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Class- XI, , CHEMISTRY, , NOTES, , ATOMIC STRUCTURE, , Fundamental particles of a atom are :, 2. DINATIONS, Electron, A fundamental particle which carries one unit negative charge and has a mass, nearly equal to 1/1837th of that of hydrogen atom., Proton, A fundamental particle which carries one unit positive charge and has a mass, nearly equal to that of hydrogen atom., Neutron, A fundamental particle which carries no charge but has a mass nearly equal to, that of hydrogen atom., T
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THOMPSON ATOMIC MODEL, 3. THOMSON MODEL, Sir J. J. Thomson, who discovered the electron, was the first to suggest a, model of atomic structure., (i) All atoms contain electrons., (ii) The atom as a whole is neutral. The total positive charge, and total negative charge must be equal., He visualised all the positive charge of the atom as being spread out uniformly, throughout a sphere of atomic dimensions (i.e. approx. 10–10 m in diameter)., The electrons were smaller particles together carrying a negative charge, equal, to the positive charge in the atom. They were studded in the atom like plums in, a pudding. The charge distribution was such, that it gave the most stable, arrangement. This model of the atom was often called the plum – pudding, model. Also the raisin pudding model or watermelon model, , Drawbacks, Though the 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
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RUTHERFORD’S α SCATTERING EXPERIMENTRutherford conducted, , - particles scattering experiments in 1909. In this, , experiment, a very thin foil of gold (0.004nm) is bombarded by a fine stream of, alpha particles. A fluorescent screen (ZnS) is placed behind the gold foil, where, points were recorded which were emerging from α-particles., Polonium was used as the source of α-particles, , Observations, Rutherford carried out a number of experiments, involving the s c a t t e r i n g, of α particles by a very thin foil of gold., Observations were:, (i) Most of the α particles (99%) passes through it, without any deviation or, deflection., (ii) Some of the α particles were deflected through small angles.
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(iii) Very few α particles were deflected by large angles and occasionally an α, particle got deflected by 180o, Conclusions, (i) An atom must be extremely hollow and must consist of mostly empty space, because most of the particles passed, through it without any deflection., (ii) Very few particles were deflected to a large extent. This indicates that:, (a) Electrons because of their negative charge and very low mass cannot, deflect heavy and positively charged α particles, (b) There must be a very heavy and positively charged body in the atom, i.e. nucleus which does not permit the passage of positively charged α, particles., (c) Because, the number of α particles which undergo deflection of 180º,, is very small, therefore the volume of positively charged body must be, extremely small fraction of the total volume of the atom. This positively, charged body must be at the centre of the atom which is called nucleus., It has been found that radius of atom is of the order of 10–10m while the, radius of the nucleus is of the order of 10–15m. Thus if a cricket ball, represents a nucleus, then the radius of atom would be about 5 km., , Rutherford’s nuclear atomic model :, (i) An atom consists of tiny positively charged nucleus at the centre and it is, surrounded by hollow portion called extra nuclear part., (ii) The positive charge of the nucleus is due to nucleons which consist of, protons and neutrons while the electrons, present in extra nuclear portion has, negligible mass and carry a negative charge., (iii) The atom is electrically neutral, as the number of electrons is equal to, number of protons in it. Thus, total positive charge of the nucleus is balanced, by the total negative charge of electrons.
