When elements react, they do so via sharing losing and gaining of electrons. This procedure of gaining, losing or sharing of electrons is generally accompanied through energy transforms. We recognize that the electrons to be divided, gained or lost are bound to the nucleus of its atom through an electrostatic force of attraction. In order to take away an electron from an atom, its force of attraction has to be conquered. This can be complete via supplying energy. The energy necessary to take away the least strongly bond electron from an isolated gaseous atom in its ground condition is recognized as the ionization energy. This procedure can be represented via the given equation:
M (g) →M+ (g) + e
Since more than 1 electron might be taking away from an atom, the energy involved for the above procedure is termed the 1st ionization energy. The 2nd ionization energy is the energy needed to take away an electron from a univalent cation that procedure is represented via this reaction.
M+ (g) → M2+ (g) + e
The 2nd ionization energy is much larger than the first ionization energy. This is since in this case an electron is being taken away from a positively charged cation. Likewise we can describe third, fourth and higher ionization energies. The S I unit of ionization energy that we will use all through this course is KILOJOULE per mole. We will be studying the factors which affect ionization energy, periodicity in ionization across the periods, trends in ionization energy, down the groups and trends in successive ionization energy.
Factors affecting ionization energies:
The ionization energy I of an outer valence electron is related to the effective nuclear charge felt via the electron or its average distance from the nucleus as stated in the equation below:
Here Z* is the effective nuclear charge, e is the charge on electron and (I/r) av is the average value of the electron from the nucleus; therefore the higher effective charge felt through the electron, the higher will be the ionization energy. As well the further the electron is from the nucleus, the lower will be the ionization energy or vice-versa.
The ionization energy as well depends on the virtual stabilities of the sub shell from that the electron is taking away. As we have seen before, entirely filled or half filled sub shells are relatively more stable, so removal of an electron from them wants more energy. The valence shell electronic configurations of noble gases are remarkably stable or then their ionization energies are the maximum in the irrespective period.
Periodicity in Ionization Energy across Periods:
The deviation in ionization energy in a meticulous group and period is best shown via a graph showing ionization energies beside atomic number. Fig illustrates the plot of 1st ionization energies of the elements of the first 6 periods against their atomic numbers. since is obvious from the figure, the first ionization energy usually raises from alkali metals to noble gases across several row of the periodic table. Other than the raise is not perfectly regular.
We have seen from a previous section that across any row of the periodic table the effective nuclear charge gradually increases or the atomic radii reduces. These two effects reinforce each and every another to amplify the ionization energies across a period. Therefore the ionization energies of the alkali metals are the lowest or those of the noble gases are the highest in their relevant periods. Though as pointed out previous the raise is not smooth or several anomalies are viewed. For instance in the elements of period 2, inspite of raise in Z* or reduce in r, the 1st ionization energies of B and 0 are lower than those of Be and N correspondingly. Though, these anomalies in the trend in ionization energy can be explained via electronic structures of these elements.
In the case of beryllium, the electron is removed from the filled 2s sub shell whereas in boron, the electron is removed from the singly occupied 2p sub shell. The 2p sub shell is higher in energy than the 2s, so the 2p electron of boron is more easily removed than the 2s electron of beryllium. When we come to nitrogen, we will find out that we have a half filled 2p sub shell (electronic configuration Is2, 2s2, 3p3) whereas in oxygen the 2p sub shell is occupied through 4 electrons. The 4th electron in this 2p sub shell is in an orbital previously occupied via the other electron, consequently it experiences substantial repulsion. Accordingly, this electron is more easily take away than one of the electrons from a singly occupied orbital in nitrogen atom. Thus the ionization energy of oxygen becomes less than that of nitrogen.
Analogous anomalies are observed in elements of period 3 wherever the first ionization energies of magnesium and phosphorous are higher than those of aluminum or sulphur correspondingly. We have previous seen that across the conversion series, the raise in effective nuclear charge or resultant reduce in atomic radius is small. Consequently raise in the first ionization energies is as well small. Though following the transition elements the 1st ionization energy drops suddenly in thallium, indium and gallium. This yet again is due to the taking away of an electron from a singly occupy np orbitals which are of comparatively higher energy than the ns orbital of Zn, Cd and Hg.
Trends in Ionization Energy down the Groups:
We have studied from a previous sub division which on moving down a group of s- or p- block elements effective nuclear charge remains almost steady. But there is a general raise in the atomic radius due to increase in the value of the principal quantum number n 1 , therefore the dominant factor in determining the ionization energies of the elements on moving down the collection is their atomic radius rather than the effective nuclear charge. Thus, as expected, the first ionization energies reduce down the groups in the case of the main group elements in the periodic table. But in the case of conversion elements opposite trends are observed. Hence, the first ionization energies of the analogous elements of 3d or 4d series are approximately similar but these are smaller than the first ionization energies of the elements of 5d series. Indeed the higher values of ionization energies of the 5d conversion elements are consistent with the relatively smaller size of their atoms.
Trends in Successive Ionization Energies:
We have already defined consecutive ionization energies, which are second, third etc. Values of 8 successive ionization energies of the first 20 elements are listed. It is evident from the values in the table that the successive ionization energies of an element inevitably become larger since the removal of consecutive electron leaves a higher charge on the nucleus to hold the remaining electrons. It is as well clear from the table that the difference between successive ionization energies of the same element is not steady. Big jumps take place when ever an electron from a sub shell of lower principal quantum number is removed for the first time. For instance, for alkali metals the second ionization energies are much higher than the first. For alkaline earth, metal the third ionization energies are much larger than the second and for the halogen, the 8th ionization energies are much greater than the 7th. These though can't be explained on the basis of amplify in nuclear charge alone. The stabilities of closed shell configuration similar to those of the noble gases are more significant in these cases. So far in our learning, we have seen that ionization energy usually increases across a period and reduces down the group in the periodic table. Therefore the tendency to form cation that is metallic character decreases across a period and amplifies down the group. For instance, in period 3 metallic character reduces from Na to Cl while in the elements of group 14, C is a non metal, Si or Ge are semi metals and metalloid whereas Sn and Pb are metals.
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