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10.5 Covalent Bonding: Lewis Structures
Weak interactions between molecule must be overcome if a compound is to remain intact.
The melting and boiling points of the compounds are different.
The octet rule requires atoms to form bonds.
The statement is correct.
The rule helps us predict when bonds will form.
The lowering of potential energy is one of the reasons that atoms form bonds.
The rule is not useful.
We know from Chapter 8 that representing electrons with dots, as we do in the Lewis model, is a drastic oversimplification.
Such is not the case.
There are two shared electron dots between the H and F atoms.
We know that they are not.
The H end of the molecule is blue, even though the Lewis structure of the molecule portrays the bonding electrons as residing charge.
There are two extremes to the categories of pure covalent and ionic.
The concept of energy required to break a bond is covered in detail in the case of H2 and F2.
If the electrons were shared equally, the average bond energy of H should be 2 and F. The bond energy of HF is not the same as the average shown here.
The square was taken with amol.
The periodic trends in electronegativity can be found in Figure 10.8: # electronegativity generally increases across a period in the periodic table.
Down a column in the periodic table, electronegativity decreases.
The most negative element is ferriine.
We have seen other periodic trends in electronegativity.
The larger the atom, the less ability it has to attract electrons to itself in a chemical bond.
The elements in order of decreasing electronegativity are P, Na, N, and Al.
If there is a large electronegativity difference between the two atoms in a bond, the electron from the metal is almost completely transferred to the nonmetal, and the bond is ionic.
The bond is polar covalent if there is an intermediate electronegativity difference between the two atoms.
The electronegativity difference between the bonding atoms can be used to classify bonds as covalent, polar covalent, and ionic.
The size of the bond's dipole moment is what we measure.
The most ionic bonds don't reach this ideal.
There is no bond that is 100% ionic.
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