Electronic Configuration, Chemistry tutorial

INTRODUCTION

The periodic law gives increase to the periodic arrangement of the element according to their atomic numbers that certainly means according to the number of electrons in their orbital. In the next chapter we will be learning the distributions of the electrons inside the atom ("the electronic configuration"), and how they administrate the properties of the elements.

The electronic configurations of isolated atoms of elements are frequently confirmed experimentally by an exhaustive analysis of atomic spectra.

Electronic configuration of atoms :

Rules governing the filling of electrons in orbital

The electronic configuration of atoms can be expected with the help of AUFBAU and the building up procedure. In the aufbau process, it is assumed that there exist a set of empty hydrogen like orbital around the nucleus of an atom. The electronic configuration of the atom in the ground state is then calculated by adding electrons one at a time to the orbitals of the lowest energy in the progression revealed by arrows.

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Fig: Order or filling of atomic orbital in polyelectronic atoms

The order in that the orbital are sealed is governed by the n +1 rule. According to this rule, in the structure up of electronic configuration of the essentials, the sub-shell with the value of n +1 fills first. This rule reminds us which the energy of sub shells of multi electron atoms depends upon the price of either the quantum numbers n or 1, but mainly on the value of n. For example, to determine which of the sub shells 5s or 4.p fills 1st. We give the rule therefore: for the 5s sub shell the value n + 1 = 5 + 0 = 5; for 4p sub shell as well the value of n + 1 = 4 + 1 = 5, but the 4p sub shell has the lower value of the principal quantum number n and thus it fills first. Filling of electrons in orbital is also governed by Hund's Rule and Pauli's Exclusion Principle.

According to the Exclusion Principle, no two electrons in similar atom can have the same value of n, 1 and m; they will differ in their m s values. In order words, an orbital can contain at the most two electrons of opposite spin. Thus there is only one orbital for any given value of n; it can since only two electrons. Thus, the three p orbital for any follow value of n can contain six electrons, the five d orbitals, for any given value of n can contain a total of ten electrons and the seven f orbital's can include 14 electrons.

Hund's rule of maximum multiplicity states that, as far as feasible in a specified atom in the ground state, electrons in the similar sub shell will occupy different orbital's and will contain parallel spins. That means that when electrons are added to orbitals of the same energy such as three p orbitals or 5d orbitals, one electron will enter each and every of the available orbital, two electrons in separate orbitals feel less repulsion than two electrons paired in the same orbitals. For example, carbon in the ground state has the configuration 1S2 2S2 2Px12Py1 rather than 1S2   2S22Px2.

So far you have studied the rules governing the filling of electrons in the orbitals of atoms. We shall now assume the electrons configurations of all the elements in the periodic table, these are following:

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Table: Ground state electronic configuration of gaseous atoms electronic Configuration of all the Elements in the Periodic Table

Period 1 :

This is the minimum of all the periods of the table. Hydrogen (Z = 1) and helium (Z = 2) are the two elements belonging to the period. The electronic configuration of hydrogen and helium are 1s1 and 1s2 correspondingly. Therefore the 1s orbital which is the only orbital subsequent to n = 1 is completely filled. The 1s2 configuration of helium is mainly represented by [He]. So anytime we see [He] electron arrangement it represents 1s2.

Period 2:

This period have elements from lithium (Z = 3) to neon (Z = 10). In lithium and beryllium, the filling of 2S orbital takes place, then in the next six elements from boron to neon, the 2p orbitals are filled. Neon thus has the electronic configuration of [He] 2S22P6 that as was done in the case of He, is symbolized by [Ne]. (An electronic configuration of [Ne] means 1S2 2S2 2P6.  At this stage, the shell having n = 2 is complete.

Period 3:

Similar to period 2, this period also consists of 8 elements from sodium (Z = 11) to argon (Z = 18) these elements 3s and 3p orbitals are successively filled in the sequence just as was done in period 2, thus argon has the electronic configuration [Ne] 2S22P6 represented as [Ar] although the third principal shell (n = 3) can hold 10 more electron in 3d orbitals filling of 4s orbital takes place first since of its lower energy.

Period 4:

This period since 18 elements from potassium (Z = 19) to krypton (Z = 36). In k and Ca, the first two elements of this period, the successive electrons go into the 4s orbitals giving them the configuration [Ar] 4S1 and [Ar] 4S2 respectively. Then in the following 10 elements (Sc, Ti, V, Cr Mn, Fe, Co, Ni, Cu and Zr) filling of hitherto unoccupied 3d orbitals takes place. Therefore the electronic configuration of zinc becomes [Ar] 3d104S2.

