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CBSE NOTES CLASS 11 CHEMISTRY CHAPTER 2

STRUCTURE OF ATOM

Atoms and subatomic particles

Cathode ray discharge tube experiment

Charge to Mass (e/m) Ratio of Electron

Millikan’s Oil Drop Method and Charge on the Electron

Canal Rays

Characteristics of positively charged particles

Some Terms Related to Atom

Nucleons

Atomic number(Z)

Mass number (A)

Isotopes

Isobars

Isotones

Isoelectronic species

Atomic Models

Thomson’s Model or Water Melon Model or Plum Pudding Model

Merits of Thomson’s Model

Shortcomings of Thomson’s Model

Rutherford’s alpha(α) particle scattering experiment

Rutherford’s Model or Planetary Model

Drawbacks of Rutherford’s Model

Why is the lead plate used in Rutherford’s alpha particle scattering experiment?

Developments Leading to the Bohr’s Model of Atom

Quantum Theory

Wave Nature of Electromagnetic Radiation

Properties of electromagnetic radiations

Wavelength

Frequency

Velocity

Wave number

Particle Nature of Electromagnetic Radiation

Black body radiation

Photoelectric effect

Observations

Threshold frequency

Work function

Kinetic energy of the ejected electrons

Einstein of energy

Dual Behavior of Electromagnetic Radiation

Spectrum

Emission spectrum

Absorption spectrum

Line spectra/atomic spectra of Hydrogen atom

Rydberg Formula

Bohr’s Atomic Model

Orbits or energy states

Frequency of radiation

The angular momentum

Ground State - Bohr radius

Energies of the stationary states associated with hydrogen like species

Radius of orbit of electron

Velocity of electron in an orbit

Frequency of revolution of electron in an orbit

Time Period of revolution of electron in an orbit

Kinetic Energy of electron in an orbit

Potential Energy of electron in an orbit

Total Energy of electron in an orbit

Wave number

Explanation of Line Spectrum for Hydrogen

Limitations of Bohr’s Model of Atom

Quantum mechanics

Schrodinger’s equation

Probability density

Shapes of Atomic Orbitals

s-orbital

p-orbital

d-orbital

Nodal surfaces or Nodes

Radial nodes

Angular nodes

Summary of Quantum Mechanical Model of Atom

Quantum numbers

Principal quantum number

Azimuthal quantum number/ Orbital angular momentum or subsidiary quantum number

Magnetic quantum number or Magnetic orbital quantum number

Electron spin quantum number

Boundary Surface Diagram Energies of Orbitals

In hydrogen and hydrogen like species

In Multi-electron Species

Ground state

Energy and atomic number

Degenerate orbitals

Shielding effect or screening effect

Filling of Orbitals in Atom

The Aufbau Principle

Pauli Exclusion Principle

Hund’s rule of maximum multiplicity

Electronic Configuration of Atoms

Stability of Completely Filled and Half Filled Subshells

Causes of Stability of Completely Filled and Half Filled Sub-shells

CBSE NOTES CLASS 11 CHEMISTRY CHAPTER 2

STRUCTURE OF ATOM

Atoms and subatomic particles

Atom is the smallest indivisible particle of the matter. Atom is made of sub-atomic particles i.e., electron, proton and neutrons.

Particle/ Species

Electron

Proton

Neutron

Nucleus

Discovery by

J. J. Thomson (1869)

Goldstein (1886)

Chadwick (1932)

Rutherford (1911)

Nature of charge

Negative

Positive

Neutral

Positive

Absolute charge

-1.6×10-19 C

1.6×10-19 C

0

Relative Charge

-1

+1

0

Mass

9.11 ×10-31 kg

1.672614×10-27 kg

1.67492×10-27 kg

Mass in u

0.00054u

1.00727u

1.00876u

Approx mass in u

0

1

1

Cathode ray discharge tube experiment

Electrons were discovered using cathode ray discharge tube experiment.

A cathode ray discharge tube made of glass is taken with two electrodes. At very low pressure and high voltage, current starts flowing through a stream of particles moving in the tube from cathode to anode. These rays were called cathode rays. When a perforated anode was taken, the cathode rays struck the other end of the glass tube at the fluorescent screen and a bright spot on the screen was observed.