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(iv) The electrons in the extra nuclear part are revolving around the nucleus in, circular paths called orbits. Thus, an atom, , resembles the solar system in, , which the sun plays the role of nucleus and the planets that of revolving, electrons and the model is known as planetary model., (v) Electrons and nucleus are held together by the electrostatic force of, attraction., (vi) Forces of attraction operating on the electron are exactly balanced by, centrifugal forces., , Drawbacks, , (i) According to classical mechanics, any charged body in motion under the, influence of attractive forces should, radiate energy continuously. If this is so, the electron will follow a spiral path, and finally fall into the nucleus and the structure would collapse. This, behaviour is never observed., (ii) 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
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ATOMIC NUMBER & MASS NUMBER, Atomic number (Z) : Atomic number of an element is equal to the number, of unit positive charges or number of protons present in the nucleus of the, atom of the element. It also represents the number of electrons in the neutral, atom. Eg. Number of protons in Na = 11 , thus atomic number of Na=11, , Mass number (A) : The elementary particles (protons and neutrons) present, in the nucleus of an atom are collectively known as nucleons. The mass, number (A) of an atom is equal to the sum of protons and neutrons. It is, always a whole number. Thus, Mass number (A) = Number of protons(Z) +, Number of neutrons(n), Therefore, number of neutrons (n) = Mass Number (A) – Number of protons, (Z), n=A–Z, The general notation that is used to represent the mass number and atomic, number of a given atoms is, , XA, Z, , Where, X – symbol of element, A – Mass number, Z – atomic number, , Isotopes, Isobars, isotones and Isoelectronic, Isotopes:, Isotopes are the atoms of the same element having identical atomic number, but different mass number. The difference is due to the difference in number of, neutrons., The chemical properties of atoms are controlled by the number of electrons., Thus, isotopes of an element show same chemical behaviour., Isotopes of Hydrogen
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Isobars:, Atoms of different elements having different atomic numbersbut same mass, numbers are called isobars., Eg
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Isotones:, Atoms of different elements which contain the same number of neutrons are, called isotones., Eg, , Isoelectronic:, , The species (atoms or ions) containing the same number of electrons are called, isoelectronic., Eg. O2–, F–, Na+, Mg+2, Al+3, Ne etc
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BOHR ATOMIC MODEL :, This model is applicable to H-atom or H-like species like He+,Li2+,Be3+., Postulates, 1) An atom consists of a small, heavy, positively charged nucleus in the centre, and the electrons revolve around it in circular orbits., 2) Electrons revolve only in those orbits which have a fixed value of energy., Hence, these orbits are called energy levels or stationary states. They are, numbered as 1,2,3,...... These numbers are known as Principal quantum, Numbers., (a) Energy of an electron is given by:, En= –RH(Z2/n2) n = 1,2,3......., where RH is Rydberg’s constant and its value is 2.18 × 10–18 J., Z = atomic number, , Thus, energies of various levels are in the order: K < L < M < N...... and so on., Energy of the lowest state(n=1) is called ground state., (b) Radii of the stationary states:, , For H-atom (Z = 1), the radius of first stationary state is called Bohr orbit (52.9, pm), (c) Velocities of the electron in different orbits:
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Since the electrons revolve only in those orbits which have fixed values of, energy, hence electrons in an atom can have only certain definite values of, energy and not any of their own. Thus, energy of an electron is quantised., 3), , 4) Like energy, the angular momentum of an electron in an atom can have, certain definite values and not any value of their own., , 5) An electron does not lose or gain energy when it is present in the same shell., 6) When an electron gains energy, it gets excited to higher energy levels and, when it de-excites, it loses energy in the form of electromagnetic radiations and, comes to lower energy values., , What does negative energy for Hydrogen atom means?, This negative sign means that the energy of the electron in the atom is lower, than the energy of a free electron at rest. A free electron at rest is an electron, that is infinitely far away from the nucleus (n = α) and is assigned the energy, value of zero. As the electron gets closer to the nucleus (as n decreases), En, becomes more and more negative. The most negative energy value is given by, n=1 which corresponds to the most stable orbit., Transition of Electron, We know that energy is absorbed or emitted when electron excites or de-excites, respectively. The energy gap between the two orbits is
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The wavelength associated with the absorption or emission of the photon is:, , This is known as Rydberg’s formula., 1.09677 × 107 m-1 is also known as Rydberg’s constant, , Line Spectrum of Hydrogen, When an electric discharge is passed through gaseous hydrogen, the H2, molecules dissociate and the energetically excited hydrogen atoms produced, emit electromagnetic radiations of discrete frequency. The hydrogen spectra, consists of several lines named after their discoverer.We get discrete lines and, not a continuous spectra because the energy of an electron cannot change, continuously but can have only definite values. Thus we can say that energy, of an electron is quantized
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LYMAN SERIES :, When an electron jumps from any of the higher states to the ground state or, first state (n = 1) ,the series of spectral lines emitted lies in the ultra violet, region and are called as Lyman series., Therefore , in Rydberg’s formula n1= 1, n2= 2,3,4,5..., BALMER SERIES :, When an electron jumps from any of the higher states to the state with n=2,the, series of spectral lines emitted lies in the visible region and are called as, Balmer series., Therefore , in Rydberg’s formula n1= 2, n2= 3,4,5,6...., PASCHEN SERIES :, When an electron jumps from any of the higher states to the state with n=3, ,the series of spectral lines emitted lies in the infrared region and are called as, Paschen series., Therefore , in Rydberg’s formula n1= 3, n2= 4,5,6..., BRACKETT SERIES :, When an electron jumps from any of the higher states to the state with n =, 4,the series of spectral lines emitted lies in the infrared region and are called, as Brackett series., Therefore , in Rydberg’s formula n1= 4, n2= 5,6,7...
