Structure of Atom Class 9 Notes Science

Structure of Atom Class 9 Notes for CBSE

Introduction

            The word atom is a Greek Word meaning ‘indivisible’ Dalton’s theory was a landmark in the history of chemistry. In 1833, Michael Faraday showed that there is a relationship between matter and electricity. This was the first major break through to suggest that atom was not a simple indivisible particle of all matter but was made up of small particles. Discovery of electrons, protons and neutrons discarded the indivisible nature of the atom proposed by John Dalton.

            The modern structure of atom is :

Figure 

            The atom is neutral in nature because it has equal no. of negatively charged electrons and positively charged protons. Neutron is neutral in nature.

Production of Cathode Rays – Discovery of Electron

            The existence of electrons in an atom was shown by J.J. Thomson in 1897 by passing electricity at high voltage through a gas at very low pressure. The purpose of applying high electrical voltage is to supply electrical energy to break the atoms of the gas into smaller particles. Before we go further, we will describe the construction of a discharge tube which is used for the production of cathode rays which led to the discovery of electron.

            A common discharge tube is a long glass tube having two metal plates sealed at its two ends (figure). These metal plates are known as electrodes. The electrode which is connected to the positive terminal of the battery is known as anode (positive electrode), and the electrode which is connected to the negative terminal of the battery is called cathode (negative electrode). Discharge tube has a side tube through which air (or other gases) can be pumped out by using a vacuum pump, so that experiments can be performed at low pressures. In the following discussion, we will use air as gas in the discharge tube.

(i)         When air inside the discharge tube is at the atomspheric pressure and a high electric voltage of 10,000 volts (or more) is applied to the electrodes, no electricity flows through the air in the discharge tube.

(ii)         If the pressure of air inside the discharge tube is reduced to about 1 mm of mercury and high voltage is applied again, electricity begins to flow through air and a light is emitted by the air inside the tube. The colour of light changes with the nature of gas taken in the discharge tube.

Figure : Production of cathode rays

(iii)        When the pressure of air in the discharge tube is reduced to about 0.001 mm of mercury and a high voltage is applied to the electrodes, the emission of light by air stops. Though the inside of the discharge tube now appears to be dark, the walls of the discharge tube at the end opposite to the cathode begin to glow with greenish light. It is now known that some invisible rays are formed at the cathode and when these rays strike the glass tube, they emit a greenish light. Since these rays are formed at the cathode, they are known as cathode rays.

Properties of Cathode Rays

            The cathode rays possess the following properties :

1. Cathode rays travel in straight lines: An object placed in the path of cathode rays casts a sharp shadow. It shows that cathode rays travel in straight lines.
2. Heating effect : When cathode rays are focused on a thin metal foil, it gets heated up to incandescence.
3. Cathode rays consist of material particles : This was indicated by the fact that a light paddle wheel placed in the path of cathode rays starts rotating.
4. Effect of electric field : When electric field is applied to a stream of cathode rays, they get deflected towards positive plate. It shows that cathode rays themselves are negatively charged.
5. Effect of magnetic field : When magnetic field is applied, perpendicular to the path of cathode rays, they get deflected in the direction expected for negative particles. This further confirmed that cathode rays are negatively charged.
6. On striking against walls of the discharge tube cathode rays produce faint greenish fluorescence.
7. Cathode rays ionize the gas through which they pass.
8. Cathode rays produce X-rays when they are made to fall on metals such as tungsten, copper, etc.
9. They can penetrate through thin metal foils.
10. The charge to mass ratio (e/m) for the particles in the cathode is independent of the nature of the gas taken in the discharge tube or the nature of the cathode.

Electron

(1)        The name electron was proposed by Stoney for the fundamental unit of electricity and used as such.

(2)        The charge on an electron (–1.602×10^–19 coulomb or – 4.8 × 10^–10 esu) was determined by Mullikan in his oil drop experiment.

(3)        Actual mass of an electron (9.11×10^–28g) was given by J.J. Thomson. Of the three fundamental particles of an atom, electron is lightest.

(4)        The specific charge (e/m ratio) of electrons (cathode rays) was determined by Thomson as 1.76 × 10⁸ coulomb/gram. The specific charge of electron decreases with increase in its velocity because increase in velocity increases the mass of electron. The e/m ratio of electron was found to be independent of the nature of gas and electrode used. Thus, electrons are present in all atoms or these are fundamental constituents of all kinds of matter.

