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Subtopics Of Class 11 Chemistry Chapter 2 Structure Of The Atom
- Sub-atomic Particles
- Discovery Of Electron
- Charge To Mass Ratio Of Electron
- Charge On The Electron
- Discovery Of Protons And Neutrons
- Atomic Models
- Thomson Model Of Atom
- Rutherford’s Nuclear Model Of Atom
- Atomic Number And Mass Number
- Isobars And Isotopes
- Drawbacks Of Rutherford Model
- Developments Leading To The Bohr’s Model Of Atom
- Wave Nature Of Electromagnetic Radiation
- Particle Nature Of Electromagnetic Radiation: Planck’s Quantum Theory
- Evidence For The Quantized* Electronic Energy Levels: Atomic Spectra
- Bohr’s Model For Hydrogen Atom
- Explanation Of Line Spectrum Of Hydrogen
- Limitations Of Bohr’s Model
- Towards Quantum Mechanical Model Of The Atom
- Dual Behaviour Of Matter
- Heisenberg’s Uncertainty Principle
- Quantum Mechanical Model Of Atom
- Orbitals And Quantum Numbers
- Shapes Of Atomic Orbitals
- Energies Of Orbitals
- Filling Of Orbitals In Atom
- Electronic Configuration Of Atoms
- Stability Of Completely Filled And Half Filled Subshells.
Structure of Atom Class 11 Notes Chemistry Chapter 2
• Discovery of Electron—Discharge Tube Experiment
In 1879, William Crooks studied the conduction of electricity through gases at low pressure. He performed the experiment in a discharge tube which is a cylindrical hard glass tube about 60 cm in length. It is sealed at both the ends and fitted with two metal electrodes as shown in Fig. 2.1.
The electrical discharge through the gases could be observed only at very low pressures and at very high voltages.
The pressure of different gases could be adjusted by evacuation. When sufficiently high voltage is applied across the electrodes, current starts flowing through a stream of particles moving in the tube from the negative electrode (cathode) to the positive electrode (anode). These were called cathode rays or cathode ray particles.
• Properties of Cathode Rays
(i) Cathode rays travel in straight line.
(ii) Cathode rays start from cathode and move towards the anode.
(iii) 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.
(iv) Cathode rays consist of negatively charged particles. When electric field is applied on the cathode rays with the help of a pair of metal plates, these are found to be deflected towards the positive plate indicating the presence of negative charge.
(v) The characteristics of cathode rays do not depend upon the material of electrodes and the nature of gas present in the cathode ray’tube.
• Determination of Charge/Mass (elm) Ratio for Electrons
J. J. Thomson for the first time experimentally determined charge/mass ratio called elm ratio for the electrons. For this, he subjected the beam of electrons released in the discharge tube as cathode rays to influence the electric and magnetic fields. These were acting perpendicular to one another as well as to the path followed by electrons.
According to Thomson, the amount of deviation of the particles from their path in presence of electrical and magnetic field depends upon following factors:
(i) 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.
(ii) The mass of the particle — lighter the particle, greater the deflection.
(iii) The deflection of electrons from their original path increases with the increase in the voltage across the electrodes or strength of the magnetic field.
By carrying out accurate measurements on the amount of deflections observed by the electrons on the electric field strength or magnetic field strength, Thomson was able to determine the value of
e/me = 1.758820 x 1011 C kg-1 where me = Mass of the electron in kg
e = magnitude of charge on the electron in coulomb (C).
• Charge on the Electron
R.A. Millikan devised a method known as oil drop experiment to determine the charge on the electrons.
• Discovery of Proton—Anode Rays
In 1886, Goldstein modified the discharge tube by using a perforated cathode. On reducing the pressure, he observed a new type of luminous rays passing through the holes or perforations of the cathode and moving in a direction opposite to the cathode rays. These rays were named as positive rays or anode rays or as canal rays. Anode rays are not emitted from the anode but from a space between anode and cathode.
• Properties of Anode Rays
(i) The value of positive charge (e) on the particles constituting anode rays depends upon the nature of the gas in the discharge tube.
(ii) The charge to mass ratio of the particles is found to depend on the gas from which these originate.
(iii) Some of the positively charged particles carry a multiple of the fundamental unit of electrical charge.
(iv) The behaviour of these particles in the magnetic or electric 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.
