covalent/molecular bonding ch. 16. the nature of covalent bonding 16-1 skip pgs 444 - 451
TRANSCRIPT
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Covalent/Molecular Bonding
Ch. 16
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The Nature of CovalentBonding
16-1Skip pgs 444 - 451
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Covalent Bonds• Covalent (molecular) bond = the attraction
of two atoms for a shared pair of electrons– Neither atom will have an ionic charge– Usually between 2 nonmetals (some involve
metalloids)!• Covalent compound = a compound whose
atoms are held together by covalent bonds• Molecule = an uncharged group of two or
more atoms held together by covalent bonds
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Single Covalent Bonds
• Single Covalent Bond = 2 atoms share one pair of electrons.– H2, F2, H2O
• Structural Formula = chemical formula that shows the arrangement of atoms.– H + H H H H – H Element Compound Structural dots dots formula
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– F + F F F F - F
– H + O O H O - H H H H
F2:
H2O:
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HCH4 C + 4H H C H
HH * H C HH
NH3 N + 3H H N HHH N HH
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Double and TripleCovalent Bonds
• Double Covalent bonds = bonds that involve 2 shared pairs of electrons.– O2 O O
• Triple Covalent bonds = bonds that involve 3 shared pairs of electrons.– N2 N N
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O2 O + O O O O O
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N2 N + N N N N N
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CO2 O + C + O O C O
O = C = OAdd to notes,above N2
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Bonding Theories
16-2Skip pgs. 452–454 + 457-459
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VSEPR Theory
• VSEPR Theory states that because electron pairs repel, molecular shape adjusts so that valence-electron pairs are as far apart as possible.– Ex: H2O bond is NOT linear!
2H + O O H H
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VSEPR GeometricsA = Central Atom, X = Attached Species,
E = Extra Pair of e-’s on A
Total # of Attached Species
Species Type
Molecular Geometry
Example
2 AX2 Linear CO2
4 AX4
AX3E
AX2E2
TetrahedralPyramidal
Bent
CH4
NH3
H2O
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A
AA
A A A
107° 105°
Triatomic120°
Pyramidal Bent
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Linear Example
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Tetrahedral Example
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Pyramidal Example
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Bent Example
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Polar Bonds +Molecules
16-3 Part I
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Bond Polarity• Bonding pairs of electrons are pulled, as in a
tug-of-war, between nuclei of atoms sharing electrons.
• If bonding pairs are shared equally it is a nonpolar covalent bond.– Atoms will have equal electronegativities (pg. 405)– Ex: N2, O2, H2, Cl2, CO2
• If bonding pairs are shared unequally it is a polar covalent bond.– Atoms have unequal electronegativity.– H2O, HCl, CO
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Polar Molecules• Polar molecule = one end
of molecule is slightly negative and other end is slightly positive.Ex: HClElectronegativity:
H = 2.1, Cl = 3.0Difference = 0.9
Ex: H2O
H = 2.1, O = 3.5Difference = 1.4
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Electronegativity Differences + Bond Types
Electronegativity Difference
Type of Bond Example
0.0 – 0.3 Nonpolar Covalent H – H (0.0)
0.4 – 1.0 Moderate PolarCovalent
∂+ ∂-H – Cl (0.9)
1.1 – 2.0 Very Polar Covalent
∂+ ∂- H – F (1.9)
> 2.0 Ionic Na+Cl- (2.1)
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• The polarity of a molecule depends on the shape + orientation of the bonds.– Ex: CO2 polarity cancels
out since it is linear = nonpolar molecule
– Ex: H2O poles add up due to its bent shape = polar molecule
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• Polar: F
H C H H
H N H H
OH H
• Nonpolar:H
H C HH
O C O
H HC CH H
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Comparing Ionic + Molecular Properties
Characteristic Ionic Cmpd Covalent/Molecular Cmpd
Representative Unit Formula Unit Molecule
Bond Formation Transfer e-’s Share pairs e-’s
Type of Element Metal + nonmetal 2 Nonmetal (possible metalloid)
Physical State Solid S, L, or G
Melting Point High (>300°C) Low (< 300°C)
Solubility in Water High High to Low
Electrical Conductivity as aqueous soln
Good Poor to none
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Attractions Between Molecules
16-3 Part II
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Attractions between Molecules• van der Waals forces = weakest
attraction (ionic + covalent are stronger).–Three types:
1. Dispersion forces2. Dipole interactions3. Hydrogen bonds.
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1) Dispersion forces = weakest of all molecular interactions; caused by motion of electrons.• Increases as the number of electrons
increases.• Halogen diatomic molecules (F2, Cl2, Br2, I2)• Fluorine + Chlorine have weak dispersion
forces (less electrons); thus are gases at STP.• Bromine (more electrons) is a liquid at STP,
and Iodine (most electrons) is a solid at STP.
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2) Dipole Interactions = occurs when polar molecules are attracted to one another.– The slightly
negative region of a polar molecule is attracted to the slightly positive region of another polar molecule
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– When placed in an electric field, dipole molecules become oriented with respect to (-) and (+) charge
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3) Hydrogen bonds = attractive forces in which a hydrogen covalently bonded to a very electronegative atoms is also weakly bonded to an unshared electron pair of another electronegative atom.-Strongest of intermolecular (van der Waals) forces
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Network Solids• Most molecules are easy to break; however, a few molecular solids are
very stable.• Network Solids = solids in which all atoms
are covalently bonded to each other.– Solid does not “melt” until 1000°C or higher,
in which it vaporizes without melting at all.– Ex: Diamond; made of carbon, each carbon
bonded to 4 other carbons. Quartz (SiO2)
C – C – C – C
C – C – C – C
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Diamond +Silicon carbide (SiC)