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PFUND SERIES :, When an electron jumps from any of the higher states to the state with n =, 4,the series of spectral lines emitted lies in the infrared region and are called, as Pfund series. Therefore , in Rydberg’s formula n1= 5, n2= 6,7..., Limitations of Bohr’s Model, 1) Inability to explain line spectra of multi-electron atoms., 2) It fails to account for the finer details (doublet-two closely spaced lines), ΔΔthe hydrogen spectra., 3) Inability to explain splitting of lines in the magnetic field (Zeeman Effect) and, in the electric field (Stark Effect)- If the source emitting the radiation is placed, in magnetic or electric field, it is observed that each spectral line splits up into, a number of lines., Zeeman Effect:, Splitting of spectral lines in magnetic field is known as Zeeman Effect, Stark Effect:, splitting of spectral lines in electric field is known as Stark Effect., 4) It could not explain the ability of atoms to form molecules by covalent bonds., 5) He ignores dual behaviour of matter and also contradicts Heisenberg, uncertainty principle., Heisenberg’s Uncertainty Principle, It is impossible to measure simultaneously the position and momentum of a, small particle with absolute accuracy. If an attempt is made to measure any of, these two quantities with higher accuracy, the other becomes less accurate., The product of the uncertainty in the position (x) and the uncertainty in, momentum (p) is always a constant and is equal to or greater than h/4π., , (x). (p) h/4 Or, , Explaination, , (x). (mv) h/4 Or, , (x). (x) h/4m
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Change of momentum and position of electron on impact with, a photon, Suppose we attempt to measure both the position and momentum of an, electron. To pin point the position of the electron we have to use light so that, the photon of light strikes the electron and the reflected photon is seen in the, microscope. As a result of the hitting, the position as well as the velocity of the, electron are disturbed., Significance of Uncertainity Principle:, It rules out the existence of definite paths or trajectories of electrons as stated, in Bohr’s Model., To go further into the atomic mysteries, we will have to understand the nature, of electromagnetic radiations and studyMaxwell’s Electromagnetic Wave, theory”. James Maxwell was the first to give a comprehensive explanation, about the interaction between the charged bodies and the behaviour of electric, and magnetic fields, , ELECTROMAGNETIC RADIATIONS, Electromagnetic Radiations are waves which are formed as a result of, oscillating magnetic and electric fields which are perpendicular to each other, and both are perpendicular to direction of motion., , They do not require any medium and can move in vacuum unlike sound waves., Light is a form of radiation and has wave characterstics., The various characterstics of a wave are:
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1) Amplitude : It is height of the crest or trough (depth) of a wave., Units :, metre (m), 2) Frequency ( γ ) : The number of waves passing through a point in one, second. Units : Hertz (Hz) or s–1, 3) Time Period : The time taken by a wave to complete one vibration is called, time period. Units : sec, 4) Velocity : The distance travelled by a wave in one second is called velocity., Units : m/s, In vacuum, all types of electromagnetic radiations travel at the same speed i.e., 3 × 108 m/s. This is called speed of light., 5) Wavelength( λ ) : The distance between two adjacent crests or troughs is, called wavelength. Units : Angstrom(Å), [1 Å=10–10m], 6) Wave Number ( ΰ) : It is the number of wavelengths per centimetre of length., Units : m-1, , ΰ, , = 1/, , ELECTRO MAGNETIC SPECTRUM, When all the electromagnetic radiations are arranged in increasing order of, wavelength or decreasing frequency the band of radiations obtained is termed, as electromagnetic spectrum.
