Bonding

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Bonding General Concepts
2
What is a Bond?

• A force that holds atoms together.
• 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.

3
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.

4
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

5
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.
• Most Noble gases do not have values.
• Difference in electronegativity between atoms
  tells us how polar.

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Electronegativity The ability of anatom in a
molecule to attract shared electrons to itself.
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Electronegativity difference
Bond Type
Zero
Covalent
Covalent Character decreases Ionic Character
increases
Intermediate
Large
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Ionic Bonds

• Electrons are transferred

• Electronegativity differences are
• generally greater than 1.7
• The formation of ionic bonds is
• always exothermic!

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Determination of Ionic Character
Electronegativity difference is not the final
determination of ionic character
Compounds are ionic if they conduct electricity
in their molten state
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Coulombs Law

• Q is the charge.
• r is the distance between the centers.
• If charges are opposite, E is negative
• exothermic
• Same charge, positive E, requires energy to bring
  them together.
• endothermic

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Size of ions

• Ion size increases down a group.
• Cations are smaller than the atoms they came
  from.
• Anions are larger.
• across a row they get smaller, and then suddenly
  larger.
• First half are cations.
• Second half are anions.

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Periodic Trends

• Across the period nuclear charge increases so
  they get smaller.
• Energy level changes between anions and cations.

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Table of Ion Sizes
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Size of Isoelectronic ions

• Iso - same
• Iso electronic 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

15
Size of Isoelectronic ions

• Positive ions have more protons so they are
  smaller.

N-3
O-2
F-1
Ne
Na1
Al3
Mg2
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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.
• React to achieve noble gas configuration

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List the following atoms in order of increasing
ionization energy Li, Na, C, O, F.

1. Li lt Na lt C lt O lt F
2. Na lt Li lt C lt O lt F
3. F lt O lt C lt Li lt Na
4. Na lt Li lt F lt O lt C
5. Na lt Li lt C lt F lt O

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Sodium losing an electron is an ________ process
and fluorine losing an electron is an _______
process.

1. endothermic, exothermic
2. exothermic, endothermic
3. endothermic, endothermic
4. exothermic, exothermic
5. more information needed

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19
Which of the following statements is true about
the ionization energy of Mg?

1. It will be equal to the ionization energy of Li.
2. It will be equal to and opposite in sign to the
   electron affinity of Mg.
3. It will be equal to and opposite in sign to the
   electron affinity of Mg.
4. It will be equal to and opposite in sign to the
   electron affinity of Mg2.
5. none of the above

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Choose the compound with the most ionic bond.

1. LiCl
2. KF
3. NaCl
4. LiF
5. KCl

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In which pair do both compounds exhibit
predominantly ionic bonding?

1. PCl5 and HF
2. Na2SO3 and BH3
3. KI and O3
4. NaF and H2O
5. RbCl and CaO

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22
Which of the following arrangements is in order
of increasing size?

1. Ga3 gt Ca2 gt K gt Cl gt S2
2. S2 gt Cl gt K gt Ca2 gt Ga3
3. Ga3 gt S2 gt Ca2 gt Cl gt K
4. Ga3 gt Ca2 gt S2 gt Cl gt K
5. Ga3 gt Ca2 gt S2 gt K gt Cl

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Which of the following species would be expected
to have the lowest ionization energy?

1. F-
2. Ne
3. O2-
4. Mg2
5. Na

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Sodium Chloride Crystal Lattice
Ionic compounds form solids at ordinary
temperatures.
Ionic compounds organize in a characteristic
crystal lattice of alternating positive and
negative ions.
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Forming Ionic Compounds

• Lattice energy - the energy associated with
  making a solid ionic compound from its gaseous
  ions.
• M(g) X-(g) MX(s)
• This is the energy that pays for making ionic
  compounds.
• Energy is a state function so we can get from
  reactants to products in a round about way.

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Calculating Lattice Energy

• Lattice Energy k(Q1Q2 / r)
• k is a constant that depends on the structure of
  the crystal.
• Qs are charges.
• r is internuclear distance.
• Lattice energy is greater with more highly
  charged ions.
• This bigger lattice energy pays for the extra
  ionization energy.
• Also pays for unfavorable electron affinity.

