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Chemical Bonding
There are basically two types of chemical bonds:
- Covalent bonds—electrons are shared by more than one nucleus
- Ionic bonds—electrostatic attraction between ions creates chemical bond Hydrogen bonding and van der Waals’ bonding are subsets of electrostatic bonding
The Octet Rule
Atoms want to have a filled valence shell—for the main group elements, this means having filled s and p orbitals. Hence the name “octet rule” because when the valence shell is filled, they have a total of eight valence electrons. We use “dot structures” to represent atoms and their electrons.
Dots around the elemental symbol represent the valence electrons. Examples
Hydrogen (1s 1 ) H· Carbon (1s 2 2s^2 2p 2 ) C
Chlorine ([Ne] 3s 2 3p 5 ) Cl
Lewis Dot Structures
· ·
·· · ··
:··
Lewis Dot Structures
None of the atoms in the previous example contained full valence shells. When creating bonds, atoms may share electrons in order to complete their valence shells.
Lewis Dot Structures
Examples F 2 molecule: F has seven valence e-^ ’s, but wants eight e-^ ’s to fill its valence shell.
:F· + ·F: :F:F: :F – F:
A line is often used to represent shared electrons.
Lewis Dot Structures
Examples CH 4 molecule: C has four valence e-^ ’s; H has one valence e-^.
H
·C· + 4 ·H H· ·C· ·H H : C : H
H
H
H·
H – C – H
H
H
ı ı
H·
H
H· ·C· ·C· ·H· H : C : C : H
H
H
Lewis Dot Structures
Examples C 2 H 4 molecule: Is this complete—do all atoms have filled valence shells?
H : C C : H
H
H
Lewis Dot Structures
Examples C 2 H 4 molecule:
H : C ··C : H
H
H
H – C – –C – H
H
H
ı
ı
Sharing of four electrons between two nuclei results in a “double” bond.
Electronegativity and Bonding
When nuclei share electrons to form a covalent chemical bond, the electrons are not necessarily shared equally—a shared electron may spend more time closer to one of the nuclei. The electronegativity of the nuclei determines how the electron is shared. Electronegativity is a measure of how strongly a “bound” electron participating in a chemical bond is attracted to a nucleus.
Electronegativity and Bonding
Electronegativity is related to electron affinity and ionization energy. Electronegativity (denoted by the greek symbol ) is highest for elements in the upper right hand side of the Periodic Table and increases from left to right and from bottom to top.
Electronegativity and Bonding
Polar Covalent Bonds
When two elements with different electronegativities bond, the resulting covalent bond will be polar , i.e ., the shared electrons will spend more time closer to the nucleus with the higher , so that end of the bond will be slightly more negative, and the other end will be slightly more positive.
H - F
=2.1 =4.
+ -
Polar Covalent Bonds
Examples Carbon monoxide: CO Is CO a polar molecule?
(O) = 3.
(C) = 2.
+ - :C O:
Polar Covalent Bonds
Examples Carbon dioxide: CO 2 is a linear triatomic molecule. Is CO 2 a polar molecule?
(O) = 3.
+ -^ (C) = 2.
O = C = O
.. .. ..
- + ..
