Bonding and Molecular Structure: Fundamental Concepts

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AP Notes Chapter 8

• Bonding and Molecular Structure Fundamental
  Concepts
• Valence e- and Bonding
• Covalent
• Ionic
• Resonance Exceptions to Octet Rule
• Bond Energy Length
• Structure, Shape Polarity of Compounds

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What is a Bond?

• A force that holds atoms together.
• Why?
• We will look at it in terms of energy.
• Bond energy the energy required to break a bond.
• Why are compounds formed?
• Because it gives the system the lowest energy.

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Covalent compounds?

• The electrons in each atom are attracted to the
  nucleus of the other.
• The electrons repel each other,
• The nuclei repel each other.
• The reach a distance with the lowest possible
  energy.
• The distance between is the bond length.

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Thus Hydrogen is Diatomic!
Bond Formation
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Covalent Character
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Why Isnt Helium Diatomic?
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• F F F2
• 2p ____ ____ ___ ___ ____ ____ 2p2s
  ____ ____ 2s
• F F

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Ionic Bonding

• An atom with a low ionization energy reacts with
  an atom with high electron affinity.
• The electron moves.
• Opposite charges hold the atoms together.

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• Li Cl1s22s1 Ne
  3s23p52s ___ 3p _____ _____
  ___1s _____ 3s _____
  Ne

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• Li Cl 2s ___ 3P _____ _____
  _____1s _____ 3s _____ Ne

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• LiCl2s ___ 3P _____
  _____ _____1s _____ 3s _____
  Ne

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Electronegativity

• Describes the relative ability of an atom within
  a molecule to attract a shared pair of electrons
  to itself.

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Electronegativity

• Pauling electronegativity values, which are
  unit-less, are the norm.

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ElectronegativityRange from 0.7 to 4.0
Figure 9.9 Kotz Treichel
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Bond A - B

• DEN ENA - ENB

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Bond Character

• Ionic Bond - Principally Ionic
  Character
• Covalent Bond - Principally
  Covalent Character

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Determining Principal Character of Bond
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F - F EN 0
Non-polar
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N - O ?EN 3.0 - 3.5
0.5
O
N
Slightly polar
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Ca - O ? EN 1.0 - 3.5
2.5
Ca
O
Ionic Bond with somecovalent character
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Electronegativity

• The ability of an electron to attract shared
  electrons to itself.
• Pauling method
• Imaginary molecule HX
• Expected H-X energy H-H energy X-X
  energy 2
• D (H-X) actual - (H-X)expected

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Electronegativity

• D is known for almost every element
• Gives us relative electronegativities of all
  elements.
• Tends to increase left to right.
• decreases as you go down a group.
• Noble gases arent discussed.
• Difference in electronegativity between atoms
  tells us how polar.

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Electronegativity difference
Bond Type
Zero
Covalent
Covalent Character decreases Ionic Character
increases
Intermediate
Large
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Dipole Moments

• A molecule with a center of negative charge and a
  center of positive charge is dipolar (two poles),
• or has a dipole moment.
• Center of charge doesnt have to be on an atom.
• Will line up in the presence of an electric field.

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How It is drawn
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Which Molecules Have Them?

• Any two atom molecule with a polar bond.
• With three or more atoms there are two
  considerations.
• There must be a polar bond.
• Geometry cant cancel it out.

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Ionic Radii -- Cations
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Ionic Radii -- Anions
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Molecular Polarity
Vector Sum of Bond Polarities
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MgBr2 Mg - Br EN 1.2 -
2.8 1.6
Mg
Br
Br
Covalent BOND w/much ionic character, BUT
NON-POLAR molecule
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Lewis Structures
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The most important requirement for the formation
of a stable compound is that the atoms achieve
noble gas e- configuration
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Valence Shell ElectronPair Repulsion
Model(VSEPR)

• The structure around a given atom is determined
  principally by minimizing electron-pair repulsions

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VSEPR
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LEWIS STRUCTURES

• draw skeleton of species
• count e- in species
• subtract 2 e- for each bond in skeleton
• distribute remaining e-

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Distinguish Between ELECTRONIC Geometry
MOLECULAR Geometry
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CH4
Bond angle 109.50
Electronic geometry tetrahedral Molecular
geometry tetrahedral
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H3O
Bond angle 1070
Electronic geometry tetrahedral Molecular
geometry trigonal pyramidal
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H2O
Bond angle 104.50
Electronic geometry tetrahedral Molecular
geometry bent
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NH2-
Bond angle 104.50
Electronic geometry tetrahedral Molecular
geometry bent
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Octet Rule holds for connecting atoms, but may
not for the central atom.
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BaI2
Bond angle 1800
Electronic geometry linear Molecular geometry
linear
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BF3
Bond angle 1200
Electronic geometry trigonal planar Molecular
geometry trigonal planar
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PF5
Bond angle 1200 900
Electronic geometry trigonal bipyramidal Molecula
r geometry trigonal bipyramidal
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SF4
Bond angle 1200 900
Electronic geometry trigonal bipyramidal Molecula
r geometry see-saw
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ICl3
Bond angle lt 900
Electronic geometry trigonal bipyramidal Molecula
r geometry T-shape
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I3-
Bond angle 1800
Electronic geometry trigonal bipyramid Molecular
geometry linear
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PCl6-
Bond angle 900
Electronic geometry octahedral Molecular
geometry octahedral
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BrF5
Bond angle 900
Electronic geometry octahedral Molecular
geometry square pyramidal
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ICl4-
Bond angle 900
Electronic geometry octahedral Molecular
geometry square planar
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Actual shape
Non-BondingPairs
ElectronPairs
BondingPairs
Shape
2
2
0
linear
3
3
0
trigonal planar
3
2
1
bent
4
4
0
tetrahedral
4
3
1
trigonal pyramidal
4
2
2
bent
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Actual Shape
Non-BondingPairs
ElectronPairs
BondingPairs
Shape
5
5
0
trigonal bipyrimidal
5
4
1
See-saw
5
3
2
T-shaped
5
2
3
linear
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Actual Shape
Non-BondingPairs
ElectronPairs
BondingPairs
Shape
6
6
0
Octahedral
6
5
1
Square Pyramidal
6
4
2
Square Planar
6
3
3
T-shaped
6
2
1
linear
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What happens when there are not enough electrons
to satisfy the central atom?
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EXAMPLES

