Bonding

1
Chapter 8

• Bonding

2

• 8.1
• Types of Chemical Bonds

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What is a Bond?

• A force that holds atoms together.
• Why do bonds occur?
• Lower potential energy
• Higher stability.
• Bond energy is the energy required to break a
  bond.
• Why are compounds formed?
• Because it gives the system the lowest potential
  energy.

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Ionic Bonding

• Electrostatic attraction of two ions.
• Ions are formed by the loss or gain of electrons
• An atom with a low ionization energy (easy to
  remove e-) reacts with an atom with high electron
  affinity (more energy released).
• Opposite charges hold the atoms together.
• Metal Nonmetal

5
Coulomb's Law

• The energy of interaction between a pair of ions.
• E 2.31 x 10-19 J nm (Q1Q2/r)
• Q is the charge on ions.
• r is the distance between ion centers.
• If charges are opposite, E is negative,
  exothermic
• Attractive force
• Same charge, positive E, requires energy to bring
  them together, endothermic.
• Repulsive force

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As the two atoms approach each other..
Energy
0
Internuclear Distance
7
The energy begins to decrease
Energy
0
Internuclear Distance
8
Until to distance reaches 0.74Å
Energy
0
Internuclear Distance
9
And then begins to increase again due to
repulsions.
Energy
0
Internuclear Distance
10
Energy
0
Bond Length is measured where there is minimal
energy
Internuclear Distance
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Energy
Bond Energy
0
Internuclear Distance
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Covalent Bonding

• Electrons are shared by atoms.
• Nonpolar covalent bonds equal sharing of
  electrons.
• Polar covalent bonds unequal sharing of
  electrons.
• One end is slightly positive, the other negative.
• Indicated using small delta d.

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(No Transcript)
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When no electric field is present the molecules
are randomly oriented.
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-

When the field is turn on, molecules line up..
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• 8.2
• Electronegativity

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Electronegativity

• The ability of an atom to attract shared
  electrons to itself.
• Linus 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 pg 353 fig 8.3.
• Period Trends Increase left to right.
• Noble gases are excluded
• Group Trends Decreases as you go down a group.

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Electronegativity

• Difference in electronegativity between atoms
  will determine the type of bond.

Electroneg. 0 0.3 1.7
4.0 difference
Nonpolar Covalent
Bond Type
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Covalent Character decreases Ionic Character
increases
Electroneg. 0 0.3 1.7
4.0 difference
Nonpolar Covalent
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• 8.3
• Bond Polarity
• And
• Dipole Moments

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Dipole Moments

• All bonds can be classified by their polarity.
• How they are positioned in a molecule can result
  in molecular polarity.
• a negative charge and a positive charge is
  dipolar (two poles), or said to have a dipole
  moment.

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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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Geometry and polarity

• Three shapes will cancel them out.
• Linear

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Geometry and polarity

• Three shapes will cancel them out.
• Planar triangles

120º
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Geometry and polarity

• Three shapes will cancel them out.
• Tetrahedral

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Geometry and polarity

• Others dont cancel
• Bent

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Geometry and polarity

• Others dont cancel
• Trigonal Pyramidal

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• 8.4
• Ions
• Electron Configurations
• And Sizes

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Ions

• Atoms tend to react to form noble gas
  configuration.
• Metals lose electrons to form cations
• Nonmetals can share electrons in covalent bonds.
  Or they can gain electrons to form anions.

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Ionic Compounds

• We mean the solid crystal.
• Ions align themselves to maximize attractions
  between opposite charges, and to minimize
  repulsion between like ions.
• Can stabilize ions that would be unstable as a
  gas.
• Pictured in the margin of pg 358.
• React to achieve noble gas configuration

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Size of ions

• Various factors affect the size of ions
• Effective nuclear charge
• Loss of (or filling of) a sublevel or energy
  level
• Shielding effect
• Cations are smaller than the atoms they came
  from.
• Anions are larger.

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Size of ions

• Period Trends across a row they get smaller,
  and then suddenly larger.
• First half are cations
• Second half are anions.
• Group Trends Ion size increases down a group.

