Periodic Table Trends and Polar Covalent Bonding

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Periodic Table Trends

Properties of elements across the periodic table change in a regular, predictable pattern.  These properties are based on periodic changes in atom size.  Regular, periodic property changes also result in similar chemical behavior for elements in the same family (column).  Going left to right across a row of the periodic table, we see that atoms become smaller.  We also see that atoms become smaller as we go from the bottom of the periodic table to the top within the same family.  The combined result is that the smallest element atoms on the periodic table are in the upper right-hand corner.  Those elements also have the highest ionization energy (IE) and electronegativity (EN) values.  IE is the energy required to remove a single valence electron from a neutral atom.  In other words, IE is the energy needed to ionize an element.  IE is normally expressed in Joules per mole of element atoms, J/mol.  EN is the relative attraction that an atom nucleus has for electrons in a covalent bond.  The EN scale was developed by the great scientist, Linus Pauling, and ranges from 0 to 4, with "4" being the most electronegative.  The diagram below shows why the atomic sizes vary the way they do across the periodic table.

As the figure shows, new electron shells are added as we go down the periodic table and electrons are added to the same shell as we move through a row to the right. The trend in atom size is directly related to trends in IE and EN values, as explained above in the diagram.

Polar Covalent Bonding

Any covalent bond between atoms of different elements will have an unequal sharing of electrons.  This results in a bond with polarity, meaning that one atom has a "partial positive" nature and the atom on the other end of the bond has a "partial negative" nature.  The polarity becomes substantial when elements two or more table "squares" away from each other are covalently bonded.

If a molecule is asymmetrical, like water (bent geometry), then the entire molecule is dipole and we say it has a "dipole moment".  The resulting molecular dipole has an uneven distribution across the structure with one end taking on a partial positive charge and the other end taking on a partial negative charge.  The following figure provides a few examples of molecular dipoles.

As the figure indicates, polar covalent bonds exist in both molecules.  Both molecules are asymmetrical with the more electronegative element to the right.  This results results in net dipoles for both molecules.  The partial negative charges have coefficients matching the number of partial charges (the H atoms) on the left sides.

That's all for this post.  The next post will discuss intermolecular, which helps explain the myriad properties observed with molecular compounds!

Have a good one!

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