• Ethene
• Acetic Acid
• Oxygen
• Nitrogen

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RESONANCE FORMAL CHARGE
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Resonance

• Sometimes there is more than one valid structure
  for an molecule or ion.
• NO3-
• Use double arrows to indicate it is the average
  of the structures.
• It doesnt switch between them.
• NO2-
• Localized electron model is based on pairs of
  electrons, doesnt deal with odd numbers.

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EXAMPLES

• Nitrate ion
• Ozone

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FORMAL CHARGE

• the charge assigned to an atom in a molecule or
  polyatomic ion

FC atom Family - LPE ½(BE) Sum FCs
atoms ion charge Closer sum FCs is to zero
more stable
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Formal Charge

• For molecules and polyatomic ions that exceed the
  octet there are several different structures.
• Use charges on atoms to help decide which.
• Trying to use the oxidation numbers to put
  charges on atoms in molecules doesnt work.

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Formal Charge

• The difference between the number of valence
  electrons on the free atom and that assigned in
  the molecule.
• We count half the electrons in each bond as
  belonging to the atom.
• SO4-2
• Molecules try to achieve as low a formal charge
  as possible.
• Negative formal charges should be on
  electronegative elements.

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Assignment of e-

• 1. Lone pairs belong entirely to atom in
  question
• 2. Shared e- are divided equally between the
  two sharing atoms

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The sum of the formal charges of all atoms in a
species must equal the overall charge on the
species.
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A useful equation

• (happy-have) / 2 bonds
• POCl3 P is central atom
• SO4-2 S is central atom
• SO3-2 S is central atom
• PO4-2 P is central atom
• SCl2 S is central atom

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Exceptions to the octet

• BH3
• Be and B often do not achieve octet
• Have less than and octet, for electron deficient
  molecules.
• SF6
• Third row and larger elements can exceed the
  octet
• Use 3d orbitals?
• I3-

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Exceptions to the octet

• When we must exceed the octet, extra electrons go
  on central atom.
• ClF3
• XeO3
• ICl4-
• BeCl2

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If nonequivalent Lewis structures exist, the
one(s) that best describe the bonding in the
species has...
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FAVORED LEWIS STRUCTURES

• 1. formal charges closest to zero
• 2. negative formal charge
• is on the most electronegative atom

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EXAMPLES

• Carbon dioxide
• Thiocyanate ion
• Sulfate ion

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BOND ENERGY LENGTH
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Bond EnergiesE (Bonds Broken) (Bonds Made)
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Bonds form between atoms because bonded atoms
exhibit a lower energy.
Thus, energy is required to break bonds and
energy is released when bonds are formed.
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Bond Order bonds to a specific set
of elementsC-C the BO1CC the BO2C
C the BO3Fractions are possible
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COVALENT BONDS

• Bond Dissociation
• Energy
• Table 9.9 (text)

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Bond Energy (kJ/mol) H-F 565 H-Cl
432 H-Br 366 H-I 299
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Bond Energy (kJ/mol) Cl-Cl 242 Br-Br
193 I-I 151
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Bond Energy (kJ/mol)
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Bond Energy (kJ/mol)
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Use bond energies to predict ?Hc for acetylene
(C2H2).
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Energy
0
Internuclear Distance
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Energy
0
Internuclear Distance
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Energy
0
Internuclear Distance
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Energy
0
Internuclear Distance
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Energy
0
Bond Length
Internuclear Distance
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Energy
Bond Energy
0
Internuclear Distance
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Bond Energy Length
(kJ/mol) (pm)
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Bond Energy Length
(kJ/mol) (pm)
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Binary Ionic Compounds

• metal(s) non-metal (g) ---gt salt(s)

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Lattice Energy

• Energy change occurring when separated gaseous
  ions are packed together to form an ionic solid
• M(g) NM-(g) --gt M-NM

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What is the lattice energy of NaCl(s)? Na(g)
Cl-(g) ---gt NaCl(s)
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Lattice Energies

• LiCl 834 BeCl2 3004
• NaCl 769 MgCl2 2326
• KCl 701 CaCl2 2223
• Li2O 2799 BeO 4293
• Na2O 2481 MgO 3795
• K2O 2238 CaO 3414

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LE Lattice Energy
Where k proportionality constant dependent on
structure of solid and on electron configuration
of the ions
Where Q1 Q2 charges on the ions
Where r the shortest distance between the
centers of the cation and anion