• The atomic theory of matter was first proposed on a firm scientific basis by John Dalton, called Dalton’s atomic theory, regarded the atom as the ultimate particle of matter. In his theory he proposed the following :

1. Matter consists of indivisible atoms.

2. All the atoms of a given element have identical properties including identical mass. Atoms

of different elements differ in mass.

3. Compounds are formed when atoms of different elements combine in a fixed ratio.

4. Chemical reactions involve reorganization of atoms. These are neither created nor

destroyed in a chemical reaction.

the experimental observations made by scientists towards the end of

nineteenth and beginning of twentieth century established that atoms can be further divided into subatomic particles, i.e., electrons, protons and neutrons.

Discovery of Electron

• In mid 1850s many scientists mainly Faraday began to study electrical discharge in partially evacuated tubes, known as cathode ray discharge tubes.
• A cathode ray tube is made of glass containing two thin pieces of metal, called electrodes, sealed in it. 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.
• The flow of current from cathode to anode was further checked by making a hole in the anode and coating the tube behind anode with phosphorescent material zinc sulphide.
• When these rays, after passing through anode, strike the zinc sulphide coating, a bright spot on the coating is developed(same thing happens in a television set).

The results of these experiments are summarised below.

• The cathode rays start from cathode and move towards the anode.
•   These rays themselves are not visible but their behavior can be observed with the help of certain kind of materials (fluorescent or phosphorescent) which glow when hit by them.
• In the absence of electrical or magnetic field, these rays travel in straight lines

Millikan’s Oil Drop Method

• In this method, 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 downward motion of these droplets was viewed through the telescope, equipped with a micrometer eye piece.
• By measuring the rate of fall of these droplets, Millikan was able to measure the mass of oil droplets.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. By carefully measuring the effects of electrical field strength on the motion of oil droplets.
• Millikan 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…

Discovery of Protons and Neutrons

Electrical discharge carried out in the modified cathode ray tube led to the discovery of particles carrying positive charge, also known as canal rays. The characteristics of these positively charged particles are listed below.

(i) unlike cathode rays, the positively charged particles depend upon the nature of gas present in the cathode ray tube. These are simply the positively charged gaseous ions.

(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 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. Later, a need was felt for the presence of electrically neutral particle as one of the constituent of atom. These particles were discovered by Chadwick (1932) by bombarding a thin sheet of beryllium by a-particles.
• When electrically neutral particles having a mass slightly greater than that of the protons was emitted. He named these particles as neutrons.

ATOMIC MODELS

• Observations obtained from the experiments mentioned in the previous sections have suggested that Dalton’s indivisible atom is composed of sub-atomic particles carrying positive and negative charges. Different atomic models were proposed to explain the distributions of these charged particles in an atom.

Thomson Model of Atom J. J. Thomson, in 1898, proposed that an atom possesses a spherical shape (radius approximately 10–10 m) in which the positive charge is uniformly distributed. The electrons are embedded into it in such a manner as to give the most stable electrostatic arrangement

• Many different names are given to this model, for example watermelon model.
• This model can be visualised as a pudding or watermelon of positive charge with plums or seeds (electrons) embedded into it. Although this model was able to explain the overall neutrality of the atom, but was not consistent with the results of later experiments.

Rutherford’s Nuclear Model of Atom Rutherford and his students (Hans Geiger and Ernest Marsden) bombarded very thin gold foil with a–particles. Rutherford’s famous a–particle scattering experiment is represented in the below Fig. A stream of high energy a–particles from a radioactive source was directed at a thin foil (thickness ∼100 nm) of gold metal. The thin gold foil had a circular fluorescent zinc sulphide screen around it.

Observations: (i) most of the a– particles passed through the gold foil undeflected.

(ii) a small fraction of the a–particles was deflected by small angles.

(iii) a very few a- particles ( 1 in 20,000)

bounced back, that is, were deflected by

nearly 180°.

Conclusions: (i) Most of the space in the atom is empty as most of the –particles passed

through the foil undeflected.

(ii) A few positively charged a– particles were deflected. The deflection must be due to enormous repulsive force showing that the positive charge of the atom is not spread throughout the atom as Thomson had presumed. The positive charge has to be concentrated in a very small volume that repelled and deflected the positively

charged a– particles.