• Thomson Model of Atom
(i) J. J. Thomson proposed that an atom may be regarded as a sphere of approximate radius 1CT8 cm carrying positive charge due to protons and in which negatively charged electrons are embedded.
(ii) In this model, the atom is visualized as a pudding or cake of positive charge with electrons embedded into it.
(iii) 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 a-particle Scattering Experiment
Rutherford in 1911, performed some scattering experiments in which he bombarded thin foils of metals like gold, silver, platinum or copper with a beam of fast moving a-particles. The thin gold foil had a circular fluorescent zinc sulphide screen around it. Whenever a-particles struck the screen, a tiny flash of light was produced at that point.
From these experiments, he made the following observations:
(i) Most of the a-particles passed through the foil without undergoing any deflection,
(ii) A few a-particles underwent deflection through small angles.
(iii) Very few mere deflected back i.e., through an angle of nearly 180°.
From these observations, Rutherford drew the following conclusions:
(i) Since most of the a-particles passed through the foil without undergoing any deflection, there must be sufficient empty space within the atom.
(ii) A small fraction of a-particles was deflected by small angles. The positive charge has to be concentrated in a very small volume that repelled and deflected a few positively charged a-particles. This very small portion of the atom was called nucleus.
(iii) The volume of nucleus is very small as compared to total volume of atom.
• Rutherford’s Nuclear Model of an Atom
(i) 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.
(ii) The nucleus is surrounded by electrons that move around the nucleus with a very high speed in circular paths called orbits.
(iii) Electrons and nucleus are held together by electrostatic forces of attraction.
• Atomic Number
The number of protons present in the nucleus is equal to the atomic number (z). For example, the number of protons in the hydrogen nucleus is 1, in sodium atom it is 11, therefore, their atomic numbers are 1 and 11. In order to keep the electrical neutrality, the number of electrons in an atom is equal to the number of protons (atomic number, z). For example, number of electrons in hydrogen atom and sodium atom are 1 and 11 respectively.
Atomic Number (z) = Number of protons in the nucleus of an atom.
= Number of electrons in a neutral atom.
• Mass Number
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).
Atoms with identical atomic number but different atomic mass number are known as Isotopes.
Isotopes of Hydrogen:
These three isotopes are shown in the figure below:
Characteristics of Isotopes
(i) Since the isotopes of an element have the same atomic number, but different mass number, the nuclei of isotopes contain the same number of protons, but different number of neutrons.
(ii) Since, the isotopes differ in their atomic masses, all the properties of the isotopes depending upon the mass are different.
(iii) Since, the chemical properties are mainly determined by the number of protons in the nucleus, and the number of electrons in the atom, the different isotopes of an element exhibit similar chemical properties. For example, all the isotopes of carbon on burning give carbon dioxide.
• Drawbacks of Rutherford Model
(i) When a body is moving in an orbit, it achieves acceleration. Thus, an electron moving around nucleus in an orbit is under acceleration.
According to Maxwell’s electromagnetic theory, charged particles when accelerated must emit electromagnetic radiations. Therefore, an electron in an orbit will emit radiations, the energy carried by radiation comes from electronic motion. Its path will become closer to nucleus and ultimately should spiral into nucleus within . 10-8 s. 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.
(ii) Rutherford’s model does not give any idea about distribution of electrons around the nucleus and about their energies.
• Developments Leading to the Bohr’s Model of Atom
Two developments played a major role in the formulation of Bohr’s model of atom. These were:
(i) Dual character of the electromagnetic radiation which means that radiations possess both wave like and particle like properties.
(ii) Experimental results regarding atomic spectra which can be explained only by assuming quantized electronic energy levels in atoms.
• Nature of Electromagnetic Radiation (Electromagnetic Wave Theory)
This theory was put forward by James Clark Maxwell in 1864. The main points of this theory are as follows:
(i) 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.
(ii) The radiations consist of electric and magnetic fields oscillating perpendicular to each other and both perpendicular to the direction of propagation of the radiation.
(iii) The radiations possess wave character and travel with the velocity of light 3 x 108 m/sec.
(iv) 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
Wavelength: It is defined as the distance between any two consecutive crests or troughs. It is represented by X and its S.I. unit is metre.
Frequency: Frequency of a wave is defined as the number of waves passing through a point in one second. It is represented by v (nu) and is expressed in Hertz (Hz).
1 Hz = 1 cycle/sec.
Velocity: 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: 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.