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The visible spectrum is a subset of this spectrum (VIBGYOR) whose range of, wavelength is 380-760nm., The wavelengths increase in the order:, Gamma Rays < X-rays < Ultra-violet rays < Visible< Infrared < Micro-waves, <Radio waves., , Electromagnetic Wave Theory:, The main points of this theory are:, (1) A source (like the heated rod) emits energy continuously in the form of, radiations (i.e. no change in wavelength, or, frequency of the emitted, radiations even on increasing the energy radiated)., (2) These radiations are Electromagnetic in nature.
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Failure of EM wave theory:, The theory failed because of 2 experiments:, , (1) Black Body Radiation :, According to Maxwell’s theory on heating a body the intensity should increase,, that is, energy radiated per unit, area should increase without having any effect on the wavelength or frequency., But we observe that when we heat an iron rod, it first turns to red then white, and then becomes blue at very high temperatures. This means that frequency, of emitted radiations is changing. An ideal body, which emits and absorbs, radiations of all, frequencies is called black body and radiation emitted by a black body is called, black body radiation, The variation of intensity with wavelength at different temperatures for a black, body is shown below:, , So it is observed that with increasing temperature the dominant wavelength in, the emitted radiations decreases, and the frequency increases., That is at higher temperatures, though the intensity rises as predicted by, Maxwell’s theory but the wavelength, decreases. If T1>T2>T3 then λ1< λ2< λ3., (2) Photoelectric Effect, :, When radiations with certain minimum frequency ( 0) strike the surface of a, metal, the electrons are ejected from the surface of the metal. This phenomena, is called photoelectric effect. The electrons emitted are called, photoelectrons., According to Maxwell’s Theory the ejection of electrons should depend on, intensity of radiation that is if electrons
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are not being ejected, then on increasing the intensity they can be ejected., , The following observations are made:, (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 the metal 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, 0 (also, known as threshold frequency) below which photoelectric effect is not, observed. At a frequency > 0, 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., Thus, these findings were contradictory to the Maxwell’s theory. The number of, electrons ejected and kinetic energy, associated with them should depend on the intensity of light. It has been, observed that though the number of, electrons ejected does depend upon the brightness of light, the kinetic energy of, the ejected electrons does not., To justify these findings Max Von Planck gave his Quantum theory.
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PLANK’S QUANTUM THEORY, The main points of this theory are:, , (i) The energy is emitted or absorbed not continuously but discontinuously in, the form of small discrete packets of, energy. Each such packet of energy is called a ‘quantum’. In case of light this, quantum of energy is called a photon, (ii) At a time only one photon can be supplied to one electron or any other, particle., (iii) One quantum cannot be divided or distributed., (iv) The energy of each quantum is directly proportional to the frequency of, radiation., , hc, E or E = h = −−, , h = Planck’s constant = 6.626 × 10-34Js, (v) The total energy emitted or absorbed by a body will be in, whole number quanta., Hence E = nhγ =, , nhc, ----λ, , This is also called “Quantisation of energy”., , Explanation of Black body radiation:, As the temperature is increased the energy emitted increases, thereby, increasing the frequency of the emitted radiations.As the frequency increases, the wavelength shifts to lower values., , Explanation of Photoelectric effect:, (i) When light of some particular frequency falls on the surface of metal, the, photon gives its entire energy to the electron of the metal atom. The electron, will be ejected from the metal only if the energy of the photon is sufficient to, overcome the force of attraction of the electron by the nucleus. So,, photoelectrons are ejected only when the, incident light has a certain minimum frequency (threshold frequency 0). The, Threshold energy required for emission is called “Work Function” that is “h0”.