(iii) Calculations by Rutherford showed that the volume occupied by the nucleus is negligibly small as compared to the total volume of the atom.

• On the basis of above observations and conclusions, Rutherford proposed the nuclear model of atom (after the discovery of protons). According to this model :

(i) The positive charge and most of the mass of the atom was densely concentrated in 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. Thus, Rutherford’s model of atom resembles the solar system in which the nucleus plays the role of sun and the electrons that of revolving planets.

(iii) Electrons and the nucleus are held together by electrostatic forces of attraction.

Drawbacks of Rutherford Model

• Rutherford model cannot explain the stability of an atom.
• Another serious drawback of the Rutherford model is that it 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.

Atomic Number and Mass Number

Atomic Number: It is the number of protons present in the nucleus and is represented by (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 respectively. 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. 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 (Z) + number of neutrons (n)

Isobars and Isotopes

• The composition of any atom can be represented by using the normal element symbol (X) with super-script on the left hand side as the atomic mass number (A) and subscript (Z) .
• Isobars are the atoms with same mass number but different atomic number for example, 146C and 147N.
• Isotopes: atoms with identical atomic number but different atomic mass number are known as Isotopes. It is evident that difference between the isotopes is due to the presence of different number of neutrons present in the nucleus. For example, considering of hydrogen atom again, 99.985% of hydrogen atoms contain only one proton. This isotope is called protium( 11H). Rest of the percentage of hydrogen atom contains two other isotopes, the one containing 1 proton and 1 neutron is called deuterium (12D, 0.015%) and the other one possessing 1 proton and 2 neutrons is called tritium (13T ).
•  Important point to mention regarding isotopes is that chemical properties of atoms are controlled by the number of electrons, which are determined by the number of protons in the nucleus.

Number of neutrons present in the nucleus have very little effect on the chemical properties of an element. Therefore, all the isotopes of a given element show same chemical behaviour

Planck’s Quantum Theory

Black Body: A body, which emits and absorbs radiations of all frequencies, is called a black body and the radiation emitted by such a body is called black body radiation.

• Planck suggested that atoms and molecules could emit (or absorb) energy only in discrete quantities. He stated that The energy is emitted or absorbed not continuously, but the energy emited or absorbed discontinuously in the form of small discrete packets. Each such packet of energy is called a ‘quantum’.

In case of light this quantum of energy is called a photon.

${E=h{\nu }={\frac {hc} {\lambda }}}$

Photoelectric Effect

• When light of some particular frequency falls on the metal surface then the electron will be ejected from the metal surface. The phenomenon is called Photoelectric effect.

The results observed in this experiment were:

(i) The electrons are ejected from the metal surface as soon as the beam of light strikes the surface, i.e., there is no time lag between the striking of light beam and the ejection of electrons from themetal surface.

(ii) The number of electrons ejected is proportional to the intensity or brightness of light.

(iii) For each metal, there is a characteristic minimum frequency, no (also known as

threshold frequency) below which photoelectric effect is not observed. At a

frequency,  n>no, the ejected electrons come out with certain kinetic energy.

• The kinetic energies of these electrons increase with the increase of frequency of the light used. Einstein (1905) was able to explain the photoelectric effect using Planck’s quantum theory.
• When a photon of sufficient energy strikes an electron in the atom of the metal, it transfers its energy instantaneously to the electron during the collision and the electron is ejected without any time lag or delay.

Greater the energy possessed by the photon, greater will be transfer of energy to the electron and greater the kinetic energy of the ejected electron. In other words, kinetic energy of the ejected electron is proportional to the frequency of the electromagnetic radiation. Since the striking photon has energy equal to hn and the minimum energy required to eject the electron is hn0 (also called work function, W0;then the difference in energy (hn– hn0 ) is transferred as the kinetic energy of the photoelectron.

Quantum Numbers

• A set of numbers which helps in getting the complete information about the electrons i.e. location, energy,type of orbital occupied etc.
• There are 4 types of quantum numbers namely, i) principal quantum number (ii) Azimuthal Quantum number (iii) Magnetic Quantum number (iv) Spin Quantum number.
• The principal quantum number n’ is a positive integer with value of n = 1,2,3……..

