Chemical Periodicity / Formal Charge
- Relationship between Electronic Structure and the Periodic Table
(Only the first four rows are shown.)
- First ionization energy - The energy required to remove the outermost electron from a neutral atom in the gas phase.
Electron affinity - The energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion.
Electronegativity - The tendency of an atom to draw the electrons in a bond towards it.
The first ionization energies, electron affinities, and electronegativities of elements generally increase as one moves closer to the top right-hand corner of the periodic table. They increase as one moves across a row from left to right because the force of attraction between the nucleus and an electron increases as the number of protons in the nucleus increases. They decrease as one moves down a group in the periodic table because the electrons in inner orbitals tend to shield the outer electrons from the attractive force of the nucleus.
Group I elements have large second ionization energies compared with their first ionization energies, while group II elements have second ionization energies not much higher than their first ionization energies. This can be explained by their respective electronic configurations.
Na has a small first ionization energy but a large second ionization energy. Cl has a large electron affinity and electronegativity. As a result, when Na and Cl2 are placed in the same container, they react: Cl readily gains an electron to become Cl-, while Na readily loses an electron to become Na+. The difference in electronegativities between Na and Cl means that the bond between them is ionic instead of covalent. The large second ionization energy of Na is the reason why NaCl2 is not formed.
Mg has small first and second ionization energies, but a large third ionization energy. Thus, magnesium chloride has the formula MgCl2. Similarly, the difference in electronegativities between Mg and Cl means that the bond between them is ionic instead of covalent.
Formal charge is the charge on an atom in its Lewis structure, or the overall charge on an ion or molecule.
Steps in determining formal charge:
1.Divide the electrons in each covalent bond between the atoms in the bond.
2.Determine whether the atom of interest has more electrons than protons (negative formal charge), less (positive formal charge), or the same number (zero formal charge).