Wave Number: It is defined as the number of waves present in 1 metre length. Evidently it will be equal to the reciprocal of the wavelength. It is represented by bar v (read as nu bar).
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:
(i) The phenomenon of black body radiation.
(ii) The photoelectric effect.
(iii) The variation of heat capacity of solids as a function of temperature.
(iv) The line spectra of atoms with reference to hydrogen.
• 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.
At a given temperature, intensity of radiation emitted increases with decrease of wavelength, reaches a maximum value at a given wavelength and then starts decreasing with further decrease of wavelength as shown in Fig 2.6.
• 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: ,
(i) The radiant energy emitted or absorbed in the form of small packets of energy. Each such packets of energy is called a quantum.
(ii) The energy of each quantum is directly proportional to the frequency of the radiation
where h is a proportionality constant, called Planck’s constant. Its value is equal to 6.626 x 10-34 Jsec.
• 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
(i) 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.
(ii) 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.
(iii) 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.
Explanation of Photoelectric Effect
Einstein in (1905) was able to give an explanation of the different points of the photoelectric effect using Planck’s quantum theory as under:
(i) Photoelectrons are ejected only when the incident light has a certain minimum frequency (threshold frequency v0)
(ii) If the frequency of the incident light (v) is more than the threshold frequency (v0), the excess energy (hv – hv0) is imparted to the electron as kinetic energy.
K.E. of the ejected electron
energy of the emitted electron.
(iii) On increasing the intensity of light, more electrons are ejected but the energies of the electrons are not altered.
• Dual Behaviour of Electromagnetic Radiation
From the study of behaviour of light, scientists came to the conclusion that light and other electromagnetic radiations have dual nature. These are wave nature as well as particle nature. Whenever radiation interacts with matter, it displays particle like properties in contrast to the wavelike 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.
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 is noticed 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:
(i) from sun or glowing electric bulb.
(ii) by passing electric discharge through a gas at low pressure.
(iii) by heating a substance to high temperature.
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. For example, sodium or its salts emit yellow light while potassium or its salts give out violet light.
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 otherwise 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.
Dark lines appear exactly at the same positions where the lines in the emission spectra appear.
• 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.
On the basis of experimental observations, Johannes Rydberg noted that all series of lines in the hydrogen spectrum could be described by the following expression:
Rydberg in 1890, and has given a simple theoretical equation for the calculation of wavelengths and wave numbers of the spectral lines in different series of hydrogen spectrum. The equation is known as Rydberg formula (or equation).
This relation is valid for hydrogen atom only. For other species,
where Z is the atomic number of the species.
Here RH = constant, called Rydberg constant for hydrogen and n1 , n2 are integers (n2 > n1)
For any particular series, the value of n1 is constant while that of n2 changes. For example,
For Lyman series, n1= 1, n2= 2, 3, 4, 5………..
For Balmer series, n1 = 2, n2 = 3, 4, 5, 6………..
For Paschen series, n1= 3, n2 = 4, 5, 6, 7………..
For Brackett series,n1 = 4, n2 = 5, 6, 7, 8………..
For Pjund series, n1 =5, n2 = 6, 7, 8, 9………..
Thus, by substituting the values of n1 and n2 in the above equation, wavelengths and wave number of different spectral lines can be calculated. When n1 = 2, the expression given above is called Balmer’s formula.
• 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:
(i) In an atom, the electrons revolve around the nucleus in certain definite circular paths called orbits.
(ii) 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………..
(iii) 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.
(iv) As long as electron is present in a particular orbit, it neither absorbs nor loses energy and its energy, therefore, remains constant.
(v) When energy is supplied to an electron, it absorbs energy only in fixed amounts as quanta and jumps to higher energy state away from the nucleus 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
1. Bohr’s theory has explained the stability of an atom.
2. Bohr’s theory has helped in calculating the energy of electron in hydrogen atom and one electron species. The mathematical expression for the energy in the nth orbit is,
3. Bohr’s theory has explained the atomic spectrum of hydrogen atom.
• Limitations of Bohr’s Model
(i) The theory could not explain the atomic spectra of the atoms containing more than one electron or multielectron atoms.
(ii) Bohr7s theory failed to explain the fine structure of the spectral lines.
(iii) Bohr’s theory could not offer any satisfactory explanation of Zeeman effect and Stark effect.
(iv) Bohr’s theory failed to explain the ability of atoms to form molecule formed by chemical bonds.
(v) It was not in accordance with the Heisenberg’s uncertainty principle.