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(ii) If the frequency of the incident light () is more than the threshold frequency, (0), the excess energy is imparted to the electron as kinetic energy. Hence,, Energy of one quantum = Threshold Energy + Kinetic Energy, 2, , h = h0 + (1/ 2) mev, (iii) When > 0, then on increasing the intensity the number of quanta incident, increases thereby increasing the number of photoelectrons ejected., (iv) When > 0, then on further increasing the frequency, the energy of each, photon increases and thus kinetic energy of each ejected electron increases., Energy can also be expressed in Electron Volt(eV), , ., , The energy acquired by an electron when it is accelerated through a potential, difference of one Volt., 1eV = 1.602 × 10–19J, , CONCLUSION :, Light has both the Wave nature (shows the phenomena of diffraction and, interference) and Particle nature (could, explain the black body radiation and photoelectric effect) .Thus, light has dual, nature. Bohr’s Model is based on “Atomic Spectra”, therefore before moving, further we will study, , WHAT IS SPECTRUM, A spectrum is a group or band of wavelengths/colours and the study of, emission or absorption spectra is known as, Spectroscopy, , Types of spectrum:, There are two types of spectrum:, 1) Emission Spectrum, 2) Absorption Spectrum, Emission Spectrum-When radiations emitted from a source are incident on a, prism and are separated into different wavelengths and obtained on a, photographic plate., (a) Continuous Emission Spectra:, There are no gaps between various wavelengths, one wavelength merges, into another.
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(b) Discontinuous Emission Spectra:, It is also known as Line Spectra or atomic spectra. In this, certain, wavelengths go missing from a group and leaves dark spaces in between, giving discontinuity to the spectrum. It is also known as fingerprint of an, element, , Absorption Spectra, When light from any source is first passed through th solution of a chemical, substance and then analysed, it is observed that there are some dark lines in, the otherwise continuous spectra.
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QUANTUM MECHANICAL MODEL OF A ATOMB, Quantum mechanics is a theoretical science that deals with the study of the, motion of microscopic objects which have both particle like and wave like, properties. The fundamental equation of quantum mechanics was developed by, Schrodinger. This equation describes a function called electron wave, function ( ). This wave function stores all the information about an electron, like energy, position, orbital etc. As such it does not have any physical, significance. The information stored in about an electron can be extracted in, terms of Quantum Numbers., Probability Density, | |2 is the probability of finding the electron at a point within an atom., Concept of Orbital, It is a three dimensional space around the nucleus within which the probability, of finding an electron of given energy is maximum (say upto 90%)., , Quantum Numbers, They may be defined as a set of four numbers with the help of which we can get, complete information about all the, electrons in an atom i.e. location, energy, type of orbital occupied, shape and, orientation of that orbital etc., The three quantum numbers called as Principal, Azimuthal and Magnetic, quantum number are derived from, Schrodinger wave equation. The fourth quantum number i.e. the Spin, quantum number was proposed later on., 1) Principal Quantum Number (n):, It tells about the shell to which an electron belongs. n = 1,2,3,4,5..... and so, on., This number helps to explain the main lines of the spectrum on the basis of, electronic jumps between these shells.
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(a) It gives the average distance of the electron from the nucleus. Larger the, value of n, larger is the distance from, the nucleus., (b) It completely determines the energy in hydrogen atom or hydrogen like, species., The energy of H-atom or H-like species depends only on the value of n., Order of energy : 1 < 2 < 3 < 4 < 5....... and so on., For multi-electron species, energy depends on both principal and azimuthal, quantum number., The maximum number of electrons present in any shell = 2n2, 2) Azimuthal Quantum Number (l) :, Also known as Orbital Angular momentum or Subsidiary, quantum number.Within the same shell, there are number of sub-shells, so, number of electronic jumps increases and this explains the presence of fine, lines in the spectrum. This quantum number tells about :, (a) The number of subshells present in a shell., (b) Angular momentum of an electron present in subshell, , (c) Shapes of various subshells present within the same shell., (d) Relative energies of various subshells., Value of l varies from 0 to n – 1, For 1st shell (n = 1): l = 0, For 2nd shell (n = 2): l = 0,1, For 3rd shell (n = 3): l = 0,1,2, For 4th shell (n = 4): l = 0,1,2,3, , The notations s,p,d,f represent the initial letters of the word sharp, principal,, diffused and fundamental. In continuation l = 4 is called g subshell and l = 5 is, called h subshell and so on.
