ELECTRONIC CONFIGURATIONS OF ATOMS
In order to represent electron population of an orbital, the principal quantum number (n) is written before the orbital symbol while the number of electrons in the orbital is written superscript near the right hand top of the orbital symbol. For example, if we have two electrons in the s-orbital of first energy level then it is written as ls2 Sometimes, electronic configurations are represented in a different manner by representing each orbital by a square or a circle and the electrons are represented by putting arrows in it as illustrated below:
ELECTRONIC CONFIGURATIONS OF ATOMS
On the basis of above rules and the sequence of energy levels, let us write electronic configurations of some elements.
Hydrogen (At. No.= 1). Since hydrogen has only one electron, it must go to ls orbital. The electronic configuration for hydrogen is
Boron (At. No. = 5). In this case, the four electrons completely fill l s and 2s orbitals. The fifth electron goes into one of the 2p orbitals.
Carbon (At. No. = 6). In case of carbon, the sixth electron also goes into 2p sub-shell, but the two electrons in 2p subshell are present in different orbitals because according to Hund’s rule pairing of electrons cannot take place in orbitals of a particular sub-shell until all the orbitals get one electron each. Hence, electronic configuration of carbon is
Nitrogen (At. No. = 7). Applying Hund’s rule, N atom has three unpaired electrons in 2p orbitals.
Oxygen, Fluorine and Neon. Beginning with oxygen, the 2p orbitals start filling by second electron till neon in which it is completely filled From sodium (At. No. 11) to argon (At. No. 18) 3s and 3p orbitals are successively filled. After 3p, the· 19th electron in potassium (At. No. 19) enters the 4s orbital instead of 3d because energy of 4s orbital is less than that of 3d. Electronic configurations of first twenty elements have been given in elements are known as the transition elements. The electronic Table 5.5. configuration of scandium (At. No. = 21), the first transition
The next ten elements (At. Nos.= 21 – 30) which follow element, for example, is Lil 2$2 2p6 3$2 3p6 4s2 3d1 while that calcium are scandium (Sc), titanium (Ti), vanadium (V), of zinc (At. No.= 30), the last element of the series, is ls2 2$2 chromium (Cr), manganese (Mn), iron (Fe), cobalt (Co), 2p6 3$2 3p6 4s2 3d10. Thus, while the fllling of 3d-orbitals begins nickel (Ni), copper (Cu) and zinc (Zn). In these elements with scandium, it ends with zinc. The electronic configurations of addition of electrons rakes place in the 3d-orbitals. All these all these 10 elements are represented in Table 5.6.
Table 5.5. Electronic Configurations of the First 20 Elements
Table 5.6. Electronic Configurations of Transition Elements
[Scandium (At. No. 21) to Zinc (At. No. 30)]
EXCEPTIONAL CONFIGURATIONS OF CHROMIUM AND COPPER
From the electronic configurations in Table 5.6, it may be noted that chromium and copper have five and ten electrons in 3d orbitals rather than four and nine electrons respectively as expected. The reason is that folly filled orbitals and exactly half-filled orbitals have extra stability. Thus, the p3, p6, c6,d10 andf14 configurations which are either fully filled or exactly half-filled, are more stable. Therefore, to acquire more stability one of the 4s electron goes into 3d orbitals so that 3d orbitals get half-filled or completely filled in chromium and copper respectively.
Expected Configuration: ls2, 2s2 2ji, 3s2, 3p6, 3d”, 4s2
Actual Configuration: 1 s2, 2s2, 2p6, 3sl, 3p6, 3d5, 4s1
Expected Configuration: ls2, 2s2, 2p6, 3s2, 3p6, 3d9, 4?
Actual Configuration: 1 s2, 2s2, 2p6, 3sl, 3p6, 3d10, 4s1
Explanation of Greater Stability of Half-filled and Fully-filled Electronic Configurations
The extra stability of half-filled and fully-filled electronic configurations can be explained in terms of symmetry and exchange energy. The half-filled and fully-filled electronic configurations have symmetrical distribution of electrons and this symmetry leads· to stability. Moreover, in such configurations electrons can exchange their positions among themselves ·to maximum extent. This exchange leads to stabilization which is expressed in terms of exchange energy.