• Dual Behaviour of Matter (de Broglie Equation)
de Broglie in 1924, proposed that matter, like radiation, should also exhibit dual behaviour i.e., both particle like and wave like properties. This means that like photons, electrons also have momentum as well as wavelength.
From this analogy, de Broglie gave the following relation between wavelength (λ) and momentum (p) of a material particle.
• Heisenberg’s Uncertainty Principle
It states that, “It is impossible to determine simultaneously, the exact position and exact momentum (or velocity) of an electron”.
• Significance of Uncertainty Principle
(i) It rules out existence of definite paths or trajectories of electrons and other similar particles.
(ii) The effect of Heisenberg’s uncertainty principle is significant only for microscopic objects and is negligible for macroscopic objects.
• Reasons for the Failure of Bohr Model
(i) The wave character of the electron is not considered in Bohr Model.
(ii) According to Bohr Model an orbit is a clearly defined path and this path can completely be defined only if both the position and the velocity of the electron are known exactly at the same time. This is not possible according to the Heisenberg’s uncertainty principle.
• Quantum Mechanical Model of Atom
Quantum mechanics: Quantum mechanics 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.
Important Features of Quantum Mechanical Model of Atom
(i) The energy of electrons in atom is quantized i.e., can only have certain values.
(ii) The existence of quantized electronic energy level is a direct result of the wave like properties of electrons.
(iii) Both the exact position and exact velocity of an electron in an atom cannot be determined simultaneously.
(iv) 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.
(v) 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.
From the value of φ2 at different points within atom, it is possible to predict the region around the nucleus where electron most probably will be found.
• Quantum Numbers
Atomic orbitals can be specified by giving their corresponding energies and angular momentums which are quantized (i.e., they have specific values). The quantized values can be expressed in terms of quantum number. These are used to get complete information about electron i.e., its location, energy, spin etc.
Principal Quantum Number (n)
It is the most important quantum number since it tells the principal energy level or shell to which the electron belongs. It is denoted by the letter V and can have any integral value except zero, i.e., n = 1, 2, 3, 4……….. etc.
The various principal energy shells are also designated by the letters, K, L, M, N, O, P ….. etc. Starting from the nucleus.
The principal quantum number gives us the following information:
(i) It gives the average distance of the electron from the nucleus.
(ii) It completely determines the energy of the electron in hydrogen atom and hydrogen like particles.
(iii) The maximum number of electrons present in any principal shell is given by 2n2 where n is the number of the principal shell.
Azimuthal or Subsidiary or Orbital Angular Quantum Number (l)
It is found that the spectra of the elements contain not only the main lines but there are many fine lines also present. This number helps to explain the fine lines of the spectrum.
The azimuthal quantum number gives the following information:
(i) The number of subshells present in the main shell.
(ii) The angular momentum of the electron present in any subshell.
(in) The relative energies of various subshells.
(iv) The shapes of the various subshells present within the same principal shell.
This quantum number is denoted by the letter T. For a given value of n, it can have any value ranging from 0 to n – 1. For example,
For the 1st shell (k), n = 1, l can have only one value i.e., l = 0 For n = 2, the possible value of l can be 0 and 1.
Subshells corresponding to different values of l are represented by the following symbols:
value of l 0 1 2 3 4 5 ……………..
Notation for subshell s p d f g h ………………..
Magnetic Orbital Quantum Number (m or m1)
The magnetic orbital quantum number determines the number of preferred orientations of the electrons present in a subshell. Since each orientation corresponds to an orbital, therefore, the magnetic orbital quantum number determines the number of orbitals present in any subshell.
The magnetic quantum number is denoted by letter m or ml and for a given value of l, it can have all the values ranging from – l to + l including zero.
Thus, for energy value of l, m has 2l + 1 values.
For l = 0 (s-subshell), ml can have only one value i.e., m1 = 0.
This means that s-subshell has only one orientation in space. In other words, s-subshell has only one orbital called s-orbital.
Spin Quantum Number (S or ms)
This quantum number helps to explain the magnetic properties of the substances. A spinning electron behaves like a micromagnet with a definite magnetic moment. If an orbital contains two electrons, the two magnetic moments oppose and cancel each other.
• Shapes of s-orbitals
s-orbital is present in the s-subshell. For this subshell, l = 0 and ml = 0. Thus, s-orbital with only one orientation has a spherical shape with uniform electron density along all the three axes.