Infrequently an electron from 4s orbitals is shifted out of revolve to the 3d orbitals due to higher permanence of half filled or completely filled orbitals , for instance, Cr (Z = 24) and Cu (Z = 29) contain the configuration [Ar] 3d54S1 and [Ar]3d10451 instead of the expected [Ar] 3d44S2 and [Ar] 3d94S2 correspondingly. After the 3d level is filled, in the next six elements of this period, which is Ga, Ge, As, se, Br and Kr, the 4p orbitals are gradually filled and Kr has the electronic configuration [Ar] 3d104S24P6 represented as [Kr].

Period 5:

The next 18 elements from rubidium (Z = 37) to Xenon (Z = 54) belong to this periods. In structure up of the atoms of these elements, 5s 4d and 5p orbital are sequentially filled just as the 4s 3d, Cd, 4p are filled in the elements of period 4. In Rb (Z = 37) and Sr (Z = 38), the 5s orbital is filled. After that in elements from y(Z = 39) to Cd (Z = 48) filling of 4d orbitals takes place. You can see from table 3.1 that once again there are minor irregularities in the distribution of electron between 4d and 5s orbitals. For example Mo (Z = 42) and Ag (Z = 47) have respectively [Kr] 4d55S1 and [Kr] 4d105S1 configurations alike to those of Cr and Cu respectively. Anomalous electronic arrangement of Nb, Ru, Rh and Pd can't be described in simple terms. We have to thus, keep in mind them as exceptions. Now in the next six elements, which is I, Sn, Sb, Te, I and Xe filling of Sp orbitals take place and thus Xe (Z = 54 attains [Kr] 4d105S25p6   configuration.

Period 6:

This period enclose 32 elements from caesium (Z = 55) to radon (Z = 86) in which the 6s, 4f, 5d and 6p orbitals are filled. The first two elements of this period have configurations analogous to those of corresponding member of the lower periods, thus caesium (Z= 55) and barium (Z =56) have [Xe] 6S1 and [Xe] 6S2 configuration correspondingly. According to aufbau principle in the next element La (Z = 57), the additional electron should enter 4f orbital. Instead it goes to the 5d orbitals and La has the configuration [Xe] 5d16S2. But why? The extra electron in the building up of La atom goes to 5d orbital instead of 4f orbital because in La atom, the 5d and 4f orbitals have almost the same energy and hence, the electron is free to enter any of these two orbitals.

In the next 14 elements from cerium (Z = 58) to lutecium (Z=71), the 4f orbital is successively filled pertaining to [Xe] 4f15d16S2   and [Xe] 4f14   45d16S2   configuration, respectively, but we should remember, it is only ce (Z= 58) Gd (Z = 64) and Lu (Z = 71) that 5d orbitals include one electron while in all the remaining Lanthanides the 5d orbitals remain vacant.

After Lutecium, successive electrons occupy 5d orbitals and the electronic configuration builds up from [Xe] 4f145d26S2 for hafnium to [Xe] 4f14 5d10 6S2 for mercury the homologue of zinc and cadmium.

Once more a minor departure from a steady raise in the number of d electrons takes places. For instance, gold has [Xe] 4f14 5d9 6S2, and as you can see, has to do with the greater stability of half filled/fully filled orbitals. At last the period is completed with consecutive occupation of the 6p orbitals from thallium, [Xe] 4f 14 5d106S2 6p1 to radon, [Xe] 4f14 5d10 6s2 6p6.  

Period 7

This period is still incomplete and includes 23 elements from francium (Z = 87) to unnitennium (Z = 109). In these elements electrons are filled in 7s, 5f and 6d orbitals. Antinium ([Ru 36d1 7S2) and Francium ([Ru] 7S1), radium ([Ru] 7S2) contain electronic configurations analogous to those of caecium, barium and lanthanum in that order. Thorium has the configuration [Ru] 6d2 7S2. Thus in the 13 elements from protactinium (Z = 91) to lawrencium (Z = 103) filling of 5f orbitals takes place successively. Though, out of these only Pa (Z = 91), U (Z = 92), Np (Z = 93), Cm (Z = 96) and Lr (Z = 103) have an electron in 6d orbitals. In the rest of the elements, the 6d orbitals remain vacant, thus the electronic configuration of Lr (Z = 103) is [Ru] 5f14 6d27S2.  The next six known elements of this period are members of 6d transition series that we have the configurations [Ru] 5f146c12 7S2 to [Ru] 5f 14 6d 7 7S.