The half-filled and completely filled degenerate orbitals provide .,extra stability to the system due to symmetrical arrangement and exchange energy as described below:
(i) Symmetrical arrangement. The electronic configurations in which all the orbitals of same sub-shell are either completely-filled or exactly half-filled have relatively more symmetrical distribution of electrons and therefore, lend more stability to the system. For example, expected configuration of chromium is 3tf 4s2. But shifting of one electron from 4s to 3d-orbitals makes the configuration more symmetrical and hence relatively more stable.
In the similar way shifting of one electron from 4s to 3d in copper also makes the configuration relatively more stable.
(ii) Stability due to exchange energy. The half-filled or fully-filled degenerate orbitals have more number of exchanges and consequently, have large exchange energy of stabilisation. The exchange means shifting of electrons from one orbital to another within same sub-shell. Let us compare the number of exchanges in 3tf 4s2 and 3cP 4s1 configurations of chromium.
Total exchanges = 3 + 2 + 1 = 6
Thus, in 3d4 arrangement electron exchanges are six which implies that there are six possible arrangements with parallel spin in 3cf configuration.
Total exchanges = 4 + 3 + 2 + 1 = 10
It is quite evident from the above description that total number of electron exchanges in 3c5 arrangement is larger which lends it relatively greater stability. In the similar way it can be established that number of .exchanges in 3d10 configuration is larger than in 3d9 configuration which makes 3d10 configuration relatively more stable.
ELECTRONIC CONFIGURATIONS OF IONS
Ions are formed by loss or gain of electrons by neutral atoms. A positive ion (or cation) is formed by loss of electrons whereas negative ion (or ariion) is formed by gain of electrons. The electronic configurations of ions are written just like the electronic configurations of atoms. However, it must be noted that while forming cations the electrons are lost from the shell with highest value of principal quantum number. For example
Fe : ls2, 2s2, 2p6, 3s2, 3p6, 3d6, 4s2
Fe2+ : ls2, 2s2, 2p6, 3s2, 3p6, 3d6
Fe3+ : ls2, 2s2, 2p6, 3s2, 3p6, 3d5
Zn : ls2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2
Zn2+: l s2, 2s2, 2p6, 3s2, 3p6, 3 d10
The anions are formed by adding electrons to the vacant orbital of lowest energy.
F ls2, 2s2, 2p5
F- ls2, 2s2, 2p6
O ls2, 2s2 2p4
O2-: ls2, 2s2 2p6
Example 5.14 An atom has 2K, 8L, 5M electrons. Write down its electronic configuration and indicate in it
(i) number of sub-shells
(ii) number of orbitals
(iii) number of unpaired electrons
(iv) number of electrons having l =1
Solution. In the given atom there are (2 + 8 + 5) =1 5 electrons. The electronic configuration of the atom is
1 s2, 2s2, 2p6, 3px1 , 3py1, 3pz1
(i) Number of subshells is 5
(ii) Number of orbitals is 9
(iii) Number of unpaired electrons is 3
(iv) Number of electrons having l = 1 is 9.
Example 5.15 An atom of an element contains 29 electrons and 35 neutrons. Deduce (i) the number of protons and (ii) the electronic configuration of the element.
Solution. (i) Number of protons = Number of electrons = 29
(ii) Electronic configuration of the element is
2, 2SZ, 2p6, 3SZ, 3p6, 3d10, 4s1
Example 1 h. Write electronic configurations of Fe2+ and fe3+ ions. Which of these has more number of unpaired electrons? Atomic no. of Fe is 26.
Solution. Fe atom has 26 electrons. Fe2+ and Fe3+ ions are formed by removal of two and three electrons respectively from Fe atom. Their electronic configurations are:
Fe2+ : 1s2, 2s2, 2p6, 3SZ, 3p6, 3d6
Fe3+: 1s2, 2SZ, 2p6, 3s2, 3p6, 3d5
Fe2+ has 4 unpaired electrons while Fe3+ has 5 unpaired electrons. Therefore, Fe3+ has more number of unpaired electrons.