(5)        Radius of the electron is found to be 42.8 × 10^–15 m.

(6)        Density of electron is found to be 2.17 × 10¹⁷ g/cm³.

(7)        Mass of one mole of electron is nearly 0.55 mg.

(8)        Charge on one mole of electron is ≈ 96500 coulomb or 1 Faraday.

Discovery of Proton Production of Anode rays

                   The formation of cathode rays has shown that all the atoms contain negatively charged particles called electrons. Now, an atom is electrically neutral, so it must contain some positively charged particles to balance the negative charge of electrons. It has actually been found by experiments that all the atoms contain positively charged particles called protons. The existence of protons in the atoms was shown by Goldstein. We will now discuss the production of anode rays (or positive rays) which led to the discovery of proton.

Figure 

                   In the production of positive rays a discharge tube having perforated cathode is used. A perforated cathode is a cathode having holes in it. These perforations or holes are to allow the positive rays to pass through them. When a high voltage of about 10000 volts is applied to a discharge tube having a perforated cathode and containing air at very low pressure of about 0.001 mm of mercury, a faint red glow is observed behind the cathode.

                   It is now known that some rays are formed at the anode and when these rays strike the walls of the discharge tube they produce a faint red light. Since these rays are formed at the anode (positive electrode), they are known as anode rays or positive rays. The properties of anode rays have been studied by performing experiments similar to those performed on cathode rays.

Properties of Anode Rays

Characteristic properties of anode rays are :

1.    Anode rays travel in straight lines.
2.    Anode rays consist of material particles.
3.    Anode rays are deflected by electric field towards negatively charged plate. This indicated that they are positively charged.
4.    When a magnetic field is applied in a direction perpendicular to the path of anode rays, they get deflected in the direction expected for positive particles. This further indicates that they are positively charged.
5.  Charge to mass ratio of the particles in the anode rays depends upon nature of the gas taken in the discharge tube.

Proton

  1.    Mass of proton is found to be 1.673 × 10^–24g.

  2.   Charge on a proton is + 1.602 × 10^–19coulomb.

  3.   The specific charge of a proton is 9.58 × 10⁴coulomb/g. However, the specific charge of the anode rays is not constant. It varies from particle to particle in a discharge tube (containing any gas other than H₂) depending upon the number of electrons lost. Specific charge on anode ray particles also changes with the nature of the gas in the tube. It is because different gases have different atomic masses. It is maximum when gas present in the discharge tube is hydrogen.

  4.    Mass of 1 mole of proton is nearly 1.007 g.

  5.   Charge on 1 mole of proton is ≈ 96500 coulomb or 1 Faraday.

  6.    The volume of a proton is nearly 1.5 × 10^–38cm³.

Discovery of Neutron

            The discovery of neutron was actually made about 20 years after the structure of atom was elucidated by Rutherford. Atomic masses of different atoms could not be explained if it was accepted that atoms consisted only of protons and electrons. Thus, Rutherford (1920) suggested that in an atom, there must be present at least a third type of fundamental particle which should be electrically neutral and possess mass nearly equal to that of proton. He proposed the name for such fundamental particle as neutron. In 1932, Chadwick bombarded beryllium with a stream of a-particles. He observed that penetrating radiations were produced which were not affected by electric and magnetic field. These radiations consisted of neutral particles, which were called neutrons. The nuclear reaction can be shown as :

 

            The mass of the neutron was determined. It was 1.675 × 10^–24 g, i.e., nearly equal to the mass of proton.

Properties of Neutron

1. Neutron is slightly heavier (0.18%) than proton. Mass of neutron is 1.008665 amu or 1.675 × 10^–24g.
2. Specific charge of a neutron is zero.
3. Density of a neutron is 1.5× 10¹⁴g/cm³  
4. Mass of 1 mole of neutron is nearly 1.0087 g.
5. Of all the elementary particles present in an atom, neutron is the heaviest and least stable particle. Isolated neutron is unstable and disintegrates into electron, proton and neutrino.