Page 25 : The number of subshells present in any principal shell is equal to the number, of the principal shell., Energies of various subshell present within the same shell is:, s<p<d<f, Angular momentum of an electron in orbital :, , (3) Magnetic Quantum Number(m):, This quantum number is required to explain the fact that when the source, producing the line spectrum is placed in a magnetic field, each spectral line, splits up into a number of line (Zeeman effect)., Under the influence of external magnetic field, electrons of a subshell can, orient themselves in a certain preferred regions of space around the nucleus, called orbitals. The magnetic quantum number determines the number of, preffered orientations of the electrons present in a subshell. Since each, orientation corresponds to an orbital, thus magnetic quantum number, determines the number of orbitals present in any subshell., Value of m ranges from – l to +l including zero.
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These orbitals of the same subshell having equal energy are called degenerate, , ., , orbitals Eg. The three p-orbitals of a particular principal shell have, , the same energy in the absence of magnetic field1 Similarly, all five, orbitals of d-subshell of a particular shell, have the same energy., Thus, for H-atom order of energy is:, 1s < 2s = 2p < 3s = 3p = 3d < 4s = 4p = 4d = 4f < .........., For multi electron atoms, the energy of the orbitals decreases with, increase in effective nuclear charge. Eg, , E2s (H) > E2s (Li) > E2s (Na) > E2s (K), The total possible values of m in a given subshell = 2l + 1, Total no. of orbitals in a given shell = n2
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4) Spin Quantum Number(s):, The electron in an atom not only moves around the nucleus but also spins, about its own axis. Since the electron in an orbital can spin either in clockwise, or anti-clockwise direction. Thus s can have only two values, , +½, , & -½, , This quantum number helps to explain the magnetic properties of substances., An orbital cannot have more than two electrons and, these electrons should be of opposite spin., Thus, maximum number of electrons in s-subshell = 2, Maximum number of electrons in p-subshell = 6, Maximum number of electrons in d-subshell = 10, Maximum number of electrons in f-subshell = 14, , Shapes of atomic orbitals:, (1) Shape of s-orbitals:, (a) They are non-directional and spherically symmetric i.e. probability of finding, the electron at a given distance is equal in all directions., (b) 1s orbital and 2s orbital have same shape but size of 2s is larger., (c) There is a spherical shell within 2s orbital where electron density is zero and, is called a node, (d) The value of azimuthal quantum number(l) is zero (l=0) and magnetic, quantum number can have only one value, i.e. m = 0
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(2) Shape of p-orbitals:, (a) It consists of two lobes present on either side of the plane that passes, through the nucleus. The p-orbital is dumb-bell shaped., (b) There are three possible orientations of electron cloud in p-orbitals., Therefore, the lobes of p-orbital may be considered to be along x,y and z axis., Hence they are designated as px,py,pz. The three p-orbitals are oriented at, right angles to one another., (c) First main energy level( Principal quantum number n = 1) does not contain, any p-orbital., (d) The three p-orbitals of a particular energy level have same energy in, absence of an external electric and magnetic field and are called degenerate, orbitals., (e) Like s orbitals, p-orbitals increase in size with increase in the energy of, main shell of an atom. Thus, value of azimuthal quantum number is one (l=1), and magnetic quantum number has three values (m= –1, 0, +1), , (3) Shapes of d-orbitals:, (a) They are designated as dxy, dyz, dzx and x2 y2 d z2. They have shape of, double dumbell., (b) All five d orbitals have same energy in the absence of, magnetic field., (c) For principal shell number 1 and 2, there are no d orbitals.
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Nodes and nodal planes :, Node:, It is a region of zero probability., There are two types of nodes:, (1) Radial or Spherical nodes: Three dimensional regions in an orbital where, probability of finding the electron becomes zero., Number of radial/ spherical nodes = n – l – 1, , (2) Planar or Angular Nodes: They are the planes cutting through the nucleus, on which probability of finding the, electron is zero., Number of Planar/Angular Nodes: l, Total Number of nodes: n - 1
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Filling of orbitals in an atom, , :, , (1) Aufbau Principle:, 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 when the, lower energy orbitals are filled., Unlike H-atom where energy of orbitals depend only on n, energy of orbitals of, multi-electron orbitals depend on both n and l. Their order of energy can be, determined using (n+l) rule., According to this rule, lower the value of (n+l) for an orbital, lower is its energy., If two different types of orbitals have the same value of (n+l), the orbital with, lower value of n has lower energy.