The bonding between atoms can have a significant effect on the bond distance between atoms. Multiple bonds between two atoms have shorter bond lengths compared to single bonds involving the same elements:
Bond Lengths
Examples average bond lengths single bond C-C 154 pm C-O 143 pm double bond C=C 133 pm C=O 120 pm triple bond CC 120 pm CO 113 pm
Bond Lengths
Examples (see Table 8-2 for a more complete list) Average bond energy (kJ mol -1^ ) single bond double bond triple bond C-H 416 C-C 356 598 813 C-O 336 750 1073 C-N 285 616 866 N-N 160 418 946 N-O 201 605 C-S 272 575
Bond Energies
Bond Energies
Example: Calculate the reaction enthalpy for the combustion of methane Step 1—Write balanced chemical equation: CH 4 + 2 O 2 CO 2 + 2 H 2 O
Step 2—Determine energy needed to break bonds: 4 C-H bonds: 4 x 416 kJ/mol = 1664 kJ/mol 2 O=O bonds: 2 x 498 kJ/mol = 996 kJ/mol total energy to break bonds = 2660 kJ/mol
Bond Energies
Example: Calculate the reaction enthalpy for the combustion of methane Step 3—Determine energy released in forming new bonds: 2 C=O bonds: 2 x -750 kJ/mol = -1500 kJ/mol 4 O-H bonds: 4 x -467 kJ/mol = -1868 kJ/mol total energy to form bonds = -3368 kJ/mol Step 4—Determine enthalpy of reaction: Hrxn = E (^) break bonds + E (^) form bonds = 2660 kJ/mol – 3368 kJ/mol = -708 kJ/mol Hrxn = -802 kJ/mol (literature value)
Example: combustion of acetylene 2 C 2 H 2 + 5 O 2 4 CO 2 + 2 H 2 O
- Which bonds are broken? 4 x C-H 4 (415 kJ mol-1^ ) = 1660 kJ mol- 2 x CC 2 (835 kJ mol -1^ ) = 1670 kJ mol- 5 x O=O 5 (495 kJ mol-1^ ) = 2475 kJ mol - total energy needed = 5805 kJ mol -
Reaction Energies
Example: 6 H 2 O + 2 N 2 4 NH 3 + 2 O (^2)
- Which bonds are broken? 12 x O-H 12 (460 kJ mol-1^ ) = 5520 kJ mol^ - 2 x NN 2 (945 kJ mol-1^ ) = 1890 kJ mol - total energy needed = 7410 kJ mol -
Reaction Energies
Example: 6 H 2 O + 2 N 2 4 NH 3 + 3 O (^2)
- Which bonds are formed? 12 x N-H 12 (390 kJ mol-1^ ) = 4680 kJ mol- 3 x O=O 3 (495 kJ mol-1^ ) = 1485 kJ mol - total energy released = 6165 kJ mol - E (^) rxn = BEreact - BEprod = 7410 kJ mol -1^ – 6165 kJ mol - = 1245 kJ mol -1^ (1267 kJ mol -1^ actual)
Reaction Energies
Reaction Energies
Energy reacants
H 2 O + N (^2)
NH 3 + O (^2) products
1245 kJ mol -
Endothermic reaction
Formal Charge of Atoms
The formal charge of an atom in a molecule is the charge the atom would have if all electrons were shared equally. Rules:
- All lone pair electrons are assigned to the atom to which they are associated.
- Half of the bonding electrons are assigned to each atom comprising that bond. The sum of these electrons is subtracted from the number of valence electrons to determine the formal charge
Formal Charge of Atoms
Example: Nitrous oxide, N 2 O .. O=N=N.. ..
:O..NN:
0 +1 - structure #
-1 +1 0 structure #
The formal charge can help predict which structure is preferred when multiple Lewis structure can be made:
- Smaller formal charges are preferred
- Negative formal charge should reside on more electronegative atoms
- Like charge should not be on adjacent atoms
Molecular Orbitals
When atomic orbitals overlap to create a covalent bond, the result is the formation of molecular orbitals. Molecular orbitals define the region of space most likely to contain bonding electrons—MO’s are drawn as 90% electron density contours just as AO’s are drawn 90% electron density contours in atoms.
Molecular Orbitals
Because electrons can be described as waves, when the AO’s overlap, the waves may either constructively interfere or destructively interfere. Constructive interference between the AO’s results in a bonding MO—destructive interference between AO’s results in an anti- bonding MO with a node along the internuclear axis. Anti-bonding orbitals are higher in energy than bonding orbitals.
Molecular Orbitals
Overlap of s-type atomic orbitals to form either bonding or anti-bonding molecular orbitals. Anti-bonding orbitals are designated with an asterisk (*).
