chapter 16 covalent bonding
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Chapter 16 Covalent Bonding. The Nature of Covalent Bonding Bonding Theories Polar Bonds and Molecules. Chapter 16.1 The Nature of Covalent Bonding. Single Covalent Bonds Double and Triple Covalent Bonds Coordinate Covalent Bonds Bond Dissociation Energies Resonance - PowerPoint PPT PresentationTRANSCRIPT
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Chapter 16Covalent Bonding
The Nature of Covalent Bonding
Bonding Theories
Polar Bonds and Molecules
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Chapter 16.1 The Nature of Covalent Bonding
Single Covalent Bonds
Double and Triple Covalent Bonds
Coordinate Covalent Bonds
Bond Dissociation Energies
Resonance
Exceptions to the Octet Rule
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Single Covalent Bonds
A bond in which two atoms share a pair of electrons (usually non-metals)
The pair of electrons is represented by a dash
Structural formulas – chemical formulas that show the arrangement of atoms in molecules or poly atomic ions
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Single Covalent Bonds
Hydrogen
H2
H H
H-H
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Single Covalent Bonds
Water
Ammonia
Methane
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Double and Triple Covalent Bonds
Double Covalent Bonds – involve two shared pairs of electrons
Triple Covalent Bonds – involve three shared pairs of electrons
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Double and Triple Covalent Bonds
Nitrogen
N2
N N
N N
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Coordinate Covalent Bond
One atom contributes both bonding electrons
Carbon Monoxide
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Bond Dissociation Energies
The total energy required to break the bond between two covalently bonded atoms
H-H + 435kJ = H + H
C-C + 347kJ = C + C
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Bond Dissociation Energies
Bond Bond Energy Bond Length
H-H 435 74
C-H 393 109
C-O 356 143
C=C 657 133
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Resonance
Structures that have two or more different electron dot structures that have the same number of electron pairs for the same molecule or ionOzone
NO3-
http://www.nku.edu/~russellk/tutorial/reson/NO3.gifhttp://www.nku.edu/~russellk/tutorial/reson/CO3.gif
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Exceptions to the Octet Rule
Occurs when the total number of valence electrons is an odd number.
NO2 (Two resonance structures)
BF3
PCl3 and PCl5SF6
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Diamagnetic
All electrons pairedThe spinning of electrons creates magnetic fields. The paired electrons spin in opposite directions, therefore their fields cancel each other outShow a weak attraction to an external magnetic field
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Paramagnetic
Contain one or more unpaired electrons
Creates a magnetic field
Show a strong attraction to an external magnetic field
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Oxygen – An Exception
O=O
Actually a mix ofO=O and O-O
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Ch 16.2 Bonding Theories
Molecular Orbitals
VSEPR Theory
Hybrid Orbitals
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Molecular Orbitals
An orbital resulting from the overlapping of orbitals from two atoms when they bondBonding orbital – molecular orbital with an energy lower than the atomic orbitalAntibonding orbital – molecular orbital with an energy higher than the atomic orbital
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Molecular Orbitals
Sigma Bond – a molecular orbital that is symmetrical along the axis connecting two atomic nuclei
Two s orbitals directly overlap
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Molecular Orbitals
Pi bonds - two parallel 'p' orbitals in close proximity can overlap sideways (laterally)
A pi bond can only form after a sigma bond has already formed
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VSEPR Theory
Electron pairs repel, so molecules adjust shape to keep the pairs as far apart as possible
Unbonded electrons repel bonded electrons more, causing the bonded electrons to be closer together
http://www.mikeblaber.org/oldwine/chm1045/notes/Geometry/VSEPR/Geom02.htm
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Ch 16.3 Polar Bonds and Molecules
Bond Polarity
Polar Molecules
Attractions Between Molecules
Intermolecular Attractions and Molecular Properties
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Bond Polarity
Nonpolar Covalent Bond – equal sharing of electrons, each atom pulls equally (same atoms)
Polar Covalent Bond – unequal sharing of electrons (different atoms)
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Bond Polarity
The more electronegative atom will have a greater pull and acquire a slightly negative charge.
HCl
H2O
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Bond PolarityElectronegativity
DifferenceType of Bond Example
0.0 – 0.4 Nonpolar Covalent H-H
0.4 – 1.0 Moderately Polar Covalent
H-Br
1.0 – 2.0 Very Polar Covalent
H-F
>2.0 Ionic NaCl
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Polar Molecules
Dipole – a molecule with two poles (different charges)
Polarity depends on shape
Atoms that lie in the same axis will cancel each other out
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Attractions Between Molecules
van der Waals forces – weak attractions between molecules – Two TypesDispersion Forces – Weak attraction due to
movement of molecules, increase with more electrons (diatomic halogens)
Dipole Interactions – polar molecules attracted to each other (water molecules)
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