Having examined the electronic configuration of elements in the periodic table, we can see from table about the elements occupying the similar group of the periodic table contain the similar valence-shell electronic configuration. In order words, the elements having the similar valence shell electronic configuration occurring periodically, this is after intervals of 2, 8, 8, 18, 18 or 32 in their atomic number. Thus periodicity in the properties of elements can simply be comprehended.

Electronic Configuration of Ions:

Whenever the gaseous iron atom having [Ar] 3d6 4S2 ground state electronic configuration looses an electron, the Fe+ ion is created. This ions has its minimum energy in the configuration [Ar] 3d7, even though the iso electronic manganese atom has the configuration [Ar] 3d5 4S2 in the ground State. Likewise, the ground state of the Fe2+ and Fe3+ ions are [Ar] 3d6 and [Ar]3d5   correspondingly rather than [Ar] 3d5 4S1 and [Ar] 3d3 4S2 that are ground states of iso electronic atoms of chromium and vanadium correspondingly. Evidently the differences in nuclear charge between Fe+ and Mn, Fe2+ and Cr and Fe3+ and V are vital in determining the orbital to be occupied by the electrons. However, along the series of ions carrying similar charge, the electronic configuration frequently changes much more regularly than the electronic configuration of the analogous atoms. Therefore for tripositive ions Sc3+ to Zn2+, the ground state electronic configuration transforms frequently from [Ar] 3d1 to [Ar] 3d10. For tripositive ions, there is a similar regular change from [Ar] for Sc3+ to [Ar] 3d9 for Zn3+ for tripositve ions of lanthium elements, there is a ordinary transform from [Xe] 4f1 for Ce3+ to [Xe] 4f14 for Lu3+. Because the chemistry of elements is basically that of their ions, the regularities in configuration of ions are a lot further significant than the irregularities in the electronic configuration of the neutral atoms.

Electronic Configuration and Division of Elements into Blocks:

Elements of the periodic table have been divided into 4 blocks s, p, d and f based on the nature of the atomic orbitals to which the differentiating and the last electron enters.

The s-block elements:  In such elements, the differentiating electron enters the 'ns' orbital. Alkali and alkaline earth metals of groups (IA) or 2 (11A) fit into this block.  As we know that valence shell electronic configuration of such groups is ns l   and ns2 correspondingly .We as well know that each and every period of the periodic table starts with alkali.

The p-Block Elements: In such elements belonging to this block, the p-orbitals are sequentially filled. Therefore the elements of the group 13 (IIIA), 14(IVA), 15(VA), 16(VIA), 17(VIIA) or 18(zero) are members of this block, because in the atoms of such elements, the differentiating electron enters the np orbitals. The ns orbital in the atoms of such elements are already wholly filled thus they have the valence shell electronic configuration ns2np 1-6.

It will be noted that the elements of s- and p- blocks are as well termed as normal representative and main group elements.

The d-Block Elements:  such elements, in which the differentiating electron enters the (n-1) d   orbitals are termed as d-block elements. These elements are situated in the middle of the periodic table between the s-and p- block elements. The electronic configuration of the atoms of the elements of this block can be symbolized by (n-1) d1-10 ns012.

Such elements which are also known transition elements are divided into 4 series analogous to the filling of 3d - 4d -5d -or 6d -orbitals whereas the 3d, 4d, and 5d series consist of 10 elements each and every, the 6d series is incomplete and has only seven elements by: Ac (Z = 89) and from Unq 9Z = 104) to Une (Z = 109). The element from Sc (Z = 21) to zn (Z = 30), Y (Z = 39) to Cd (Z =48), La (Z = 57) and from Hf (Z = 72) to Hg (Z = 80) are the members of 3d, 4d, and 5d series correspondingly.

It will be noted that d-Block elements are as well called as transition elements.

The  f-Block Elements: Such elements in which the extra electron enters (n-2)f orbitals are termed the f-block elements. The atoms of these elements contain the general configuration (n-2) f 1-14 (n-1) d 0.1 ns2. Such elements fit into two series depending on the filling of 4f and 5f orbitals. Elements from Ce (Z = 58) to Lu (Z = 71) are the members of the 4f series, whereas those from Th (Z =90) to Lr (Z = 103) belong to the 5f series. Elements of 4f series which follow lanthanium in the periodic table are known as LANTHANIDES whereas those of 5f series following actinium are called ACTINIDES. All such elements are collectively said to as INNER- TRANSITION elements since of filling of electrons in an inner (n-2) f sub shell.

It will be noted that f-block elements are as well called as inner transition elements.

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