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In presence of electrical or magnetic field, behaviour of cathode rays is similar to that shown by negatively charged particles.

Results of Cathode ray experiment

  1. The cathode rays start from cathode and move towards the anode.

  2. 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.

  3. In the absence of electrical or magnetic field, these rays travel in straight lines

  4. In the presence of electrical or magnetic field, the behaviour of cathode rays is similar to that expected from negatively charged particles, suggesting that the cathode rays consist of negatively charged particles, called electrons.

  5. 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. Thus, we can conclude that electrons are basic constituent of all the atoms.

Charge to Mass (e/m) Ratio of Electron

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Thomson measured the ratio of electrical charge (e) to the mass of electron (me) by using cathode ray tube and applying electrical and magnetic field perpendicular to each other as well as to the path of electrons Deviation of the particles from their path in the presence of electrical or magnetic field depends upon,

  1. The magnitude of the negative charge on the particle, greater the magnitude of the charge on the particle, greater is the interaction with the electric or magnetic field and thus greater is the deflection.

  2. The mass of the particle - lighter the particle, greater the deflection.

  3. The strength of the electrical or magnetic field - stronger the electrical or magnetic field, greater will be the deflection.

Where me is the mass of the electron in kg and e is the magnitude of the charge on the electron in coulomb (C).

Millikan’s Oil Drop Method and Charge on the Electron

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Oil droplets in the form of mist, produced by the atomiser, were allowed to enter through a tiny hole in the upper plate of electrical condenser. The air inside the chamber was ionized by passing a beam of X-rays through it. The electrical charge on these oil droplets was acquired by collisions with gaseous ions.

The fall of these charged oil droplets can be retarded, accelerated or made stationary depending upon the charge on the droplets and the polarity and strength of the voltage applied to the plate.

In the chamber, the forces acting on oil drop are gravitational force, electrostatic force due to electrical field and a viscous drag force when the oil drop is moving.

It was concluded that the magnitude of electrical charge, q, on the droplets is always an integral multiple of the electrical charge, e, that is, q = n e, where n = 1, 2, 3... .

He found the charge on the electron to be e = –1.6 × 10–19 C.

The mass of the electron (me) was determined by combining these results with Thomson’s value of e/me ratio.

me=ee/me= 9.1094×1031 kg 

Canal Rays and Discovery of Protons

Electrical discharge carried out in the modified cathode ray tube, containing gas, led to the discovery of particles carrying positive charge, also known as canal rays.

Characteristics of positively charged particles

  1. The positively charged particles depend upon the nature of gas present in the cathode ray tube. These are simply the positively charged gaseous ions.

  2. The charge to mass ratio of the particles is found to depend on the gas from which these originate.

  3. Some of the positively charged particles carry a multiple of the fundamental unit of electrical charge.

  4. 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.

Neutrons were discovered by James Chadwick by bombarding a thin sheet of beryllium by α- particles. They are electrically neutral particles having a mass slightly greater than that of the protons.

Some Terms Related to Atom

Nucleons: The protons and neutrons present in the nucleus are collectively known as nucleons.

Atomic number (Z) = number of protons in the nucleus

= number of electrons in a neutral atom

The total number of nucleons is termed as mass number (A) of the atom.

Mass number (A) = number of protons (Z) + number of neutrons (n)

Isotopes: The atoms of the same element having the same atomic number but different mass number, e.g.,

H (Protium)11, D (Deutrim)12, T (Tritrium)13.

Chemical properties of atoms are controlled by the number of electrons, which are determined by the number of protons in the nucleus. Therefore, all the isotopes of a given element show same chemical behaviour.

Isobars: Isobars are the atoms of different elements having the same mass number but different atomic number, e.g.,

Ar (Argon)1840, Ca 2040 (Calcium).

Isotones: Isotones are the atoms of different elements having the same number of neutrons, e.g.,

C 614, N 715, etc.