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Estimate ?Hf for Sodium Chloride
Na(s) ½ Cl2(g) ? NaCl(s)
Lattice Energy -786 kJ/mol
Ionization Energy for Na 495 kJ/mol
Electron Affinity for Cl -349 kJ/mol
Bond energy of Cl2 239 kJ/mol
Enthalpy of sublimation for Na 109 kJ/mol
Na(g) Cl-(g) ? NaCl(s) -786
kJ
Na(g) ? Na(g) e- 495 kJ
½ Cl2(g) ? Cl(g) ½(239 kJ)
Cl(g) e- ? Cl-(g) - 349 kJ
Na(s) ? Na(g) 109 kJ
Na(s) ½ Cl2(g) ? NaCl(s) -412 kJ/mol
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Lattice Energies of Alkali Metals Halides
(kJ/mol)
F- Cl- Br- I-
Li 1036 853 807 757
Na 923 787 747 704
K 821 715 682 649
Rb 785 689 660 630
Cs 740 659 631 604
Lattice Energies of Salts of the OH- and O2- Ions
(kJ/mol)
OH- O2-
Na 900 2481
Mg2 3006 3791
Al3 5627 15,916
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What about 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.
• They reach a distance with the lowest possible
  energy.
• The distance between is the bond length.

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Covalent Bonds
Polar-Covalent bonds

• Electrons are unequally shared
• Electronegativity difference between .3 and 1.7

Nonpolar-Covalent bonds

• Electrons are equally shared
• Electronegativity difference of 0 to 0.3

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Covalent Bonding Forces

• Electron electron
• repulsive forces

• Proton proton
• repulsive forces

• Electron proton
• attractive forces

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Bond Length Diagram
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The Octet Rule
Combinations of elements tend to form so that
each atom, by gaining, losing, or sharing
electrons, has an octet of electrons in its
highest occupied energy level.
Diatomic Fluorine
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Formation of Water by the Octet Rule
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Comments About the Octet Rule

• 2nd row elements C, N, O, F observe the octet
  rule.
• 2nd row elements B and Be often have fewer than 8
  electrons around themselves - they are very
  reactive.
• 3rd row and heavier elements CAN exceed the octet
  rule using empty valence d orbitals.
• When writing Lewis structures, satisfy octets
  first, then place electrons around elements
  having available d orbitals.

37
Lewis Structures

•  Shows how valence electrons are arranged among
  atoms in a molecule.
•  Reflects central idea that stability of a
  compound relates to noble gas electron
  configuration.

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Completing a Lewis Structure -CH3Cl

•   Make carbon the central atom

•   Add up available valence electrons

•   C 4, H (3)(1), Cl 7 Total 14

•   Join peripheral atoms
• to the central atom
• with electron pairs.

H
..
..
..
C
H
..

•   Complete octets on
• atoms other than
• hydrogen with remaining
• electrons

..
Cl
..
..
H
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Multiple Covalent BondsDouble bonds
Ethene
Two pairs of shared electrons
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Multiple Covalent BondsTriple bonds
Ethyne
Three pairs of shared electrons
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Resonance

•  Resonance is invoked when more than one valid
  Lewis structure can be written for a particular
  molecule.

Benzene, C6H6

•   The actual structure is an average of the
  resonance
• structures.

•   The bond lengths in the ring are identical,
  and
• between those of single and double bonds.

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Resonance Bond Length and Bond Energy

•  Resonance bonds are shorter and stronger than
  single bonds.

•   Resonance bonds are longer and weaker than
  double
• bonds.

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Resonance in Ozone, O3
Neither structure is correct.
Oxygen bond lengths are identical, and
intermediate to single and double bonds
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Resonance in Polyatomic Ions
Resonance in a carbonate ion
Resonance in an acetate ion
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Localized Electron Model
Lewis structures are an application of the
Localized Electron Model
L.E.M. says Electron pairs can be thought of as
belonging to pairs of atoms when bonding
Resonance points out a weakness in the Localized
Electron Model.
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Models
Models are attempts to explain how nature
operates on the microscopic level based on
experiences in the macroscopic world.
Models can be physical as with this DNA model
Models can be mathematical
Models can be theoretical or philosophical
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Fundamental Properties of Models

• A model does not equal reality.
• Models are oversimplifications, and are therefore
  often wrong.
• Models become more complicated as they age.
• We must understand the underlying assumptions in
  a model so that we dont misuse it.

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VSEPR Valence Shell Electron Pair
Repulsion
X E Overall Structure Forms
2 Linear AX2
3 Trigonal Planar AX3, AX2E
4 Tetrahedral AX4, AX3E, AX2E2
5 Trigonal bipyramidal AX5, AX4E, AX3E2, AX2E3
6 Octahedral AX6, AX5E, AX4E2
A central atom
X atoms bonded to A
E nonbonding electron pairs on A
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VSEPR Linear
AX2
CO2
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VSEPR Trigonal Planar
AX3
BF3
AX2E
SnCl2
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VSEPR Tetrahedral
AX4
CCl4
AX3E
PCl3
AX2E2
Cl2O
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VSEPR Trigonal Bi-pyramidal
AX5
PCl5
AX4E
SF4
AX3E2
ClF3
AX2E3
I3-
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VSEPR Octahedral
AX6
SF6
AX5E
BrF5
AX4E2
ICl4-