Fundamental Particles of an Atom

 

Name	Mass	Charge	e/m
Electron (e)	9.1 × 10-31 kg or 5.5 × 104 amu	– 1.602 ×10 -19 C or – 4.8 × 1010 esu	1.76 × 108 C/g
Proton (p)	1.673 × 1027 kg or 1.007 amu	+ 1.602 × 10-19 C or + 4.8 × 10-10 esu	9.58 × 104 C/g
Neutron (n)	1.675 × 1027 kg 1.008 amu	Neutral	Zero

Thomson’s Model     

            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 electronic arrangement. Many different names are given to this model, for example, plum pudding, raisin pudding or watermelon. This model can be visualised as a pudding or watermelon of positive charge with plums or seeds (electrons) embedded into it. An important feature of this model is that the mass of the atom is assumed to be uniformly distributed over the atom. Although this model was able to explain the overall neutrality of the atom, but was not consistent with the results of later experiments. Thomson was awarded Nobel Prize for physics in 1906, for his theoretical and experimental investigations on the conduction of electricity by gases.

            Drawback :  It was unable to explain how positively charged particles are sided from negatively charged particle, without getting combined to neutralize each other.

Rutherford’s Experiment

            Alpha Particle Scattering Experiment :  Rutherford, in 1911, performed an experiment which led to the downfall of Thomson’s model. The experiment involved the bombardment of a thin sheet of gold (thickness ~ 100 nm or 10^–5 cm) by α -particles. These particles were obtained in the form of a narrow beam by passing through a slit in a lead plate. A circular fluorescent screen coated with zinc sulphide (ZnS) was placed around the foil to detect the deflection suffered by α-particles. Whenever an α-particle struck the screen, a tiny flash of light was produced at that point.

Figure : Setup of Rutherford’s scattering experiment

            Rutherford observed that

            (i)   Most of the α-particles (nearly 99%) passed through the gold foil undeflected.

            (ii)   Some of the α-particles  were deflected by small angles.

            (iii)  A very few α-particles (1 in 20000) were either deflected by very large angles or were actually reflected back along their path.

            These observations could not be explained by Thomson’s model. According to Thomson’s model, mass and charge in an atom is uniformly distributed throughout its volume. On the basis of this model it was expected that α-particles in passing through the foil would experience only a weak electric field and hence, they should suffer only slight deflections at the most.

            In order to explain the observations of α-particle scattering experiment, Rutherford assumed that the solid gold foil consists of layers of individual atoms which are touching each other so that there is hardly any empty space between them. As such, the α-particles striking the gold foil must pass through the atoms. Rutherford explained his observations as follows :

            (i)   Since most of the α-particles pass through the foil undeflected, it indicates that the most of the space in an atom is empty.

            (ii)  α-particles being positively charged and having considerable mass, could be deflected only by some heavy, positively charged centre. The small angle of deflection of α-particles indicated the presence of a heavy positive centre in the atom. Rutherford named this positive centre as nucleus.

            (iii) α-particles which make head-on collision with heavy positive centre are deflected through large angles. Since the number of such α-particles is very small, the space occupied by the heavy positive centre must be very small.

Rutherford’s Model of an Atom

           On the basis of scattering experiment, Rutherford put forward nuclear model of  an atom. Main points of this model are :

1. Most of the mass and all the positive charge of an atom is concentrated in a very small region called nucleus. Size of the nucleus is extremely small compared with the size of the atom. Radius of the nucleus is of the order of 10^–15m, whereas radius of atom is of the order of 10^–10 m.
2. The positive charge on the nucleus is due to protons. The magnitude of the charge on the nucleus is different for atoms of different elements.
3. The nucleus is surrounded by electrons which are revolving around it at very high speeds.
4. Total negative charge on the electrons is equal to the total positive charge on the nucleus, so that atom, on the whole, is electrically neutral.
5. Most of the space inside an atom is empty.

            Nuclear model of atom can be compared with the solar system. In an atom, electrons revolve around the nucleus in just the same way as the planets revolve around the sun. Due to this comparison, revolving electrons are sometimes called planetary electrons and Rutherford’s nuclear model of atom is known planetary model of atom.

Failure of Rutherford’s Model

            A major drawback (or defect) of Rutherford’s model of the atom is that it does not explain the stability of the atom. This point will become more clear from the following discussion :

            In the Rutherford’s model of an atom, the negatively charged electrons are revolving around the positively charged nucleus in circular paths. Now, we know that if an object moves in a circular path, then its motion is said to be accelerated. This means that the motion of an electron revolving around the nucleus is accelerated.