Isoelectronic species: The atoms or ions, which have the same number of electrons. For example, Ne, Na+, F, Al3+, Mg2+, etc.

Atomic Models

Atomic models are structures showing arrangement of different sub atomic particles within the atom.

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Thomson’s Model or Water Melon Model or Plum Pudding Model

This model was proposed by Joseph James Thomson in 1897. This model is also known as apple pie model.

Merits of Thomson’s Model -Thomson’s model explained that atoms are electrically neutral.

Shortcomings - It could not explain the results of α (alpha) particle scattering experiment carried out by Rutherford.

Rutherford’s Model or Planetary Model

This model was given by Ernest Rutherford in 1911, based on his α particle scattering experiment

Alpha(α) particle scattering experiment

In this experiment, fast moving alpha (α)-particles were made to strike a thin gold foil.

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Observations

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Conclusions

  1. Most of the space inside the atom is empty because most of the α-particles passed through the gold foil without getting deflected.

  2. Very few particles were deflected from their path, indicating that the positive charge of the atom occupies very little space.

  3. A very small fraction of α-particles were deflected by 1800, indicating that all the positive charge and mass of the gold atom were concentrated in a very small volume within the atom.

Rutherford’s Model

  1. All the positively charged particles are present in a small space in the centre of the atom. This small space is called nucleus.

  2. Electrons (negatively charged) revolve around the nucleus in circular orbits with a high speed.

  3. The size of the nucleus is very small compared to the size of atom.

Drawbacks of Rutherford’s Model

  1. It is possible to have infinite number of orbits. In practice it is not the case.

  2. The electron moving with acceleration must continuously lose energy and eventually will fall into the nucleus, which is contrary to the fact.

  3. Rutherford model says nothing about the electronic structure of atoms i.e., how the electrons are distributed around the nucleus and what are the energies of these electrons.

Developments Leading to the Bohr’s Model of Atom

  1. Dual character of the electromagnetic radiation which means that radiations possess both wave like and particle like properties,

  2. Experimental results regarding atomic spectra which can be explained only by assuming quantized electronic energy levels in atoms.
  1. Wave Nature of Electromagnetic Radiation

    When electrically charged particle moves under acceleration, alternating electrical and magnetic fields are produced and transmitted. These fields are transmitted in the forms of waves called electromagnetic waves or electromagnetic radiation.

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Properties of electromagnetic radiations

  1. The oscillating electric and magnetic fields produced by oscillating charged particles are perpendicular to each other and both are perpendicular to the direction of propagation of the wave.

  2. They do not require a medium to travel. That means they can even travel in vacuum.

  3. Different types of electromagnetic radiations differ from one another in wavelength (or frequency). Collection of electromagnetic waves over a range of wavelengths (frequencies), is called electromagnetic spectrum. Example, Radio waves, light, xrays, gamma rays etc.

  4. Wavelength - The distance between two neighbouring crests or troughs of wave. It is denoted by λ.

  5. Frequency (ν) - The number of waves which pass through a particular point in one second. Or the number of oscillations per second.

  6. Velocity (c) - The distance travelled by a wave in one second. In vacuum all types of electromagnetic radiations travel with the same velocity. Its value is 3 ×108 m sec-1. This is also the velocity of light, c.

    c = νλ


  7. Wave number- Wave number, ν̅ is defined as the number of wavelengths per unit length. The SI unit is m–1. However commonly used unit is cm–1.

ν̅=1λ

Particle Nature of Electromagnetic Radiation: Planck’s Quantum Theory

  1. Some of the observations which could not be explained with the help of even the electromagnetic theory (wave theory)

  2. The nature of emission of radiation from hot bodies (black body radiation)

  3. Ejection of electrons from metal surface when radiation strikes it (photoelectric effect)

  4. Variation of heat capacity of solids as a function of temperature

  5. Line spectra of atoms with special reference to hydrogen

Quantum Theory

Photoelectric effect

The phenomenon of ejection of electrons from the surface of metal, when light of suitable frequency strikes it, is called photoelectric effect. The ejected electrons are called photoelectrons.

