Principles of Chemistry

1
Principles of Chemistry
CHAPTER 12
CHEMICAL BONDING
2
Types of Chemical Bonds

• Bond - a force that holds groups of two or more
  atoms together
• atoms function as a unit
• Bond energy - energy required to break a bond
• Ionic bonding - Strong Bonding
• Result from the opposite attractions of charged
  ions
• Some atoms readily give up electrons

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3
Ionic Bond
e-
Nonmetal-
Metal
Nonmetal
Metal
Ionic Compound
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4
Bonding between similar atoms

• Situation arises with atoms of similar
  electronegativities interact with one another
• Wont form an Ionic bond
• Ionic bonds require one atom to give up electrons
  and the other to have a strong affinity for
  electrons
• Covalent bond - electrons are shared by both
  atoms nuclei
• electrons are attracted to simultaneously by both
  protons

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5
Covalent Bond
H

H
H
H
Covalent molecule
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Bonding between similar atoms

• Polar Covalent Bond - Electrons are shared as
  they are in a covalent bond however they are
  shared unequally
• One atom is slightly more electronegative
• Electrons tend to spend more time around one atom
  than another
• One atom gains a slight positive charge (d), the
  other gains a slight negative charge (d-)
• These slight charges occur because of unequal
  electron sharing

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7
Polar Covalent Bond
H

H
F
F
d
d-
Polar Covalent Molecule
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8
Electronegativity

• Electronegativity - The relative ability of an
  atom in a molecule to attract shared electrons to
  itself
• Explains unequal electron sharing
• Ionic vs. Covalent Bonding
• Periodic Trends
• Top Right - Greatest electronegativity
• Bottom Left - Least electronegativity
• General trend with exceptions (Ru, Rh, Pd, etc)

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9
Review Periodic Organization
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10
Electronegativity

• Electronegativity
• Determined by measuring polarities of the bonds
  between various atoms
• Polarity of a bond depends on the difference
  between the electronegativity values of the atoms
  forming the bond

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11
Bond Polarity Determination
Values from Figure 12.3
H - H S - H Cl - H O - H
F - H
Least Polar
Greater Polarity
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12
Bond Polarity and Dipole Moments

• Dipole Moment - A polar molecule has a center of
  positive charge and a center of negative charge
• Represented by arrow points
• Arrow points in the direction of the negative
  charge center
• Arrows tail (Fletching) indicates the positive
  center of charge

H
F
d
d-
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13
Bond Polarity and Dipole Moments

• Dipole Moment
• Any diatomic molecule that has a polar bond
• Some polyatomic molecules have dipole moments
• Water as a polar Molecule
• Molecules are attracted to one another
• Causes high boiling point
• Dissolve ionic compounds

2d-
O
H
H
d
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14
Red High Electron DensityBlue Low Electron
Density
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15
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16
Review Ne electron configuration
1s
10
2s
2p
Ne
3d
3s
3p
4d
4s
4p
4f
Neon
5d
5s
5p
5f
..
electrons
2
2
6



10
Ne


..
1s2
configuration
2s2
2p6
valence electrons
2s2
2p6

8
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17
Stable Electron Configurations
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18
Stable Electron Configurations

• Main group metals form ions by losing enough
  electrons to achieve the configuration of the
  previous nobel gas
• Nonmetals form ions by gaining enough electrons
  to achieve the configuration of the next noble gas

In almost all stable chemical compounds of the
representative elements, all of the atoms have
achieved a noble gas electron configuration
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19
Review Noble gases
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20
Review Noble gases
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21
Predicting Formulas

• When a nonmetal and a group 1, 2 or 3 metal react
  to form a binary ionic compound, the ions form in
  such a way that the valence-electron
  configuration of the nonmetal is completed to
  achieve the configuration of the next nobel gas
  and the valence orbitals of the metal are emptied
  to achieve the configuration of the previous
  nobel gas.
• When two nonmetals react to form a covalent bond,
  they share electrons in a way that completes the
  valence-electron configurations of both atoms.
  That is, both nonmetals attain noble gas electron
  configurations by sharing electrons

