Covalent Bonding
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Atoms bond in such a way as to obtain a stable electron configuration
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Group 8A are particularly stable and uncreative; they possess an octet of
electrons in their valence shells.
Octet rule: the number of electrons associated with bond pairs and lone
pairs for each of the Lewis symbols in a Lewis structure will be eight (an
octet)
Covalent bond: the sharing of electrons between non-metallic elements to
achieve stable octets; results in the formation of discrete molecules
Three forces must be considered in the formation of covalent bonds (the
sharing of electrons between two nuclei):
1. The attraction of electrons between the two nuclei
2. The repulsion of the two nuclei
3. The repulsion of the shared electrons
The effects of these forces can be represented by a potential energy
diagram for the formation of H2.
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When potential energy is at a minimum this represents the formation of a
stable molecule. For H2, this takes place when the distance between the
two nuclei (the internuclear distance) is 74 pm.
Electron-Dot Formulae of Covalent Compounds
1. Draw the Lewis symbol for each of the atoms in the molecule.
2. Unpaired electrons (bonding electrons) are available for sharing; to form
covalent bonds.
3. Paired electrons (lone pairs) do not take part in covalent bonds.
4. The atom with the most bonding electrons is placed in the centre with the
other atoms bonded to it (this is the central atom).
5. In the resulting Lewis diagram all electrons must be paired an each atom
(except H) must be surrounded by an octet of electrons.
Example:
Structural Formulas
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Bonding representations are often simplified by omitting the lone
electron pairs and by substituting a dash for each bonding pair of
electrons.
The resulting representation is a structural formula.
Multiple Bonds
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Once you have placed all the atoms around the central atom in a Lewis
structure, there may be bonding electrons "left over".
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These bonding electrons can pair up to form multiple bonds between two
atoms.
Double Covalent Bond: two pairs of electrons are shared between bonded
atoms.
Triple Covalent Bond: three pairs of electrons are shared between bonded
atoms.
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When drawing Lewis diagrams that include multiple bonds, all electrons
must still be paired, and each atom (except hydrogen) must be
surrounded by an octet of electrons.
Polar Covalent Bonding
Electronegativity: a measure of the electron attracting-power of a bonded
atom; metals have low electronegativities; non-metals have high
electronegativities.
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Different atoms have different abilities to attract electrons (we saw
this when we discussed electron affinity)
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Experimental evidence indicates that covalently bonded atoms usually
exhibit unequal attractions for shared electrons.
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Linus Pauling developed a scale of electronegativies.
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Pauling arbitrarily assigned the most electronegative atom, fluorine, a
value of 4.0.
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In a covalent bond, the shared electrons are more strongly attracted to
the atom with the higher electronegativity value.
Examining the periodic table reveals two electronegativity trends:
1. Electronegativities increase from left to right within a period
2. Electronegativities decrease from top to bottom in a group.
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Cesium and francium have the lowest electronegativities. They also have
low ionization energies. It is therefore relatively easy to transfer an
electron from cesium or francium to another element.
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Fluorine has the highest electronegativity. Fluorine also has a high
ionization energy. This means that it takes a large amount of energy to
remove an electron from fluorine.
Electronegativity Difference (EN): the difference in the
electronegativities of two atoms that are bonded together; used to assess
the degree of bond polarity.
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That is, we can compare the electronegativity of two atoms in a bond to
determine which atom more strongly attracts the electron pair.
EN  2.0  ionic bond
EN  0  non-polar covalent bond
0 EN 2.0  polar covalent bond
Example: hydrogen fluoride
EN = 4.0 - 2.1 = 1.9 polar covalent bond
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The bonding electrons are found closer to fluorine therefore there will
be a distortion of the electron cloud (area of probability of finding an
electron) toward the fluorine.
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Since the electrons are closer to fluorine there will be a partial negative
charge (indicated by -) on the fluorine atom.
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Also, since the electrons are closer to fluorine, hydrogen has an
unneutralized proton, leading to a partial positive charge (+) on hydrogen.
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Hydrogen fluoride has a partial positive and a partial negative end; this is
known as a bond dipole. The presence of a bond dipole is indicated in a
diagram by an arrow pointing toward the more electronegative atom.
Valence-Shell Electron-Pair Repulsion (VSEPR) Theory and Molecular
Polarity
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Discrete molecules have definite three-dimensional shapes
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Electron pairs repel each other, whether they are in chemical bonds
(bond pairs), or unshared (lone pairs).
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Electron pairs assume orientations about an atom to minimize repulsions.
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The particular geometry of a molecule results from this requirement.
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VSEPR theory proposes a set of rules to allow us to predict molecular
geometry based on the arrangement of electron pairs in the valence shell
of the central atom.
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Lone pairs spread out more in space than do bonding pairs, which are
more localized between atomic nuclei.
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The repulsion between lone pairs is greater than the repulsion between a
lone pair and a bonding pair, or between two bonding pairs.
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This is reflected in the geometry of molecules with lone pairs.
Applying VSEPR Theory
1. Draw the Lewis diagram of the molecule.
2. Determine the number of electron groups around the central atom and
identify them as being either bond pairs or lone pairs.
3. Establish the electron group geometry around the central atom.
4. Determine the molecular geometry from the positions around the central
atom occupied by the other atomic nuclei.
Note: For purposes of applying VSEPR theory only, multiple bonds are
treated as though they are single bonds.
Polar Bonds and Molecular Shapes
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In order to determine if a molecule is polar or not, we need to know both
the geometry of the molecule and the bond dipoles of the molecule.
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If each of the bond dipoles in a molecule cancels with each other, then
we have a non-polar covalent molecule.
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If, however, the bond dipoles are oriented in such a way that they do not
cancel, then partial positive (+) and partial negative (-) charges will be
present and we will have a polar molecule.
In general a molecule will be polar if:
1. lone pairs of electrons on the central atom in a molecule are present
2. different elements surrounding the central atom are present
3. the molecule is a binary molecule involving two different elements