The principal quantum number determines the Size.

• Size of an orbital increases with increase of principal quantum number ‘n’.
• Azimuthal quantum number: this is represented by ‘l and itis also

known as orbital angular momentum or subsidiary quantum number.

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)

• For example, when n = 1, value of l is only 0. For n = 2, the possible value of l can be 0 and 1. For n = 3, the possible l values are 0, 1 and 2.

Value for l :               0 1 2 3 4 5 …………

notation for:                s p d f g h …………

sub-shell

• Magnetic orbital quantum number: magnetic Quantum numbers are represented by ‘ml’. and the  values ranges from -l to +l including zero.
• Magnetic quantum number determines the number of orbitals present in any subshell. For any sub-shell (defined by ‘l’ value) 2l+1 values of ml are possible and these values are given by :

ml = – l, – (l –1), – (l–2)… 0,1… (l – 2), (l–1), l

Thus for l = 0, the only permitted value of ml = 0, [2(0)+1 = 1, one s orbital].

• For l = 1, ml can be –1, 0 and +1 [2(1)+1 = 3, three p

orbitals].

• For l = 2, ml = –2, –1, 0, +1 and +2,

[2(2)+1 = 5, five d orbitals]. It should be noted that the values of ml are derived from l and  that the value of l are derived from n.

• Spin Quantum Number: Uhlenbeck and Samuel Goudsmit proposed the presence of the fourth quantum number known as the electron spin quantum number (ms). two spin states of the electron and are normally represented by two arrows,   (spin up) and   (spin down).
• Two electrons that have different ms values the two possible spin values are +1/2 and -1/2
• An orbital cannot hold more than two electrons and these two electrons should have opposite spins. ml designates the orientation of the orbital.
• For a given value of l, ml has (2l+1) values, the same as the number of orbitals per subshell ms refers to orientation of the spin of the electron.

Shapes of Atomic Orbitals

• In general, it has been found that ns-orbital has (n – 1) nodes, that is, number of nodes increases with increase of principal quantum number n. In other words, number of nodes for 2s orbital is one, two for 3s and so on.

Shape of the s- orbital: All the s-orbitals are spherically symmetric, that is, the probability of finding the electron at a given distance is equal in all the directions.

• It is also observed that the size of the s orbital increases with increase in n, that is, 4s > 3s > 2s > 1s and the electron is located further away from the nucleus as the principal quantum number increases.

Shape of p-orbital: Each p orbital consists of two sections called lobes that are on either side of the plane that passes through the nucleus. The probability density function is zero on the plane where the two lobes touch each other

Shape of d- orbital : There are five ml values (–2, –1, 0, +1 and +2) for l = 2 and thus there are five d orbitals. The five d-orbitals are designated as dxy, dyz, dxz, dx2–y2 and . The shapes of the first four d-orbitals are similar to each other, where as that of the fifth one, , is different from others  –

Energies of Orbitals

• The energy of an electron in a multielectron atom, unlike that of the hydrogen atom, depends not only on its principal quantum number (shell), but also on its azimuthal quantum number (subshell) that is, energy of electron in an orbital, as mentioned earlier, depends upon the values of n and l.
• The lower the value of (n + l) for an orbital, the lower is its energy. If two orbitals have the same value of (n + l), the orbital with lower value of n will have the lower energy.

Aufbau Principle: The aufbau 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 :

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 4f,

5d, 6p, 7s

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 spin.” This means that the two electrons can have the same value of three quantum numbers n, l and ml, but must have the opposite spin quantum number.

Hund’s Rule of Maximum Multiplicity:

• This rule deals with the filling of electrons into the orbitals belonging to the same subshell this rule states that pairing of electron takes place only after all the orbitals are half filled.
• Since there are three p, five d and seven f orbitals, therefore, the pairing of electrons will start in the p, d and f orbitals with the entry of 4th, 6th and 8th electron, respectively. It has been observed that half filled and fully filled degenerate set of orbitals acquire extra stability due to their symmetry.

For the further notes on topics Electro magnetic Radiation, line spectrum of hydrogen, bhor’s model and drawbacks, dual nature of matter, and Heisenberg’s uncertainity principleplease refer to the above given pdf in which they were explained in a detailed manner.

Categories: General