            According to the electromagnetic theory of physics, if a charged particle undergoes accelerated motion, then it must radiate energy (or lose energy) continuously.  Now, if we apply this electromagnetic theory to the Rutherford’s model of an atom, it will mean that the negatively charged electrons revolving around the nucleus with accelerated motion will lose their energy continuously by radiation. Thus, the energy of revolving electrons will decrease gradually and their speed will also go on decreasing. The electrons will then be attracted more strongly by the oppositely charged nucleus due to which they will come more and more close to the nucleus. And ultimately the electrons should fall into the nucleus by taking a spiral path (as shown in figure). This should make the atom very unstable and hence the atom should collapse.

            But this does not happen at all. We know that the electrons do not fall into the nucleus of an atom. Rather, atoms are very stable and do not collapse on their own. The Rutherford’s model, however, does not explain the stability of an atom.

            Rutherford did not gave any idea about the arrangement of electrons around the nucleus.

Figure : Electron following a spiral path & falling into the nucleus

Bohr’ s Atomic Model            

The important postulates on which Bohr’s model is based are the following :

(1)        The atom has a nucleus, where all the protons and neutrons are present. The size of the nucleus is very small. It is present at the centre of the atom.

(2)        Negatively charged electrons are revolving around the nucleus in the same way as the planets are revolving around the sun. The path of the electron is circular. The force of attraction between the nucleus and the electron is equal to centrifugal force of the moving electron.

Force of attraction towards nucleus = centrifugal force of the moving electron.

(3)        Out of infinite number of possible circular orbits around the nucleus, the electron can revolve only on those orbits, whose angular momentum is an integral multiple of  \frac{h}{{2\pi }} i.e.,  mvr = n \frac{h}{{2\pi }} where m = mass of the electron, v = velocity of electron, r = radius of the orbit and n = 1, 2, 3,…….number of the orbit. The angular momentum can have values such as \frac{h}{{2\pi }},\frac{{2h}}{{2\pi }} , \frac{{3h}}{{2\pi }}etc., but it cannot have a fractional value. Thus, the angular momentum is quantized. The specified or circular orbits (quantized) are called stationary orbits.

(4)        By the time, the electron remains in any one of the stationary orbits, it does not lose energy. Such a state is called ground or normal state.

(5)        Each stationary orbit is associated with a definite amount of energy. The greater the distance of the orbit from the nucleus, more shall be the energy associated with it. These orbits are also called energy levels and are numbered as 1, 2 3, 4,…………or K, L, M, N,………………..from nucleus outwards.

            i.e., E₁ < E₂ < E₃ < E₄…………………

            (E₂ – E₁) > (E₃ – E₂) > (E₄ – E₃)…………..

(6)        The emission or absorption of energy in the form of radiation can only occur when an electron jumps from one stationary orbit to another.

            ΔE = E_high – E_low = h

            Energy is absorbed when the electron jumps from inner to outer orbit and is emitted when it moves from outer to an inner orbit.

            When the electron moves from inner to outer orbit by absorbing definite amount of energy , the new state of the electron is said to be excited state.

            Using the above postulates, Bohr calculated the radii of various stationary orbits, the energy associated with each orbit and explained the spectrum of hydrogen atom.

Atomic Number

            The number of protons in an atom is equal to the number of electrons, since atom on the whole, is electrically neutral. Thus, atomic number of an element is equal to the number of protons present in the nucleus or the number of electrons present outside the nucleus when it is neutral. For example, number of protons in hydrogen atom and carbon atom are 1 and 6 respectively. So, their atomic numbers are 1 and 6 respectively. It is generally denoted by the letter Z. Thus,

                                    Atomic number (Z)  = Nuclear charge or number of protons (p)

                                                               = Number of electrons (e)

Mass Number

            Nucleus consists of protons and neutrons and these are collectively known as nucleons. Since the electrons are of negligible mass, the entire mass of the atom is due to the nucleus i.e., nucleons. The sum of the neutrons and protons is known as mass number.

                                    Mass number = No. of protons + No. of neutrons

            Mass number is generally represented by the letter A.