N-3
O-2
F-1
Li1
Be2
B3
C4
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Size of Isoelectronic ions

• Iso - same
• Isoelectronic ions have the same of electrons
• Al3, Mg2, Na1, Ne, F-1, O-2, and N-3
• All have 10 electrons.
• All have the configuration 1s22s22p6

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Size of Isoelectronic ions

• Positive ions have more protons so they are
  smaller.
• Size decreases as the nuclear charge increases

N-3
O-2
F-1
Ne
Na1
Al3
Mg2
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• 8.5
• Formation of Binary
• Ionic Compounds

38
Forming Ionic Compounds

• Lattice Energy - the energy released when one
  mole of ionic crystal is formed from its ions.
• M(g) X-(g) MX(s)
• Exothermic, negative

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Forming Ionic Compounds

• Ionic crystals have a lower potential energy and
  greater stability than the individual ions.
• The formation of an ionic compound consists of
  several intermediate steps and each step involves
  energy.

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Forming Ionic Compounds

• General rule
• Energy is absorbed to break bonds and released to
  form bonds.
• Energy is a state function so we can get from
  reactants to products in a round about way.

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Na(s) ½F2(g) NaF(s)

• First sublime Na Na(s) Na(g)
  DH 109 kJ/mol
• Ionize Na(g) Na(g) Na(g) e- Ei
  495 kJ/mol
• Dissociation F2 Bond ½F2(g) F(g)
• Ed 77 kJ/mol
• Affinity of F F(g) e-
  F-(g) Ea -328 kJ/mol

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Na(s) ½F2(g) NaF(s)

• Lattice energy Na(g) F-(g) NaF(s) EL
  -923 kJ/mol
• OVERALL RXN AND ENERGY
• Na(s) ½F2(g) NaF(s)
• ?E - 570 kJ/mole

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Calculating Lattice Energy

• Lattice Energy k(Q1Q2 / r)
• k is a constant that depends on the structure of
  the crystal.
• Due to a variable k value general trends are
  used to understand lattice energy
• As ionization energy ? lattice energy ?
• As the charge Q ? lattice energy ?

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• 8.6
• Partial Ionic character of Covalent Bonds

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Partial Ionic Character

• To calculate the ionic character
• Compare measured dipole of XY
• bonds to the calculated dipole of XY-
• for the completely ionic case.
• dipole Measured X-Y x 100
  Calculated XY-

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As the difference in Electronegativity increases
75
Ionic Character
50
25
the Ionic Character increases
Electronegativity difference
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You Can Not Reach 100 Ionic Character?

• What about polyatomic ions?
• An ionic compound will be defined as any
  substance that conducts electricity when melted.
• Also use the generic term salt.

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• 8.8
• Covalent Bond Energies
• And
• Chemical Reactions

49
Covalent Bond Energies

• Energy is also involved in the formation of
  covalent compounds in the form of bonds breaking
  and bonds forming.
• The bond energy (required to break bonds) is
  inversely proportional to bond length.

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Covalent Bond Energies

• Bond length is directly proportional to atomic
  size. (page 373)
• single bond, 1 pair of e- shared.
• double bond, 2 pair of e- shared.
• triple bond, 3 pair of e- shared.
• More bonds, shorter bond length.
• Bond energy and the enthalpy (?H) can be
  calculated.

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Using Bond Energies

• We can find change in enthalpy, DH, for a
  reaction.
• It takes energy to break bonds, endothermic, ().
• We use energy to form bonds, exothermic, (-).
• ?H ? D (bonds broken) - ? D (bonds formed)

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Find the energy for this
2 CH2 CHCH3
2NH3
O2




2 CH2 CHC º N
6 H2O
C-H 413 kJ/mol
O-H 467 kJ/mol
CC 614kJ/mol
OO 495 kJ/mol
N-H 391 kJ/mol
CºN 891 kJ/mol
C-C 347 kJ/mol
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• 8.9
• The Localized Electron Bonding Model

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Localized Electron Model

• A molecule is composed of atoms that are bound
  together by sharing pairs of electrons using the
  atomic orbitals of the bound atoms.
• Three Parts
• Valence electrons using Lewis structures
• Prediction of geometry using VSEPR
• Description of the types of orbitals (Chapt 9)

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• 8.10
• Lewis Structures

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Lewis Structure

• Shows how the valence electrons are arranged.
• One dot for each valence electron.
• A stable compound has all its atoms with a noble
  gas configuration.
• Hydrogen follows the duet rule.
• The rest follow the octet rule.
• Bonding pair is the one between the symbols.