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Observations

Dual Behavior of Electromagnetic Radiation

The light possesses both particle and wave like properties, i.e., light has dual behavior. Whenever radiation interacts with matter, it displays particle like properties (for example, black body radiation and photoelectric effect). Wave like properties are exhibited when it propagates (for example, interference and diffraction).

Spectrum

When a white light is passed through a prism, it splits into a series of coloured bands known as spectrum.

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Rydberg’s Formula - Wave number of emitted radiation, when an electron falls from energy level n2 to n1, is given by

ν̅=109,677 1n12-1n22

Where n1 = 1, 2........, n2 = n1 + 1, n1 + 2......,

and R = Rydberg’s constant = 109677 cm-1

Spectral Lines for Atomic Hydrogen

Series

n1

n2

Spectral Region

Layman

1

2, 3, …

Ultraviolet

Balmer

2

3, 4, …

Visible

Paschen

3

4, 5, …

Infrared

Brackett

4

5, 6, …

Infrared

Pfund

5

6, 7, …

Infrared

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Bohr’s Atomic Model

This model was given by Neils Bohr in 1913.

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According to this model,

  1. The electron in the hydrogen atom can move around the nucleus in a circular path of fixed radius and energy, known as orbits, stationary states or allowed energy states. These orbits are arranged concentrically around the nucleus.

  2. The energy of an electron in the orbit does not change with time. While revolving in discrete orbits the electrons do not radiate energy.
    • The electron will move from a lower stationary state to a higher stationary state when required amount of energy is absorbed by the electron.

    • Energy is emitted when electron moves from higher stationary state to lower stationary state.

    • The energy change does not take place in a continuous manner.

  3. The frequency of radiation absorbed or emitted when transition occurs between two stationary states that differ in energy by ΔE, is given by ,

    ν=ΔEh=E2- E1h (Bohrs frequency rule)

    Where E1 and E2 are the energies of the lower and higher allowed energy states respectively..


  4. The angular momentum of an electron in a given stationary state can be expressed as

    mev r=nh2π,  where n=1,2,3 

Explanation of Line Spectrum for Hydrogen

Limitations of Bohr’s Model of Atom

  1. It fails to account for the finer details (like doublet, that is two closely spaced lines) of the hydrogen atom spectrum observed by using sophisticated spectroscopic techniques.

    [Doublets occur due to presence of two electrons with different spin orientation in each orbital]


  2. This model is also unable to explain the spectrum of atoms other than hydrogen, for example, helium atom which possesses only two electrons.

  3. Bohr’s theory was also unable to explain the splitting of spectral lines in the presence of magnetic field (Zeeman effect) or an electric field (Stark effect).

  4. It could not explain the ability of atoms to form molecules by chemical bonds.

Dual Behavior of Matter - de Broglie’s wavelength

de Broglie proposed that matter, like radiation, should also exhibit dual behaviour i.e., both particle and wavelike properties. This means that just as the photon has momentum as well as wavelength, electrons should also have momentum as well as wavelength. de Broglie’s wavelength (i.e., the wavelength of matter wave associated with a particle of mass m) is given by,

λ=hmv=hp

Heisenberg’s Uncertainty Principle:

It states that it is impossible to determine simultaneously, the exact position and exact momentum (or velocity) of an electron. The product of their uncertainties is always equal to or greater than h4π.

Mathematically,

Δx × Δp h4π, where  Δx=uncertainity in position and Δp=  uncertainity in momentum  

Significance of Heisenberg Uncertainty Principle

  • Heisenberg’s uncertainty principle rules out the existence of definite paths or trajectories of electrons and other similar particles.

  • The effect of Heisenberg Uncertainty Principle is significant only for motion of microscopic objects and is negligible for that of macroscopic objects.

  • The precise statements of the position and momentum of electrons have to be replaced by the statements of probability, that the electron has at a given position and momentum.

Failure of Bohr’s model

  • It ignores the dual behavior of matter.

  • It contradicts Heisenberg’s uncertainty principle.

Classical and Quantum Mechanics

It is based on Newton’s laws of motion. It successfully describes the motion of macroscopic particles but fails in the case of microscopic particles.