            Therefore, from the knowledge of atomic number and mass number of an element, the number of electrons, protons and neutrons can be easily predicted. We know

                                                Atomic number (Z)  = No. of protons (p)

                                                                           = No. of electrons (e)

                                                Mass number (A) = No. of protons (p) + No. of neutrons (n)

            Therefore, for an atom with mass number A and atomic number Z:

            Number of electrons = Z, Number of protons = Z

            Number of neutrons = A – Z

            For example, lithium has an atomic number (Z) = 3 and mass number (A) = 7. Therefore,

            Number of electrons = Atomic number = 3, Number of protons = Atomic number = 3

            Number of neutrons = Mass number – Atomic number = A – Z = 7 – 3 = 4.

Isotopes           

What are Isotopes ?

            Isotopes may be defined as follows :

            Atoms of the same element, having the same atomic number, but different mass numbers are called isotopes of that element.

            Since all isotopes of an element have the same atomic number, all the isotopes should contain the same number of protons inside their nuclei. Also, since different isotopes of an element have different mass numbers, the number of neutrons in the nuclei of isotopes of an element should be different. So, isotopes may also be defined as follows :

            The atoms of the same element which have the same number of protons but different number of neutrons inside their nuclei are called isotopes of that element.

            Isotopes are described by writing the mass number of that isotope as a superscript on the left side (top-left side) of the symbol of the element. The atomic number is written as a subscript on the left side (bottom-left side) of the symbol. For example, an isotope of the element X is described as,

                                          {}_{Atomic\,\,number}^{Mass\,\,number}\,X  or  {}_Z^A\,X

Isotopes of hydrogen

            Hydrogen (H) has three isotopes having mass numbers 1, 2 and 3, but all having atomic number equal to 1. These three isotopes of hydrogen can be described as follows :

{}_1^1\,H      {}_1^2\,H          {}_1^3\,H    

                                                   1 electron           1 electron          1 electron

                                                   1 proton             1 proton            1 proton

                                                   0 neutron            1 neutron           2 neutrons

                                                   (Protium)            (Deuterium, D)   (Tritium, T)

What are the characteristics of Isotopes :

            (i)   The isotopes of an element have the same number of protons inside their nuclei. As a result, all the isotopes of an element contain the same number of electrons.

            (ii)  Different isotopes of an element have different mass number.  So, isotopes show the following characteristics :

            (a) Since the isotopes of an element have the same number of protons and electrons, all the isotopes of an element show the same chemical properties, same electronic configuration, and the same number of valence electrons.

            (b)  The isotopes of an element have different masses. So, the properties which depend upon atomic mass should be different for different isotopes. Many physical properties, e.g., melting point, boiling point, density, etc., depend upon atomic mass. So, different isotopes of an element show different physical properties.

            Application of radioactive isotopes

            Isotopes or radioisotopes are useful in many ways. Some main uses of radioactive isotopes (called radioisotopes) are :

For estimating the age of old archaeological samples

            Radiocarbon dating, or in general, radioisotopic dating method is used for estimating the age of old archaeological samples. For example, age of the earth, moon, rocks and mineral deposits can be determined by using the principle of radioisotopic dating.

For the treatment of diseases :

            Radioisotopes are widely used for the treatment of diseases like cancer.

• The isotope, cobalt-60 \left( {{}_{27}^{60}Co} \right) is used for destroying malignant cells in patients suffering from cancer.
• The isotope, iodine-131 \left( {{}_{53}^{131}I} \right) is used for studying disorders of the thyroid gland.
• The isotope, sodium-24 \left( {{}_{11}^{24}\,Na} \right) is used for examining the circulation of blood.

For estimating the age of glaciers

            The age of glaciers, snow-field, and even wines can be estimated by radioisotopic dating. In these cases, the radioactivity level of tritium (an isotope of hydrogen having mass number of 3, {}_1^3\,H) is measured.

What is radiocarbon dating

            This technique was developed by Willard Libby. He was awarded the Nobel Prize for this work.

            Estimating the age of a carbon-containing object by measuring the concentration (or activity) of {}_6^{14}\,C  in it is called radiocarbon dating.

            Radioactive carbon-14 \left( {{}_6^{14}\,C} \right) gets converted to radioactive carbon dioxide ^{14}C{O_2}. This radioactive ^{14}C{O_2} is taken up by plants during photosynthesis. {}_6^{14}\,C is radioactive and decays by\beta  - emission.