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Rules

• Count up all valence e- for all atoms.
• Arrange the satellite atoms around the central
  atom (least electronegative) connected by a
  single dashed line. (each line represents two
  electrons)
• Place the rest of the valence e- around the
  satellite atoms to fill the octet rule (except
  H).
• COUNT and CHECK!!!
• If more valence e- are needed use multiple bonds.
• If electrons are left over place around the
  central atom.

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A useful equations

• (happy-have) / 2 bonds
• POCl3
• SO4-2
• SO3-2
• PO4-2
• SCl2

59

• 8.11
• Exceptions
• to the
• Octet Rule

60
Exceptions to the octet

• The second row elements, C, N, O, and F should
  always follow octet rule.
• The second row elements B and Be often have fewer
  than 8 electrons. These compounds are very
  reactive.

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Exceptions to the octet

• The octet rule because their valence orbitals (2s
  and 2p) will only hold 8 electrons.
• Third row elements and beyond will meet the octet
  rule and often exceed 8 electrons using empty d
  orbitals.

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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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• 8.12
• Resonance

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Resonance

• Sometimes there is more than one valid structure
  for an molecule or ion.
• NO3-

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Resonance

• NO2-
• Localized electron model is based on pairs of
  electrons, doesnt deal with odd numbers.

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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 one is
  most predominant.
• Trying to use the oxidation numbers to put
  charges on atoms in molecules doesnt work.

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Formal Charge

• Formal Charge
• ( valence e- on free atom)
• - (valence e- assigned)
• Valence e- assigned
• (number of lone pair electrons)
• ½ (number of shared electrons)

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Formal Charge

• SO4-2

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Formal Charge

• SO4-2
• Valence e- assigned to each O
• 6 ½(2) 7
• Formal Charge 6 7 -1
• The formal charge on each oxygen is -1

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Formal Charge

• SO4-2
• Valence e- assigned to each S
• 0 ½(8) 4
• Formal Charge 6 4 2
• The formal charge on each Sulfur is 2

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Formal Charge

• SO4-2

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Formal Charge

• SO4-2
• Valence e- assigned to single bond O
• 6 ½(2) 7
• Formal Charge 6 7 -1
• --------------------------------------------------
  -----
• Valence e- assigned to double bond O
• 4 ½(4) 6
• Formal Charge 6 6 0

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Formal Charge

• SO4-2
• Valence e- assigned to each S
• 0 ½(12) 6
• Formal Charge 6 6 0

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Formal Charges

• Atoms in molecules try to achieve a formal charge
  as close to zero as possible.
• This resonance structure is the most probable one

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Examples

• XeO3

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• 8.13
• Molecular Structure
• The VSEPR Model

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VSEPR

• Lewis structures tell us how the atoms are
  connected to each other.
• They dont tell us anything about shape.
• The shape of a molecule can greatly affect its
  properties.
• Valence Shell Electron Pair Repulsion Theory
  allows us to predict geometry

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VSEPR

• Molecules take a shape that puts electron pairs
  as far away from each other as possible.
• Have to draw the Lewis structure to determine
• electron pairs
• bonding
• nonbonding lone pair
• Lone pair take more space.
• Multiple bonds count as one pair.

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VSEPR

• The number of pairs determines
• bond angles
• underlying structure
• The number of atoms determines
• actual shape

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VSEPR
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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
82
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
84
No central atom

• You can predict the geometry of each angle.
• Just build it piece by piece.

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How well does it work?

• Does an outstanding job for such a simple model.
• Predictions are almost always accurate.
• Like all simple models, it has exceptions.