Plus One Chemistry Notes Chapter 2 Structure of Atom is part of Plus One Chemistry Notes. Here we have given Kerala Plus One Chemistry Notes Chapter 2 Structure of Atom.
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|Chapter Name||Structure of Atom|
|Category||Plus One Kerala|
Kerala Plus One Chemistry Notes Chapter 2 Structure of Atom
The atomic theory of matter was first proposed by John Dalton. His theory, called Dalton’s atomic theory, regarded the atom as the ultimate particle of matter.
Sub atomic Particles :
The atoms can further divided into sub atomic particles; electrons, protons and neutrons.
Discovery of Electron
The experiments of Michael Faraday in discharge tubes showed that when a high potential is applied to a gas taken in the discharge tube at very low pressures, certain rays are emitted from the cathode. These rays were called cathode rays.+
Anode rays or Canal rays. Anode rays travelled in a direction opposite to that of the cathode rays. These rays are also called canal rays or positive rays.
Discovery of Protons and Neutrons
Electrical discharge carried out in the modified cathode ray tube led to the discovery of canal rays.
Thomson Model of Atom
Thomson proposed the first model of the atom. This model is also known as plum pudding, raisin pudding or water lemon. In this model, the positive charge is spread over a sphere with the electrons embedded in it such that the atom as a whole is electrically neutral.
Drawback of Thomson Model of Atom
This model of atom could account for the electrical neutrality of atom, but it could not explain the results of gold foil experiment carried out by Rutherford.
Rutherford’s Nuclear Model of Atom
Rutherford and his students (Hans Geiger and Ernest Marsden) bombarded very thin gold foil with α-particles. The experiment is known as α- particle scattering experiment.
- Most of the space in the atom is empty as most of the a-particles passed through the undeflected.
- A few postively charged a-particles were deflected.
- Very few were deflected back (180°).
- 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.
- The electrons move around the nucleus with a very high speed in circular paths called orbits.
- Electrons and the nucleus are held together by electrostatic forces of attraction.
Atomic number and Mass number
Atomic number (Z) = Number of protons or electrons in a neutral atom.
Mass number (A) = Number of protons (Z) + Number of neutrons (n)
Note: An element is thus represented by
Isobars and Isotopes
Isobars: These are the atoms with same mass number but different atomic number,
Isotopes: These are the atoms with same atomic number but different mass number.
eg., isotopes of hydrogen protium (1H1), deuterium (1D2), Tritium (1T3).
Developments Leading to the Bohr’s Model of Atom
Neils Bohr improved the model proposed by Rutherford. Two developments played a major role in the formulation of Bohr’s model of atom. These were:
- Electromagnetic radiation possess both wave like and particle like properties (Dual character).
- Experimental results regarding atomic spectra which can be explained only by assuming quantized electronic energy levels in atoms.
Wave Nature of Electromagnetic Radiation
Light is the form of radiation and it was supposed to the made of particles known as corpuscles.
As we know, waves are characterised by wave-length (λ), frequency (υ) and velocity of propagation (c) and these are related by the equation.
The wavelengths of various electromagnetic radiations increase in the order,
ϒrays < x-rays < UV rays < Visible < IR < Microwaves < Radio waves.
Wavelength (λ), frequency (υ) and wave number ( ).
is the distance between two adjacent crests or troughs.
is defined as the number of wave lengths travelled in one second.
Particle nature of Electromagnetic radiation: Planck’s Quantum Theory
Planck suggested that atoms and molecules emit (or absorb) energy only in discrete quantities and not in a continuous manner, at belief popular a that time. Planck gave the name quantum to the smallest quantity of energy that can be emitted or absorbed in the form of electromagnetic radiation. The energy (E) of a quantum of radiation is proportional to its frequency (υ) and is expressed by equation,
E = hυ
When a beam of light falls on certain metal plates like K, Cs etc. electrons are emitted which conduct electricity. This effect (discovered by Hertz) is known as photoelectric effect.
Threshold frequency (υ0)
Electrons are emitted from a metal surface only when the frequency (υ) of the incident radiation exceeds a certain characteristic value. This frequency is called the threshold frequency.
Energy of incident photons (hυ) = Work function (hυ0) + 1/2 me v2 K.E); where υ0 is the threshold frequency
Dual behaviour of Electromagnetic Radiation
Einstein proposed that light has a dual character i.e., pa
rticle and wave nature. The particle nature is explained on the basis of photo electric effect and wave nature is explained on the basis of interference and X-ray diffraction of aluminium foil.
Line spectrum of Hydrogen
The emission spectrum of hydrogen atom is obtained by passing electrical discharge through gas at very low pressure. Hydrogen atoms emit bluish light.
When the ray of this bluish light is passed through a prism, a discontinous line spectrum consisting of several sharp lines is obtain. Such a spectrum is called line spectrum of hydrogen.
Bohr model for Hydrogen atom
Bohr model for Hydrogen atom says that:
- The energy of an electron does not change with time.
- The frequency of radiation absorbed or emitted when transition occurs between two stationary states.
E1 and E2 are the energies of the lower and higher allowed energy states respectively.given stationary state can be expressed as in equation.