            {}_6^{14}\,C is being continuously formed in the higher atmosphere and consumed due to \beta  - emission decay. As a result, an equilibrium concentration of  {}_6^{14}\,C is maintained in all the living plants. However, when a plant dies, it can no longer fix up radioactive ^{14}C{O_2}. As a result, the concentration of {}_6^{14}\,C in it starts decreasing. The half-life of {}_6^{14}\,C  is 5760 years. Thus, in 5760 years, the concentration of {}_6^{14}\,C  is lowered to half (50%) of its initial concentration, and another 5760 years, its concentration gets lowered to 25% (50% of the 50%) of the initial concentration. Thus, in 11,520 years, the {}_6^{14}\,C concentration is reduced to one-fourth of its initial concentration. Thus, by measuring the concentration of {}_6^{14}\,C in a dead carbon-containing object, and knowing the concentration of {}_6^{14}\,C in a living plant, we can estimate the age of the object (the age of the object means the number of years ago when the plant would have died).

Isobars

            The word ‘isobar’ has been derived from the Greek word meaning ‘equally heavy’ (isos = equal, barys
= heavy).

            The atoms of different elements which have the same mass number but different atomic number are called isobars.

            As the mass number is equal to the sum of the number of protons and neutrons inside the nucleus of an atom, therefore,

            isobars may also be defined as follows :

            The atoms of different elements which have different number of protons but equal sum of the number of protons and neutrons are called isobars.

            {}_{18}^{40}Ar,\,\,{}_{19}^{40}K\,\,\,and\,\,\,\,{}_{20}^{40}Ca are some typical examples of isobars. Each of these have the same mass number but different atomic number. The nuclear composition of these isobars is as follows :

Isotones

            Atoms of different elements having same number of neutrons but different mass numbers are called isotones.

            For example, {}_{14}^{30}Si (14 protons, 16 neutrons), {}_{15}^{31}P (15 protons, 16 neutrons) and {}_{16}^{32}S (16 protons, 16 neutrons) are isotones because all have 16 neutrons.

Quantum Numbers                 

A large number of electron orbitals are possible in an atom. These can be distinguished by their size, shape and orientation.

                  To describe each electron in an atom in different orbitals, we need a set of three numbers known as quantum numbers. These are designated as n, and m. In addition to these three members, another quantum number is also needed which specify the spin of the electron. These four numbers are called quantum numbers.

                  Four quantum numbers are :

            (i)    Principal quantum number (n)

            (ii)   Azimuthal quantum number ()

            (iii)  Magnetic quantum number (m)

            (iv)  Spin quantum number (s)

(i)         Principal quantum number (n)

• This quantum number determines the main energy shell or shell  in which the electron is present.
• It is denoted by n.
• n can have whole number values starting from 1 such as n = 1, 2,3,4…………….
• This quantum number is also regarded as shell or orbit. The shell with n = 1 is called first shell.

 

• This quantum number also helps to determine the average distance of the electron from the nucleus. As the value of n increases, the distance of the electron from the nucleus and also the energy of that shell increases.

(ii)        Azimuthal quantum number ()

• This quantum number determines the sub energy level or sub shell in which the electron is present.
• This quantum is denoted by
• This quantum number is also known as angular momentum quantum number or subsidary quantum number.
• Sub shells are s, p, d, f.
• Maximum number of  electrons that can occupy in these sub shells are :

• Value of l is from zero to (n – 1)

                  Where n = principal quantum number

(iii)       Magnetic quantum number (m)

• This quantum number refers to the different orientations of electron cloud in a particular sub shell. These different orientations are called orbitals.
• The number of orbitals in a particular sub shell within a principal energy level is given by the values allow to m, which in turns depends on the values of .
• The possible values of m ranges from +  through 0 to – , thus making a total of (2 + 1) values. Thus in a sub shell, the number of orbitals is equal to (2 + 1)

(iv)       Spin quantum number

• This quantum number determines the spin orientation of the electron.
• This quantum number is designated by s.
• The spin quantum number can have only two values as + \frac{1}{2} and - \frac{1}{2}
• The  + \frac{1}{2}  value indicates the clockwise spin.
• The - \frac{1}{2} value indicates the anticlockwise spin. 

Electronic Configuration                 

The filling of orbitals in an atom is a hypothetical process in which the atom is built up by feeding electrons in orbitals, one at a time and by placing each new electron in the lowest available energy orbital. The distribution of electrons in different orbitals is known as electronic configuration of the atom. For the sake of presentation, the following symbols are commonly used :

Figure : Presentation of one and two electrons in an orbital

                  A box for orbital (square or circular); an arrow for an electron, the direction of the arrow giving the orientation of its spin. Two arrows are shown for two electrons with opposite orientations of spin.