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22
Predicting Formulas
Ca Ar4s2O He2s22p4
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Predicting Formulas
Ca Ar4s2O He2s22p4
Metal
Nonmetal
Ca - 2e- ? Ca2O 2e- ? O2-
Review
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24
Predicting Formulas
Ca Ar4s2O He2s22p4
Metal
Nonmetal
Ca Ar4s2 ? 1s22s22p63s23p6 2e-O He2s22p4
2e- ? 1s22s22p6
Ca2O2-
Ca2O2- 0 net charge CaO empirical formula
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25
Predicting Formulas
Ca Ar4s2O He2s22p4
Metal
Nonmetal
Ca Ar4s2 ? 1s22s22p63s23p6 2e-O He2s22p4
2e- ? 1s22s22p6
Ca2O2-
Valence 8
Ca2O2- 0 net charge CaO empirical formula
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26
Ionic Structures

• When Metals and Nonmetals react
• Large amounts energy required to break the bonds
• Resulting compound is very stable
• Empirical formula
• Simple ratio NaCl
• Real structures of NaCl consists of huge equal
  numbers of Na and Cl- packed together

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27
Ionic Structures
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28
Polyatomic Ions

• Binary atomic ions KCl, NaCl, KBr, etc
• Polyatomic ions NH4, NO3-
• Charged species
• Individual Polyatomic ions are held together by
  covalent bonds
• Ammonium ion has four N - H covalent bonds
• Ammonium species as a whole is charged

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29
Common Polyatomic Ions
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30
Lewis Structures

• Bonding involves just the valence electrons
• Lewis Structure - is a representation of a
  molecule that shows the valence electrons
  arranged among the atoms in the molecule
• From the idea most important requirement for the
  formation of a stable compound is that the atoms
  achieve noble gas electron configurations
• Also known as Lewis dot diagrams and electron dot
  diagrams

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31
Lewis Structures

• Hydrogen forms stable molecules when it shares
  two electrons
• H - H
• Hellium does not form bonds because its valence
  orbital is already filled.
• He 1s2

.
.
H
H
2H 1s1
H2 1s2 Filled its valence shell
H H
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32
Lewis Structures

• The second row nonmetals carbon through fluorine
  form stable molecules whne they are surrounded by
  enough electrons to fill the valence orbitals.
• These electrons obey the octet rule - surrounded
  by 8 electrons

..
..
..
..
.
.
.
.
.
.
.
Cl

Cl

Cl

Cl

C
.
..
..
..
..
..
Cl


..
..
..
C 1s22s22p2 4 valence electronsCl
1s22s22p63s23p5 7 valence e-



Cl

Cl
C
..
..
..
Cl


..
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Lewis Structures

• Neon does not form bonds because it already has
  an octet of valence electrons filled.

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34
Lewis Structures

• We must include all the valence electrons from
  all atoms
• the total number of valence electrons is the sum
  of all valence electrons from all the atoms in
  the molecule
• Atoms that are bound to each other share one or
  more pairs of electrons
• The electrons are arranged so that each atom is
  surrounded by enough electrons to fill the
  valence orbitals of that atom. (2 e- for
  hydrogen, 8 e- for second row nonmetals)

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35
Molecules with Multiple Bonds

• Sum up the valence electrons
• Form a bond between the carbon and each oxygen
• Distribute the remaining electrons to achieve the
  Octet Rule (with the exception of Hydrogen)

Try CO32-
C 1s22s22p2 (4)O 1s22s22p4 (6)2e- (2)
C O O O 2e- (4) (6) (6) (6) (2) 24
v.e. remember because there are 2 additional
electrons this molecule will have a 2- charge
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Molecules with Multiple Bonds

• Form a bond between the carbon and each oxygen

2-
C
O
O


6 v.e.
O
24 a.e. - 6 18 v.e.
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37
Molecules with Multiple Bonds

• Distribute the remaining electrons to achieve the
  Octet Rule (with the exception of Hydrogen).
  Check work. This is trial and error.