WRITING THE ELECTRONIC CONFIGURATION IN CONDENSED FORM
It may be noted that configurations of atoms can also be written in condensed form by taking the configurations of noble gases as the core. The configurations of inert gases representing core are written as [He]2, [Ne]10, [Ar]18, [Kr]36, [Xe]54 and [Rn]86. For example, electronic configurations of scandium having atomic number 21 may be written as:
21Sc: [Ar)18, 3d1, 4s2.
I. Objective Type Questions
Select the most appropriate choice from the options given as
(a), (b), (c) and (d) after each question:
1. 'Plum pudding' model of atom was proposed by
(a) Rutherford (b) Dalton
(c) Bohr (d) J. J. Thomson.
2. An isotope of magnesium is represented as 26 /12 Mg . The number of nucleons in the isotope is
(a) 12 (b) 14
(c) 26 (d) 38.
3. Naturally occurring gold does not have an isotope. The number of peaks in the mass spectrum of naturally occurring gold would be
(a) 3 (b) 2
(c) 1 (d) 4.
4. The credit of discovering neutron goes to
(a) Rutherford (b) Thomson
(c) Goldstein (d) Chadwick.
5. The protons and neutrons are collectively called
(a) deuterons (b) positrons
(c) mesons (d) nucleons.
6. The ratio of the radius of the atom to the radius of the nucleus is of the order of
(a) 104 (b) 106
(c) w-5 (d) w-8
7. 13 /6 C and 12 /6 c differ from each in respect of number of
(a) electrons (b) protons
(c) neutrons (d) none of these.
8. Rutherford experiment which established the nuclear model of the atom used a beam of
(a) 13-particles which impinged on a metal foil and got absorbed
(b) y-rays which impinged on a metal foil and ejected electrons
(c) helium atoms. which impinged on a metal foil and got scattered
(d) helium nuclei which impinged on a metal foil and got scattered.
9. The charge to mass (elm) ratio of proton is
(a) 9.55 x w-4 C/g (b) 9.55 x 104 C/g
(c) 1.76 x 108 C/g (d) 1.76 x 10-8 C/g.
10. Isotopes of an element have
(a) similar chemical properties but different physical propetties
(b) similar chemical and physical properties
(c) similar physical properties but different chemical prope1ties
(d) different chemical and physical properties.
11. For a given principal level n= 4, the energy of its sub-shells is in the order
(a) s<p< d<f (b)s>p>d>f
(c)s<p<f<d (d) f<p<d<s.
12. The electrons, identified by quantum numbers and l, (i)n = 4 , l = 1 (ii) n= 4, l= 0 (iii) n = 3, l = 2 (iv) n= 3, l= 1 can be placed in order of increasing energy, from the lowest to highest, as
(a) (iv) < (ii) <(iii)< (i) (b) (ii) < (iv) < (z) <(iii)
(c) (i) <(iii)< (iv) < (iv) (d) (iii)< (i) < (iv) < (ii).
13. A sub-shell with n= 6 , l = 2 can accommodate a maximum of
(a) 12 electrons (b) 36 electrons
(c) 10 electrons (d) 72 electrons.
14. Which of the following sets of quantum numbers is not possible?
(a) n= 2, l= 1, m=1 - 1, ms =- 1/ 2
(b) n= 2 l= 0 m= 0 m= + 1/ 2
(c) n= 3, l= 2, m=1 - 2, ms = +1/2
(d)n=3 1=2 m= -3m=+1/ 2 ·
15 The correct configuration of 29Cu is
(a) [Ar] 4s1 (b) [Ar] 4s2
(c) [Ar] 3d104s1 (d) [Ar] 3d6 4s2.
16. The atomic orbitals are progressively filled in order of increasing energy. This principle is called
(a) Hund’s rule (b) Aufbau principle
(c) Exclusion principle (d) de-Broglie rule.
17. According to Aufbau principle, the 19th electron in an atom goes into the
(a) 4s-orbital (b) 3d-orbital
(c) 4p-orbital (d) 3p-orbital.
18. In a multi-electron atom, the energy of the electron in a particular orbital is determined by
(a) n only (b) l only
(c) n and l (d) n. land m r
19 The orbital diagram in which both Pauli’s exclusion principle and Hund’s rule are violated is
20 Which one of the following pairs of ions have the same electronic configuration?
(a) CI3+, Fe3+ (b) Fe3+, Mn2+
(c) Fe3+, Co3+ (d) Sc3+, Cr3+.