                  Alternatively, electronic configuration is expressed by indicating the principal quantum number and its respective orbital along with the number of electrons present in it. For example, the notation 3p_x indicates that there is one electron in p_x orbital of third principal shell.

Figure : Representation of electrons

            The filling of orbitals is governed by the following principles :

1. Aufbau principle : The Aufbau principle states that in the ground state of an atom, an electron enters the orbital of lowest energy first and subsequent electrons are fed in the order of increasing energies. The word ‘aufbau’ is German word and it means ‘building up’. The building up of the orbitals means the filling up of orbitals with electron. From an energy level diagram of multi-electron atoms, the following sequence is observed for orbitals in the increasing order of energy :

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

            According to Aufbau principle, the orbital should be filled in the above sequence.

            It is very important to remember that the sequence of energy levels pertains up to 3p and then 4s-orbital comes first instead of 3d. In fact, the energy of an orbital is determined by the quantum numbers n and with the help of important rule known as (n + ) rule or Bohr Bury’s rule. According to this :

            (i)    Orbitals fill in the order of increasing value of n + . For example, 3s-orbital (n + = 3 + 0 = 3) will be filled before 4s (n + = 4 + 0 = 4) orbital. Similarly, out of 3d and 4s, the 4s (n + = 4 + 0 = 4) orbital will be filled before 3d
 (n + = 3 + 2 = 5) orbital.

            (ii)   If the two orbitals have same value of (n + ), then the orbital with lower value of n will be filled first. For example, 2p-orbital (n + = 2 + 1 =3) and 3s-orbital (n + = 3 + 0 =3) have the same (n + ) value but 2p-orbital has lower value of ‘n’ and therefore, it will be filled first.

                  This rule also helps to account for the fact that certain orbitals with higher value of n but lower value of have less energy than orbital with lower value of . For example, let us apply this rule to 4s- and 3d-orbitals. For 3d-orbital, (n + ) value = 3 + 2 = 5 while of  4s-orbital, (n + ) value = 4 + 0 = 4. Thus, 4s-orbital has lesser energy than 3d-orbital and, therefore, is filled first.

        The sequence of energy levels can be easily remembered by the systematic diagram as shown in Figure 

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

Figure : Sequence of filling of electrons in different energy levels

2. Pauli’s exclusion principle

            According to this principle, an orbital can accommodate maximum of two electrons.

            This principle,  thus, limits the accommodation of electrons in an orbital. This means that an orbital can have 0,1 or 2 electrons. Moreover, if an orbital has two electrons, they must be of opposite spins.

3. Hund’s rule of maximum multiplicity

            According to this rule, electron pairing will not take place in orbitals of same energy (same sub shell) until each orbital is singly filled. This suggests that it is difficult for an electron to enter an orbital which already has an electron than to enter an unoccupied orbital of same energy. This principle is very important in guiding the filling of p, d and f orbitals, which have more than one kind of orbitals. For example, we know that there are three p-orbitals (p_x, p_y and p_z) of the p-sub shell in a principal energy level. According to Hund’s rule, each of the three p-orbitals must get one electron of parallel spin before any one of them receives the second electron of opposite spin.

Electronic Configuration of Some Elements

            Based upon the above rules and the sequence of energy levels, let us write the electronic configurations of some elements.

     Hydrogen (At. No. = 1)

            Since hydrogen has only one electron, it must go to 1s-orbital which has lowest energy.  

    Helium (At. No. = 2)

            In helium atom, the second electron can also go into 1s-orbital. The two electrons must have opposite spins (Pauli’s exclusion principle).

      Lithium  (At. No. = 3)

            Since 1s-orbital is filled with two electrons, it cannot have any more electrons. Therefore, the third electron goes to the next lowest energy orbital, namely 2s-orbital.

                                                            Li : (Z = 3) 1s² 2s¹          or                                     

      Beryllium (At No = 4)

            In Berylium atom the fourth electron goes into 2s orbital and it get completely filled.

                                                            Be : (Z = 4) 1s²  2s²       or                                                  

     Boron  (At . no = 5)

            Now the fifth electron in Boron atom comes into 2p shell because after 2s – orbital, 2p comes.

