Explain the reason for the variation of the first ionization energies of the third period elements.
Ionization energy is the energy required to remove one electron from an atom in its gaseous state.
In general, ionization energy increases across a period on the periodic table since atomic radius decreases across the period. The electrons are closer to the nucleus, and it becomes harder to remove an electron. Ionization energy decreases down a group since atomic radius increases down the group and there are additional shells that shield the attractional force of the nucleus on the valence electrons.
The trend across a period has two main exceptions, between group 2 and 3 and between group 5 and 6.
The ionization energy for group 3 is lower than the ionization energy for the element in group 2 and in the same period. This is because, for group 3, the outermost electron is in the p-orbital, while the outermost electron is group 2 is
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From the infrared spectrum given below, identify the possible functional groups corresponding to the peaks indicated in the arrows.
Explain the type of hybridization in phosphorous pentafluoride
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CH2 = C = CH2
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- 10 B + n → .... + α
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- 241Am + ...→ 243Bk + 2n
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