- The angular momentum of an electron in a
- The stationary states for electron are numbered n = 1, 2, 3, … These integral numbers are known as principal quantum numbers.
- The radii of the stationary states are expressed as,
- The most important property associated with the electron, is the energy of its stationary state. It is given by the expression
Rn is called Rydberg constant and its value is 2.18 x 10-18
- Bohr’s theory can also be applied to the ions containing only one electron, similar to that present in hydrogen like species.
eg., He+ Li2+, Be3+ and so on
Demerits of Bohr’s Model
- Inability to explain line spectra of multi electron atom.
- Inability to explain splitting of lines in the magnetic field (Zeeman effect) and in the electric filed (stark effect).
- Inability to explain the shapes of molecules.
- Inability to explain de Broglie concept of dual nature of matter and Heisenbergs uncertainty principle.
Explanation of Line Spectrum of Hydrogen
The frequency (υ) associated with the absorption and emission of the photon can be evaluated by using equation.
Towards Quantum Mechanical Model of the Atom
Two important developments which contributed significantly in the formulation of a more suitable and general model for atoms were:
- Dual behaviour of matter
- Heisenberg’s uncertainty principle.
Dual Behaviour of Matter
The French physicist, de Broglie proposed that matter, like radiation, should also exhibit dual behaviour i.e., both particle and wave like properties. This means that just as the photon, electrons should also have momentum as well as wavelength, de Broglie, from this analogy, gave the following relation between wave length (λ) and momentum (p) of a material particle.
Heisenberg’s Uncertainty Principle
It states that it is impossible to determine simultaneously and precisely both the exact position and exact momentum or velocity of a sub atomic particle like electron.
where Δp is the uncertainty in momentum of the particle, Δx is the uncertainty in position.
From this equation, it is seen that if Δp increases, the Δx decreases and vice versa. It is also can be written as follows:
Quantum Mechanical model of Atom
A new model of atom was developed on the basis of de Broglie hypothesis and uncertainty principle. This model is known as quantum mechanical model of atom.
Important features of the Quantum mechanical model of atom:
- The energy of electrons in atoms can have only certain specific values (quantized).
- Wave nature of electrons in the quantized energy levels.
- The idea of uncertainty in the position of electrons in the atom.
- An atomic orbital is the wave function ψ for an electron in an atom.
- Probability density ψ2 determines the probability of finding the moving electron in a given region.
An orbital may be defined as the region of space around the nucleus where there is maximum probability of finding an electron.
Quantum numbers are the set of four numbers with the help of which we can get complete information about the electrons in an atom. In order to specify energy, size, shape, orientation and spin of electron, four quantum numbers are required.
Write the correct set of four quantum numbers for the valence electron
(outer- i most electron) of potassium (Z =19).
K (19): 1s2 2s2 2p6 3s2 3p6 4s1
Valence electron: 4s1
n = 4,1 = 0, m = 0, s = + 1/2 or-1/2
In which of the following set of quantum numbers an electron will have the highest energy?
ii. 4,2,-1, -1/2
iii. 4,1, 0, -1/2
iv. 5, 0, 0, 1/2
4, 2, -1, -1/2 , i.e. 4d-electron will be highest energy because (n+1), i.e. 4 + 2 = 6 is highest.
Nodes or Nodal surfaces
Surfaces at which the probability of finding the electron goes to zero is called node.
s-orbitals have (n-1) nodes, p-orbitals have (n-2) nodes.
d-orbitals have (n-3) nodes. In general (n-/-1) nodes.
Energies of orbitals
The order of energy of orbitals are given below:
1s < 2s = 2p < 3s = 3p = 3d < 4s = 4p = 4f
The orbitals having same energy are called degenerate.
Using ‘s’, ‘p’, ‘d’, ‘f’ notations, describe the orbital with the following quantum numbers.
Filling of Orbitals in Atom
Aufbau Principle. According to this principle in the ground state of an atom, an electron will occupy the orbital of lowest energy and orbitals are occupied by electrons in the order of increasing energy.
Memory tip: For remembering order of energies
Hund’s rule of Maximum Multiplicity
This rule states that electron pairing in orbitals of same energy will not take place until each available orbital of a given subshell is single occupied (with parallel spin).
eg., N(7) has electronic configuration 1s2 2s2 2px1 2py2 2pz1 according to Hund’s rule and not 1s2 2s2 2px2 2py1.
Pauli’s Exclusion Principle.
It states that no two electrons in an atom can have the same values for all the four quantum numbers’. The maximum number of electrons that can be accomodated in a shell is 2n2, where n is the principal quantum number of the shell.
Extra Stability of Completely Filled and Half- Filled Orbitals
The electronic configuration of most of the atoms follows the basic rules. However certain elements such as Cr or Cu do not follow the rules. In such elements, the two sub-shells 4s and 3s slightly differ in energy, i.e., 4s is slightly lower in energy than 3d orbital.
- Cr atomic number 24 :
– 1s2, 2s2, 2p1, 3s2, 3p3, 3d4, 4s1
- Cu atomic number 29 :
– 1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s1
What is the energy in joules required to shift the electron of the hydrogen atom from the first Bohr orbit to the Fifth Bohr orbit. What is the wavelength of the light emitted when the electron returns to the ground state?
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