The probability of Is electron is found to be maximum near the nucleus and decreases with the increase in the distance from the nucleus. In 2s electron, the probability is also maximum near the nucleus and decreases to zero probability. The spherical empty shell for 2s electron is called nodal surface or simply node.
• Shapes of p-orbitals
p-orbitals are present in the p-subshell for which l = 1 and m1 can have three possible orientations – 1, 0, + 1.
Thus, there are three orbitals in the p-subshell which are designated as px, py and pz orbitals depending upon the axis along which they are directed. The general shape of a p-orbital is dumb-bell consisting of two portions known as lobes. Moreover, there is a plane passing through the nucleus along which finding of the electron density is almost nil. This is known as nodal plane as shown in the fig.
From the dumb-bell pictures, it is quite obvious that unlike s-orbital, a p-orbital is directional in nature and hence it influences the shapes of the molecules in the formation of which it participates.
• Shapes of d-orbitals
d-orbitals are present in d-subshell for which l = 2 and m[ = -2, -1, 0, +1 and +2. This means that there are five orientations leading to five different orbitals.
• Aufbau Principle
The principle states: In the ground state of the atoms, the orbitals are filled in order of their increasing energies.
In other words, electrons first occupy the lowest energy orbital available to them and enter into higher energy orbitals only after the lower energy orbitals are filled.
The order in which the energies of the orbitals increase and hence the order in which the orbitals are filled is as follows:
Is, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, id, 5p, 6s, if, 3d, 6p, 7s, 5f 6d, 7p
The order may be remembered by using the method given in fig. 2.11.
• Pauli Exclusion Principle
According to this principle, no two electrons in an atom can have the same set of four quantum numbers.
Pauli exclusion principle 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.
• 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.
• Causes of Stability of Completely Filled and Half Filled Subshells
The completely filled and half filled subshells are stable due to the following reasons:
1. Symmetrical distribution of electrons: The completely filled or half filled subshells have symmetrical distribution of electrons in them and are therefore more stable.
2. The stabilizing effect arises whenever two or more electrons with same spin are present in the degenrate orbitals of a subshell. These electrons tend to exchange their positions
and the energy released due to their exchange is called exchange energy. The number of exchanges that can takes place is maximum when the subshell is either half filled or completely filled.
-As a result the exchange energy is maximum and so is the stability.
CBSE Class 11 Chemistry Chapter-2 Important Questions
1 Marks Questions
1.Name the sub – atomic particles of an atom.
Ans: Electron, proton and neutron.
2.Name the scientist who first formulated the atomic structure.
Ans: John Dalton, a British teacher in 1808 first proposed a firm scientific basis known as Dalton’s
3.What is the e/m ratio of an electron?
Ans: According to Thomson’s experiment, e/m ratio for an electron is
4.What is the charge (e) of an electron?
Ans: From Millikan’s experiment, the charge of an electron (e) is
5.(i) What is the mass of a proton?
(ii) What is the charge of a proton?
Ans:(i) The mass of a proton is
(ii) The charge of a proton is
6.(i) What is the mass of a neutron?
(ii) What is the charge of a neutron?
Ans: (i) The mass of a neutron is
(ii) Neutron is electrically neutral i.e. it has no charge as an electron or a proton has.
7.Name the scientist who first gave the atomic model.
Ans: J.J. Thomson, in 1898 first proposed the atomic model called raising-pudding model.
8.What is an isotope?
Ans: Atoms of the same elements having same atomic number but different mass number are called isotopes.
9.What are isobars?
Ans: Atoms of different elements which have same mass number but different atomic nos.
10.What are isotones?
Ans: Atoms of different elements which contains the same number of neutron.
11.What is an atomic number?
Ans: Atoms of different elements which contains the same number of neutron.
12.What is a mass number?
Ans: Maas number of an element is the sum of number of proton and neutron present in the nucleus of an atom.
13.Give the drawbacks of J.J. Thomson’s experiment.
Ans: (i) It could not explain the origin of the spectral lines of hydrogen and other atoms,
(ii) It failed to explain scattering of in Rutherford’s scattering experiment.
14.Why Rutherford’s model could not explain the stability of an atom?
Ans:According to the electromagnetic theory of Maxwell, charged particles when accelerated should emit electromagnetic radiation. Therefore, an electron in an orbit will emit radiation; the orbit will then continue to shrink which does not happen in an atom.
15.Define photoelectric effect.
