
    gets filled first, then to the next higher energy level, and so on i.e. in the pattern of their increasing energy. The correct sequence of filling of electrons in an orbital depends on two factors:

    • Orbital having the lowest value of (n+1) is filled first.
    • In the case where two orbitals have the same (n+1) value, the orbital with the lowest ‘n’ value will be filled first by electrons.

    The order in which the energies of the orbitals increase and hence the order in which the orbitals are filled is as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, if, 3d, 6p, 7s, 5f 6d, 7p.

    Pauli Exclusion Principle

    According to this principle, no two electrons in an atom can have the same set of four quantum numbers. It can also be stated as: Only two electrons may exist in the same orbital and these electrons must have
    opposite spins.

    Hund’s Rule Of Maximum Multiplicity

    It states that: pairing of electrons in the orbitals belonging to the same subshell (p, d or f) does not take place until each orbital belonging to that subshell has got one electron each i.e., it is singly occupied.