1. (Z = 5) 1s²2s² 2p¹_x  or                                         

       Carbon (At. No. = 6)

            In carbon atom, the sixth electron also goes into the 2p-orbitals because it can accommodate six electrons. Here Hund’s rule applies, i.e., the electrons enter the orbitals of same energy with parallel spin until all are singly filled. Therefore, the sixth electron cannot enter the 2p_x orbital, rather it can go into either 2p_y or 2p_z in accordance with Hund’s rule :

C : (Z = 6) 1s² 2s² 2p_x¹ 2p_y¹        or 

      

      Nitrogen (At. No. = 7)

     Applying Hund’s rule, nitrogen atom has three unpaired electrons in 2p-orbitals as :                                            

N : (Z = 7) 1s² 2s² 2p_x¹ 2p_y¹ 2p_z¹  or                                                            

        Oxygen (Z = 8), Fluorine (Z = 9) and Neon (Z = 10)

    Beginning with oxygen, the 2p-orbitals start getting filled by second electron till each of these is completely filled.

Ne : (Z = 10) 1s² 2s² 2p_x² 2p_y² 2p_z²   or                                                          

Valency and Valence Electrons

            The electrons present in the outer most shell of an atom are known as valence electrons because they decide the valency (combining capacity) of the atom.

            Valence electrons take part in chemical reaction.

            In order to find out the number of valence electrons in atom of the element, we should write down the electronic configuration of the element by using its atomic number. The outermost shell will be the valence shell and the number of electrons present in it will give the number of valence electrons.

            For example,

                        Element X has atomic number = 11

                                                                        K          L          M

                       Electronic configuration =          2          8          1

            Here M shell is the outermost shell or valence shell of the atom and it has 1 electron in it. Therefore, number of valence electron in X is one.

Inertness of Noble Gases

            There are some elements which do not combine with other elements. These elements are :

            Helium, Neon, Argon, Krypton, Xenon and Radon. They are known as noble gases or inert gases because they do not react with other elements to form compounds. Since the noble gases are chemically unreactive, we must conclude that the electron arrangement in their atoms are very stable, which do not allow the outermost electrons to take part in chemical reaction.

            All the noble gases have completely filled outermost shells. The atoms having 8 electrons in their outermost shell, are very stable and hence, chemically unreactive.

            Note :  2 electrons in the outermost shell is considered to be a stable arrangement of electrons only when the atom has just one shell, K shell and there are no other electron shell in the atom.

            Since the atoms of inert gases are very stable or unreactive, they can exist in the free state as individual atoms. So, the inert gases are monoatomic. For example : Helium, Neon, Argon etc. exist in the form of monoatomic molecules, He, Ne, Ar, etc.

Cause of Chemical Combination

            The atoms combine with one another to achieve the inert gas electron arrangement and become more stable.

            An atom can achieve the inert gas electron arrangement in three ways.

            (i)    by losing one or more electrons (to another atom).

            (ii)   by gaining one or more electrons (from another atom).

            (iii)  by sharing one or more electrons.

Relation Between Valency And Valence Electrons

            (i)    The valency of a metal element is equal to the number of valence electrons in its atom.

                  That is :     

                                    Valency of a metal = Number of Valence Electrons in its Atom

            (ii)   Valency of a non metal element is usually equal to eight minus the number of valence electrons in its atom.

                  That is :

                                    Valency of a non-metal = 8 –Number of Valence Electrons in its Atom

Types of Valency

            There are two types of valency :

            (i)    Electrovalency        (ii)         Covalency

1. Electrovalency : The number of electrons lost or gained by one atom of an element to achieve the nearest inert gas electronic configuration is known as its electrovalency.

            For example :

            Atomic number of Mg = 12    Electronic configuration = 2, 8, 2

            Since one Mg atom loses 2 electrons to achieve the inert gas electronic configuration, therefore, the  electrovalency of Mg = + 2 or (2).

2. Covalency : The number of electrons shared by one atom of an element to achieve the nearest inert gas electronic configuration is known as its covalency.

                        For example : Atomic number of Nitrogen = 7

                        Electronic Configuration = 2, 5

3. Since one nitrogen atom shares 3 electrons to achieve the inert gas configuration, therefore, the covalency of nitrogen is 3.

 

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