21. The number of d-electrons retained in Fe3+ (At. no. of Fe= 26) ion is
(a)6 (b) 3
(c) 4 (d) 5.
22 .Consider the following sets of quantum numbers:
n l m1 ms
(z) 3 0 0 + 1/2
(ii) 2 2 1 + 1/2
(iii) 4 0 -2 - 1/12
(iv) 1 0 -1 -1/2
(v) 3 2 3 + 112
Which of the following sets of quantum numbers is not possible?
(a) (i) and (iii) (b) (ii), (iii) and (iv)
(c) (i), (ii), (iii) and (iv) (d) (iv), (iii), (iv) and (v).
23. Which of the following sets of quantum numbers represents the highest energy of an atom?
(a) n = 3, l = 1, m=1 1, m=s + 112
(b) n= 3, l = 2, m1 = 1, m=s + 112
(c) n = 4, l= 0 , m=1 0 , ms = + 1/2
(d) n = 3, l= 0 , m=1 0, ms = + 1/2.
24 The electrons of the same orbitals can be distinguished by
(a) principal quantum number
(b) azimuthal quantum number
(c) spin quantum number
(d) magnetic quantum number.
15. A body of mass 10 mg is moving with a velocity of 100m s-1. The wavelength of the de-Broglie wave associated with it would be
(a) 6.63 x l0-37 m (b) 6.63 x 10-31 m
(c) 6.63 x l0-34 m (d) 6.63 x 10-35 m.
26. Which of the following sets of the quantum numbers is permitted?
(a) n = 4, l = 2, m1 = + 3, ms = + -1/2
(b) 1 n = 3,l = 3, m1 = + 3, ms = + 1/2
(c) n =4 , l =0 , m1 = 0, ms = +1/2
(d) n = 4, l = 3, m1 = + 1, ms = 0.
27. The maximum number of electrons that can be accommodated
in fifth energy level is
(a) 10 (b) 25
(c) 5 0 (d) 32.
U. Fill in the Blanks———–
28. Complete the following sentences by supplying appropriate words:
(i) The radius of atoms is of the order of …… m.
(ii) The neutral sub-atomic particles are …… .
(iii) In the atom represented as 23 /11 na, the number of neutrons is …… .
(iv) Isotopes have same …… number.
(v) A naturally occurring element which has two isotopes will show …… peaks in its mass spectrum.
(vi) Alpha particle contains …… protons.
(vii) s-orbitals are …… in shape.
(viii) The orbitals having equal energy are called …… orbitals.
( ix) The number of unpaired electrons in an atom of Mn (Z = 25) is …… .
(x) Specific charge of electron is …… .
III. Discussion Questions
29. Discuss Dalton’s atomic theory. What are its limitations?
30. What are cathode rays? Give four characteristics of cathode rays.
31. How will you demonstrate that cathode rays carry negative charge?
32. What is the charge and charge to mass (elm) ratio of particles in cathode rays?
33. How will you demonstrate with the help of an experiment that:
(i) Cathode rays travel in straight lines
(ii) Cathode rays consist of material particles.
34. The charge to mass ratio for an electron is 1.76 x 108 coulombs per gram. The charge On the electron is 1.602x l0-!9 coulombs. Calculate mass of the electron.
35. If an atom contains one electron and one proton, will it carry any charge or not?
36. Name the three subatomic particles of an atom.
37. On the basis of Thomson’s model of an atom explain how the atom is neutral as a whole.
38. Compare the properties of the three fundamental sub-particles of an atom.
39. What were the main observations of Rutherford’s scattering experiment?
40. What is a proton? How does it differ from a neutron?
41. Give experimental evidence to show that:
(i) Most of the space inside the atom is empty.
(ii) An atom contains heavy positively charge centre.
(iii) Size of the nucleus is very small.
(iv) Nucleus of an atom is positively charged.
42. What experiment established the presence of atomic nucleus? What features of nucleus were deduced from this experiment?
43. Describe the essential features of the model of atom proposed by E. Rutherford. How is it different from that proposed by J.J. Thomson?
44. Discuss the working of mass spectrometer. How does it help us in determining the number of isotopes and relative atomic mass of an element?
45. Describe the significance of each term in the symbol