Reason: Classical mechanics ignores the concept of dual behaviour of matter especially for sub-atomic particles and the Heisenberg’s uncertainty principle.

Quantum mechanics

It is a theoretical science that deals with the study of the motions of the microscopic objects that have both observable wave like and particle like properties.

Schrödinger’s equation

Quantum mechanics is based on a fundamental equation which is called Schrodinger’s equation.

For a system (such as an atom or a molecule whose energy does not change with time) the Schrödinger equation is written as,

Ĥ ψ = E ψ,

Where Ĥ is a mathematical operator called Hamiltonian, E is the energy of the system and ψ is the wave function.

When Schrödinger equation is solved for hydrogen atom, the solution gives the possible energy levels the electron can occupy and the corresponding wave function(s) of the electron associated with each energy level.

These quantized energy states and corresponding wave functions are characterized by a set of three quantum numbers

  • principal quantum number n,

  • azimuthal quantum number l and

  • magnetic quantum number ml.

The wave function is a mathematical function whose value depends upon the coordinates of the electron in the atom and does not carry any physical meaning. The wave functions of hydrogen or hydrogen like species with one electron are called atomic orbitals.

Probability density - ψ gives the amplitude of wave. The value of ψ has no physical significance.

The probability of finding an electron at a point within an atom is proportional to the |ψ|2 at that point. It is called probability density.

Schrödinger equation for Multi-electron Atoms

The Schrödinger equation cannot be solved exactly for a multi-electron atom and approximate methods are used. In fact there is very little difference between the orbitals in multi-electron atoms and that in hydrogen.

However the orbitals in multi-electron atoms are contracted due to increased nuclear charge.

Quantum numbers

There are a set of four quantum numbers which specify the energy, size, shape and orientation of electron in an orbital.

To specify an orbital only three quantum numbers are required while to specify an electron all four quantum numbers are required.

Value of n = 1 2 3 4 ...

Shell Name = K L M N ...

Principal quantum number (n) identifies the shell, determines the size and to large extent the energy of the orbital. The number of allowed orbitals in nth shell is given by ‘n2

Azimuthal quantum number ‘lis also known as orbital angular momentum or subsidiary quantum number. It defines the subshell, three dimensional shape of the orbital and energy in case multi-electron atoms and orbital angular momentum.

For a given value of n, l can have n values ranging from 0 to n – 1, that is, for a given value of n, the possible value of l are : l = 0, 1, 2, .........(n–1).

Value of l

0

1

2

3

4

5

Notation for subshell

s

p

d

f

g

h

The orbital angular momentum is given by

ll+12 h2π

Magnetic quantum number or Magnetic orbital quantum number (ml) gives information about the spatial orientation of the orbital with respect to standard set of co-ordinate axis. For any sub-shell (defined by ‘l’ value), 2l+1 values of ml are possible, ranging from –l to l, that is, – l, – (l –1), – (l–2)... 0,1...(l – 2), (l–1), l.

Electron spin quantum number (ms) refers to orientation of the spin of the electron within an orbital. It can have two values +½ and -½. +½ identifies the clockwise spin and -½ identifies the anti-clockwise spin.

Shell (n)

l = 0 to n-1

Subshell notation

Energy of orbital

ml = -l to +l

Orbital notation

Total orbitals

Total electrons

1

0

1s

1

0

1s

1

1

2

2

0

2s

2

0

2s

1

4

8

1

2p

3

-1, 0, 1

2px, 2py, 2pz

3

3

0

3s

3

0

3s

1

9

18

1

3p

4

-1, 0, 1

3px, 3py, 3pz

3

2

3d

5

-2 to +2

3dxy, 3dxz, 3dx2-y2, 3dyz, 3dz2

5

4

0

4s

4

0

4s

1

16

32

1

4p

5

-1, 0, 1

4px, 4py, 4pz

3

2

4d

6

-2 to +2

4dxy, 4dxz, 4dx2-y2, 4dyz, 4dz2

5

3

4f

7

-3 to +3

7

Summary of Quantum Mechanical Model of Atom

Quantum mechanical model of atom is the picture of the structure of the atom, which emerges from the application of the Schrödinger equation to atoms.