Ans: It is the phenomenon in which the surface of alkali metals like potassium and calcium emit electrons when a beam of light with high frequency is made to fall on them.
16.How does the intensity of light effect photoelectrons?
Ans: The number of electron ejected and kinetic energy associated with them depend on the brightness of light.
17.What is threshold frequency?
Ans: The minimum frequency below which photo electric effect is not observed is called threshold frequency (vo)
18.Name the scientist who demonstrated photoelectric effect experiment.
Ans: In 1887, H. Hertz demonstrated photo electric effect.
19.What did Einstein explain about photoelectric effect?
Ans: Einstein in 1905 was able to explain the photoelectric effect using Planck’s quantum theory of electromagnetic radiation.
20.Calculate energy of 2mole of photons of radiation whose frequency is
Ans: Energy (E) of one photon = E = hv
Energy of 2 mole of photon
21.States Heisenberg’s Uncertainty Principle.
Ans: It states that It is impossible to determine simultaneously the exact position and exact momentum (or velocity) of an electron.
22..How would the velocity be effected if the position is known?
Ans: If the position of the electron is known with high degree of accuracy (is small), then the velocity of the electron will be uncertain
23.We don’t see a car moving as a wave on the road why?
Ans: According to de Broglie’s relation, the mass of the car is very large and its wavelength or wave character is negligible. Therefore, we do not see a car moving like a wave.
24.Give the de – Broglie’s relation.
Ans: According to de Broglie, every particle in motion is associated with a wavelength and other wave characteristics. He deduced the relation that wavelength of a particle in motion is equal to the Planck’s constant (h) divided by the momentum (p) of the particle.
Where m is the mass, is the velocity after particles
25.Calculate the uncertainty in the velocity of a wagon of mass 4000kg whose position is known accurately of
Ans: Uncertainty in velocity is given by
The uncertainty in the velocity of the wagon is
26.What is the physical significance of up?
Ans: represent probability of finding an electron.
27.Which orbital is non – directional?
Ans:S – orbital is spherically symmetrical i. e it is non – directional.
28.What is the meaning of quantization of energy?
Ans:Quantization of energy means the energy of energy levels can have some specific values and not all the values.
29.Why is energy of 1s electron lower than 2s electron?
Ans.1s electron being close to the nucleus experiences more force of attraction than 2s– electron which is away from the nucleus.
30.What is nodal surface or nodes?
Ans:The region where the probability of finding an electron is zero i. e. Y2 = 0
31.How many spherical nodal surfaces are there in 4s – sub-shell?
Ans:In ns orbital, the number of spherical nodal surfaces are(n – 1), hence is 4s (4 – 1) = 3 nodal surfaces are present.
2 Marks Questions
1.What is the mass (m) of an electron?
Ans: mass of an electron (m) =
So, the mass of an electron is or th of the mass of a hydrogen atom.
2.Which experiment led to the discovery of electrons and how?
Ans:The cathode ray discharge tube experiment performed by J.J. Thomson led to the discovery of negatively charged particles called electron.
A cathode ray tube consists of two thin pieces of metals called electrodes sealed inside a glass tube with sealed ends. The glass tube is attached to a vacuum pump and the pressure inside the tube is reduced to 0.01mm. When fairly high voltage (10, 000V) is applied across the electrodes, invisible rays are emitted from the cathode called cathode rays. Analysis of this rays led to the discovery electrons.
3.Give the main properties of canal ray experiment.
Ans:The canal ray experiment led to the discovery of –
(i)The anode rays, travel in straight line
(ii)They are positively charged as they get deflected towards the –ve end when subjected to an electric and magnetic field.
(iii)They depend upon the nature of gas present in the cathode tube.
(iv)The charge to mass ration (e/m) of the particle is found to depend on the gas from which they originate.
(v)They are also material particles
The analysis of these proportions led to the discovery of positively charged proton.
4.Find out atomic number, mass number, number of electron and neutron in an element?
Ans: The mass no. of
The atomic no. of
No. of proton is = Z – A = 40 – 20 = 20
No. of electron its (A) = 20
No. of proton is (A) = 20
5.Give the main features of Thomson’s Model for an atom.
Ans: J.J. Thomson proposed that an atom consists of a spherical sphere (radius of about 10-10m)in which the positive charges are uniformly distributed the electrons are embedded into it in such a manner so as to give stable electrostatic arrangement.