8
..
2-


..
..
8


8
C
O
O
..
..
26 v.e.
O


..
8
24 a.e. - 26 v.e. -2 v.e. We have used too many
Valence Electrons!
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38
Molecules with Multiple Bonds

• Distribute the remaining electrons to achieve the
  Octet Rule (with the exception of Hydrogen).
  Check work. This is trial and error.

8
..


..
2-
8

8
C
O
O
..
..
24 v.e.
O


..
8
24 a.e. - 24 v.e. 0 v.e.
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39
Resonance Structures

• You may have noticed that 3 structures are
  possible

8
..


..
2-
8

8
C
O
O
..
..
8
O


..
..


..
2-
8
C
8


8
O
O
..
..
8
O

..
..


..
2-
8

8
C
O
O
8
..
..
O


..
8
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40
Resonance Structures

• More than one Lewis structure can be drawn for
  the molecule
• In this course we are not concerned with choosing
  the best resonance structure
• Important topic in more advanced chemistry courses

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41
Some exceptions to the octet rule

• Many elements usually obey the octet rule
• If an atom that can exceed the octet rule is
  bound to atoms that obey the octet rule
• the a.e. should be distributed to form octets on
  all atoms
• the remaining a.e. can be placed on atoms that
  can exceed the rule
• Boron tends to form compounds in which the Boron
  atom has fewer than 8 electrons.

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42
Molecular Structure

• Molecular Structure or Geometric Structure
• Different perspective than Lewis Structures
• Refers to the 3-D arrangement of the atoms
• Water as an example
• 3-D shape is bent or V-shaped
• 105 bond angle
• CO2
• Linear bond angle, 180

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43
Molecular Structure

• Trigonal Planar - Planar or flat angle
• such as BF3
• bond angle 120
• Tetrahedral Structure
• 3-D Triangle with 4 faces CH4

F
..
..
O
B
C
O
O
..
..
H
H
F
F
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44
VSEPR Model

• The structures of molecules play an important
  role in determining their properties
• Taste is related to structure
• Chemical receptors in Mammals are related to
  structure
• Predict approximate structure of a molecule using
  VSEPR model
• Valence Shell Electron Pair Repulsion Model
• Predicting structures formed from nonmetals

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45
VSEPR Model

• Main Idea the structure around a given atom is
  determined by minimizing repulsions between
  electron pairs
• Bonding and Nonbonding electron pairs should be
  positioned as far apart as possible

CO2 Best Arrangement places the pairs 180
degrees away from each other
..
..
C
O
O
..
..
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VSEPR Model

• Considering (4) atoms as in the case of this
  example, simply dividing the atoms on a flat
  surface would separate the atoms 90 degrees away
  from one another
• Is there any possible way to increase the space?

4 atoms Best Arrangement use 4-D space, this
increases the spacing from 90 degrees to 109.5
degrees This is known as a Tetrahedral Arrangement
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47
VSEPR Rules

• Two pairs of electrons on a central atom in a
  molecule are always placed 180 apart. This is a
  linear arrangement of pairs
• Three pairs of electrons on a central atom in a
  molecule are always placed 120apart. This is a
  trigonal planar arrangement of pairs
• Four pairs of electrons on a central atom in a
  molecule are always placed 109.5apart. This is a
  tetrahedral arrangement of electron pairs.

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48
VSEPR Rules

• When every pair of electrons on the central atom
  is shared with another atom, the molecular
  structure has the same name as the arrangement of
  electron pairs.
• When one ore more of the electron pairs around a
  central atom are unshared (lone pairs), the name
  for the molecular structure is different from the
  arrangement of electron pairs

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Molecular Structure Molecules with Double Bonds

• Each double bond in this molecule acts
  effectively as one repulsive unit
• Think of bonds as a cloud between two atoms
• Each double bond should be treated as a single
  bond in VSPR
• 4 electrons in a double bond are tied together

..
..
C
O
O
..
..
C
O
O
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VSEPR Summery

• When using the VSEPR model to predict the
  molecular geometry, a double bond is counted the
  same as a single electron pair

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References

• Zumdahl, Steven. DeCoste, Donald. Introductory
  Chemistry A foundation. Sixth Edition. Houghton
  Mifflin Company. ISBN 0-618-80327.
• Wikipedia (some graphics)

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