  1. The energy of electrons in atoms is quantized (i.e., can only have certain specific values), for example when electrons are bound to the nucleus in atoms.

  2. Existence of quantized electronic energy levels is due to the wave like properties of electrons and allowed solutions of Schrödinger wave equation.

  3. Both the exact position and exact velocity of an electron in an atom cannot be determined simultaneously (Heisenberg uncertainty principle). The path of an electron in an atom therefore, can never be determined or known accurately. That is why we talk of only probability of finding the electron at different points in an atom.

  4. An atomic orbital is the wave function ψ for an electron in an atom. Whenever an electron is described by a wave function, we say that the electron occupies that orbital. Since many such wave functions are possible for an electron, there are many atomic orbitals in an atom.

    In each orbital, the electron has a definite energy. An orbital cannot contain more than two electrons.

    In a multi-electron atom, the electrons are filled in various orbitals in the order of increasing energy. For each electron of a multi-electron atom, there shall, therefore, be an orbital wave function characteristic of the orbital it occupies. All the information about the electron in an atom is stored in its orbital wave function ψ and quantum mechanics makes it possible to extract this information out of ψ.

  1. The probability of finding an electron at a point within an atom is proportional to the square of the orbital wave function i.e., |ψ|2 at that point. |ψ|2 is known as probability density and is always positive. From the value of |ψ|2 at different points within an atom, it is possible to predict the region around the nucleus where electron will most probably be found.

Shapes of Atomic Orbitals

The orbital wave function or ψ for an electron in an atom has no physical meaning. It is simply a mathematical function of the coordinates of the electron. However the square of the wave function (i.e., ψ2) at a point gives the probability density of the electron at that point.

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s orbital is spherical

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px orbital is along x-axis, py orbital is along y-axis and pz orbitalis along z-axis

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dxy orbital is in xy plane (each lobe bisecting the angle between x and y-axes)

dzx orbital is in zx plane (each lobe bisecting the angle between z and x-axes)

dyz orbital is in yz plane (each lobe bisecting the angle between y and z-axes)

lobes of x2-y2 orbital are along x and y-axes.

Two lobes of z2 orbital are along z-axis with ring of high electron density in the xy plane. This is basically combination of two solutions of the Schrödinger equation.

Nodal surfaces or Nodes

The region where this probability density function reduces to zero is called nodal surfaces or simply nodes.

Total number of nodes = n – 1

Radial nodes: Radial nodes occur when the probability density of wave functionfor the electron is zero on a spherical surface of a particular radius.

Number of radial nodes = n – l – 1

Angular nodes: Angular nodes occur when the probability density function for the electron is zero along the directions specified by a particular angle.

Number of angular nodes = l

  • Explain the probability density and nodes for 1s, 2s and other orbitals.

Boundary Surface Diagram

Boundary surface diagram is a boundary surface or a contour surface drawn in a space for an orbital on which the value of probability density |ψ|2 is constant. The boundary surface diagram of constant probability density is considered as a good and acceptable approximation of shape of orbital if the boundary surface encloses the region or volume with probability density of more than 90%.

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Energies of Orbitals

  • In hydrogen and hydrogen like species - The energy of an electron in a hydrogen atom is determined solely by the principal quantum number, n. Thus the energy of the orbitals are in the order 1s < 2s = 2p < 3s = 3p = 3d <4s = 4p = 4d = 4f.

  • In Multi-electron Species - The energy of an electron in a multi-electron atom depends on n + l.

  • Energies of the orbitals in the same subshell decrease with increase in the atomic number (Zeff). For example, energy of 2s orbital of hydrogen atom is greater than that of 2s orbital of lithium and that of lithium is greater than that of sodium and so on, that is, E2s(H) > E2s(Li) > E2s(Na) > E2s(K).

  • Ground state - the energy state for an electron corresponding to the most stable condition in which an electron is most strongly held by the nucleus The 1s in a hydrogen atom is the ground state. 2s, 2p etc in hydrogen atom are excited states.