This model is also called raisin pudding model.
6.What did Rutherford conclude from the observations of scattering experiment?
Ans: Rutherford proposed the nuclear model of an atom as
(i) The positive charge and most of the mass of an atom was concentrated in an extremely small region. He called it nucleus.
(ii) The nucleus is surrounded by electrons that move around the nucleus with a very high speed in orbits.
(iii) Electron and nucleus are held together by electrostatic forces of attraction.
7.What is the relation between kinetic energy and frequency of the photoelectrons?
Ans: Kinetic energy of the ejected electron is proportional to the frequency of the electromagnetic radiation.
8.What transition in the hydrogen spectrum would have the same wavelength as the Balmer transition, n = 4 to n = 2 of He+ spectrum?
Ans: For the Balmer transition, n = 4, to n = 2 in a He+ ion, we can write.
For a hydrogen atom
Equating equation (ii) and (i), we get
This equation gives n1 = 1 and n = 2. Thus the transition n = 2 to n = 1 in hydrogen atom will have same wavelength as transition, n = 4 to n = 2 in He+
9.Spectral lines are regarded as the finger prints of the elements. Why?
Ans: Spectral lines are regarded as the finger prints of the elements because the elements can be identified from these lines. Just like finger prints, the spectral lines of no two elements resemble each other.
10.Why cannot the motion of an electron around the nucleus be determined
Ans: Because there is an uncertainty in the velocity of moving electron around the nucleus (Heisenberg’s Uncertainty Principle).
11.Calculate the uncertainty in the momentum of an electron if it is confined to a linear region of length
Ans: According to uncertainty Principle
12.Give the mathematical expression of uncertainty principle.
Ans:Mathematically, it can be given as
Where is the uncertainty in position and is the uncertainty in momentum (or velocity) of the particle.
13.Which quantum number determines
(i) energy of electron
(ii) Orientation of orbitals.
Ans. (i) Principal quantum number (n), and
(ii) Magnetic quantum number (m).
14.Arrange the electrons represented by the following sets of quantem number in decreasing order of energy.
1. n = 4, l = 0, m = 0, s = +1/2
2. n = 3, l = 1, m = 1, s = -1/2
3. n = 3, l = 2, m = 0, s = +1/2
Ans.(i)Represents 4s orbital
(ii) Represents 3p orbital
(iii)Represents 3d orbital
(iv)Represents 3s orbital
The decreasing order of energy 3d > 4s > 3p > 3s
n = 3, l = 0, m = 0, s = -1/2
3 Marks Questions
1.What designations are given to the orbitals having
(i) n = 2, l = 1 (ii) n = 2, l = 0 (iii) n = 4, l = 3
(iv) n = 4, l = 2 (v) n = 4, l = 1?
Ans. (i) Here, n = 2, and l = 1
Since l = 1 it means a p-orbital, hence the given orbital is designated as 2p.
(ii) Here, n = 2 and l = 0
Since l = 0 means s – orbital, hence the given orbital is 2s.
(iii) Here, n = 4 and l = 3
Since, l = 3 represents f – orbital, hence the given orbital is a 4f orbital.
(iv) Here, n = 4 and l = 2
Since, l = 2 represents d – orbital, hence the given orbital is a 4d – orbital.
(v) n = 4 and l = 1
since, l = 1 means it is a p – orbital, hence the given orbital can be designated as – 4p orbital.
2.Write the electronic configuration of (i) Mn4+, (ii) Fe3+ (iii) Cr2+ and Zn2+ Mention the number of unpaired electrons in each case.
Ans.(i) Mn (z = 25), Mn4+ (z = 21)
The electronic configuration of Mn4+ to Given by
1s2 2s2 2p6 3s2 3p6 3d3
As the outermost shell 3d has 3 electrons, thus the number of unpaired
electrons is 3.
(ii) Fe (z = 26), Fe3+ (z = 23)
The electronic configuration of Fe3+ is given lay
1s2 2s2 2p6 3s2 3p6 3d5
The number of unpaired electron is 5.
(iii) Cr (z = 24), Cr2+ (z = 22)
The electronic configuration of Cr2+ is
1s2 2s2 2p6 3s2 3p6 3d4
The number of unpaired electron is 4.
(iv) Zn (z = 30), Zn2+ (z = 28)
The electronic configuration of Zn2+ is
1s2 2s2 2p6 3s2 3p6 3d10
The number of unpaired electron is 0.