the nature of the chemical bond - weebly
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The Nature of the Chemical Bond
Figure 1: The formation of a single covalent bond in a hydrogen molecule by the overlap of two 1s orbitals or individual hydrogen atoms. This represents a new, lowerenergy state of the two atoms.
ANIMATION OF ELECTRON DENSITY IN THE BOND REGION
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Figure 2: A hydrogen atom has only one occupied orbital, the 1s orbital. For simplicity, only the one halffilled p orbital of the fluorine atom is shown. The final combined, bonding orbital contains a pair of electrons and the fluorine atom now has a complete octet.
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Figure 3: Two overlaps are possible to produce two shared pairs of electrons forming two covalent bonds. As before, the oxygen atom completes its octet of electrons.
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Hybrid Orbitals
Consider the 2nd energy level for a groundstate carbon atom.
It would appear that carbon has one lone pair and only 2 bonding electrons. Why does it form four bonds? Linus Pauling was the first to suggest “electron promotion”. One s electron gets promoted to the empty p orbital. He justified this by suggesting that the energy gained by the molecule when it bonds would be greater than the energy required for promotion to a slightly higher energy level. However, experimental evidence suggests that the electron orbitals are all equivalent in shape and energy there isn’t one that’s more “s” than the others. As a result, the four bonds for carbon in molecules such as methane are explained by hybridization to four identical sp3 atomic orbitals.
These hybridized orbitals only exist when bonding occurs and are NOT found in an isolated atom.
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Example 1 What are the bonding orbitals and the structure of the BF3 molecule?
Solution
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Example 2 Provide the groundstate and the promoted state configurations for beryllium and then describe the bonding and structure of a BeH2 molecule.
Solution
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Double and Triple Covalent Bonds
sigma (σ)bond:a bond created by the end to end overlap of atomic orbitals (like those below)
Figure 4: Sigma bonds form with the overlap of (a) s orbitals (b) p orbitals and (c) hybrid orbitals.
ANIMATION OF SIGMA BOND FORMATION
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pi (π) bond: a bond created by the sidebyside (or parallel) overlap of atomic orbitals (usually p orbitals)
Figure 5: p orbitals form with the sidebyside overlap of orbitals
ANIMATION OF PI BOND FORMATION
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Double Bonds
The carbon atom is the most common central atom in molecules with double and triple covalent bonds. It is thought that there is a partial hybridization of the available orbitals leaving one or two p orbitals with single unpaired electrons.
Figure 6: For this carbon atom, the sp2 hybrids are planar at 120° to each other and the p orbital is at right angles to the plane of the hybrid orbitals.
For example, after promoting an electron in carbon’s 2s orbital to a 2p orbital, we form only three sp2 hybridized orbitals but three of these are hybrids and one is a “normal” p orbital.
ANIMATION OF ETHYLENE BONDS
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Figure 7:
(a) The sigma bonds for a C2H4 molecule use the sp2 hybrid orbitals.
(b) The two halffilled p orbitals of the adjacent carbon atoms overlap sideways.
(c) The complete bonding orbitals for a C2H4 molecule.
Therefore, a double bond consists of a σ bond and a π bond.
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Triple Bonds
For C2H2:
Figure 8: Instead of mixing all four orbitals, valence bond theory suggests that only two are mixed to form sp hybrid orbitals and two unhybridized p orbitals for a carbon atom.
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Figure 9: (a) The sigma bonds for a C2H2 molecule use the sp hybrid orbitals.
(b) The two pairs of halffilled p orbitals of the adjacent carbon atoms overlap sideways.
(c) The complete bonding orbitals for a C2H2 molecule.
Therefore, triple bonds involve a σ bond and 2 π bonds.
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Sigma Bonds in Ethylene.mov
Attachments Page 1
Figure 1: The formation of a single covalent bond in a hydrogen molecule by the overlap of two 1s orbitals or individual hydrogen atoms. This represents a new, lowerenergy state of the two atoms.
ANIMATION OF ELECTRON DENSITY IN THE BOND REGION
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Figure 2: A hydrogen atom has only one occupied orbital, the 1s orbital. For simplicity, only the one halffilled p orbital of the fluorine atom is shown. The final combined, bonding orbital contains a pair of electrons and the fluorine atom now has a complete octet.
3
Figure 3: Two overlaps are possible to produce two shared pairs of electrons forming two covalent bonds. As before, the oxygen atom completes its octet of electrons.
4
Hybrid Orbitals
Consider the 2nd energy level for a groundstate carbon atom.
It would appear that carbon has one lone pair and only 2 bonding electrons. Why does it form four bonds? Linus Pauling was the first to suggest “electron promotion”. One s electron gets promoted to the empty p orbital. He justified this by suggesting that the energy gained by the molecule when it bonds would be greater than the energy required for promotion to a slightly higher energy level. However, experimental evidence suggests that the electron orbitals are all equivalent in shape and energy there isn’t one that’s more “s” than the others. As a result, the four bonds for carbon in molecules such as methane are explained by hybridization to four identical sp3 atomic orbitals.
These hybridized orbitals only exist when bonding occurs and are NOT found in an isolated atom.
5
6
Example 1 What are the bonding orbitals and the structure of the BF3 molecule?
Solution
7
Example 2 Provide the groundstate and the promoted state configurations for beryllium and then describe the bonding and structure of a BeH2 molecule.
Solution
8
Double and Triple Covalent Bonds
sigma (σ)bond:a bond created by the end to end overlap of atomic orbitals (like those below)
Figure 4: Sigma bonds form with the overlap of (a) s orbitals (b) p orbitals and (c) hybrid orbitals.
ANIMATION OF SIGMA BOND FORMATION
9
pi (π) bond: a bond created by the sidebyside (or parallel) overlap of atomic orbitals (usually p orbitals)
Figure 5: p orbitals form with the sidebyside overlap of orbitals
ANIMATION OF PI BOND FORMATION
10
Double Bonds
The carbon atom is the most common central atom in molecules with double and triple covalent bonds. It is thought that there is a partial hybridization of the available orbitals leaving one or two p orbitals with single unpaired electrons.
Figure 6: For this carbon atom, the sp2 hybrids are planar at 120° to each other and the p orbital is at right angles to the plane of the hybrid orbitals.
For example, after promoting an electron in carbon’s 2s orbital to a 2p orbital, we form only three sp2 hybridized orbitals but three of these are hybrids and one is a “normal” p orbital.
ANIMATION OF ETHYLENE BONDS
11
Figure 7:
(a) The sigma bonds for a C2H4 molecule use the sp2 hybrid orbitals.
(b) The two halffilled p orbitals of the adjacent carbon atoms overlap sideways.
(c) The complete bonding orbitals for a C2H4 molecule.
Therefore, a double bond consists of a σ bond and a π bond.
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Triple Bonds
For C2H2:
Figure 8: Instead of mixing all four orbitals, valence bond theory suggests that only two are mixed to form sp hybrid orbitals and two unhybridized p orbitals for a carbon atom.
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Figure 9: (a) The sigma bonds for a C2H2 molecule use the sp hybrid orbitals.
(b) The two pairs of halffilled p orbitals of the adjacent carbon atoms overlap sideways.
(c) The complete bonding orbitals for a C2H2 molecule.
Therefore, triple bonds involve a σ bond and 2 π bonds.
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Sigma Bonds in Ethylene.mov
Attachments Page 1