  • Degenerate orbitals - Orbitals having the same energy are called degenerate orbitals. 2s and 2p are degenerate orbitals in hydrogen.

Shielding effect or screening effect:

  • Due to the presence of electrons in the inner shells, the electron in the outer shell will not experience the full positive charge on the nucleus. This is called shielding of the outer shell electrons from the nucleus by the inner shell electrons or screening effect.

  • Due to the screening effect, the net positive charge experienced by the electron of outer shell is less than the total charge on the nucleus and is known as effective nuclear charge, Zeff. Effective nuclear charge experienced by the orbital decreases with increase of azimuthal quantum number (l).

  • Screens effect of orbitals s > p > d > f….

Filling of Orbitals in Atom

The filling of electrons into the orbitals of different atoms takes place according to,

  1. The Aufbau principle

  2. Pauli’s exclusion principle,

  3. The Hund’s rule of maximum multiplicity and

  4. The relative energies of the orbitals.
The Aufbau Principle

In the ground state of the atoms, the orbitals are filled in order of their increasing energies.

Orbitals with lower value of (n + l) have lower energy. If two orbitals have the same value of (n + l) then orbital with lower value of n will have lower energy.

The order in which the orbitals are filled is,

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 4f, 5d, 6p, 7s...

Pauli Exclusion Principle:

“No two electrons in an atom can have the same set of four quantum numbers.”

or

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“Only two electrons may exist in the same orbital and these electrons must have opposite spin.”

  • The maximum number of electrons in any shell can be 2n2.

Hund’s rule of maximum multiplicity

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.

Electronic Configuration of Atoms

The distribution of electrons into orbitals of an atom is called its electronic configuration.

The electronic configuration of different atoms can be represented in two ways,

  1. sa pb dc...... notation. In this notationwe can also write the electronic configuration as [X] sa pb dc, where X is the noble gas preceding the element in the periodic table, eg. electronic configuration of Na can be written as,

    1s2 2s2 2p6 3s1 or [Ne] 3s1

(ii) Orbital diagram, each orbital of the subshell is represented by a box and the electron is represented by an arrow (↑) a positive spin or an arrow (↓) a negative spin.

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  • Take examples and fill the orbitals.

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Stability of Completely Filled and Half Filled Subshells

The completely filled and completely half filled sub-shells are stable compared to other configurations.

The ground state electronic configuration of the atom of an element always corresponds to the state of the lowest total electronic energy.

The electronic configurations of most of the atoms follow the basic rules given above, however, in certain elements such as Cu, or Cr, where the two subshells (4s and 3d) differ slightly in their energies, an electron shifts from a subshell of lower energy (4s) to a subshell of higher energy (3d), provided such a shift results in all orbitals of the subshell of higher energy getting either completely filled or half filled.

The valence electronic configurations of Cr and Cu, therefore, are 3d5 4s1 and 3d10 4s1 respectively and not 3d4 4s2 and 3d9 4s2.

  • Check the other starred elements in the list in the book.

Causes of Stability of Completely Filled and Half Filled Sub-shells

The extra stability of half-filled and completely filled subshell is due to:

  1. Symmetrical distribution of electrons (Relatively small shielding)

    Symmetry leads to stability. The completely filled or half filled subshells have symmetricaldistribution of electrons in them and are therefore more stable. Electrons in the same subshell (here 3d) have equal energy but different spatial distribution. Consequently, their shielding of one another is relatively small and the electrons are more strongly attracted by the nucleus.


  2. Larger exchange energy

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The configuration is more stable whenever two or more electrons with the same spin are present in the degenerate orbitals of a subshell.

These electrons tend to exchange their positions and the energy released due to this exchange is called exchange energy.

The number of exchanges that can take place is, maximum when the subshell is either half filled or completely filled (see figure).

As a result the exchange energy is, maximum and so is the stability.

The exchange energy is the basis of Hund’s rule that electrons which enter orbitals of equal energy have parallel spins as far as possible.

  1. Smaller coulombic repulsion energy,

    The coloumbic force is less when only one electron is present in one orbital compared to when two electrons are present in the same orbital. As a result the configuration is more stable.