Chemical Bonding

1
Chemical Bonding

• Chapters 8 and 9

2
Chemical Bonds

• What is a bond?
• A force that holds atoms together
• We will look at it in terms of energy
• Bond Energy is the NRG required to break a bond
• Why are compounds formed?
• Bonds give the system the lowest NRG

3
Bond Energy is the energy required to break a
bond.

• REMEMBER
• It always takes energy to break a chemical bond.
• AND,
• To form a bond, requires a lowering of energy.

4
Types of Bonds

• Ionic Bonds electrostatic forces that exist
  between ions of opposite charge
• Covalent Bonds sharing of electrons between two
  atoms
• Metallic Bonds found typically in transition
  metals each atom is bonded to several
  neighboring atoms bonding electrons are
  relatively free to move throughout the structure
  of the metal.

5
Ionic Bonding

• An atom with a low ionization energy reacts with
  an atom with high electron affinity.
• Typically a metal with a nonmetal
• The electron transfers atoms.
• The ions each achieve a Noble Gas electron
  configuration low energy state.
• Opposite charges hold ions together.

6
Crystal Lattice

• A repeating and ? crystal lattice results.

NaCl, sodium chloride
7
Ionic Compounds

• Makes a solid crystal.
• Ions align themselves to maximize attractions
  between opposite charges,
• and to minimize repulsion between like ions.
• Chemical formula is actually the empirical
  formula, called the formula unit
• ClNaClNaClNaClNaClNa
• NaClNaClNaClNaClNaCl NaCl
• ClNaClNaClNaClNaClNa

8
Features of Ionic Compounds

• Brittle, hard, crystalline solids at room
  temperature
• High melting points
• Metals bonded to non-metals
• Elements from opposites sides of the Periodic
  Table
• Dissolve in water to form ions
• Conduct an electric current in water

9
Coulombs Law

• Expresses the NRG of interaction between a pair
  of ions.
• E 2.31 x 10-19 J nm (Q1Q2) / r
• E energy of interaction between a pair of ions
  (in Joules)
• r distance (in nm) between ion centers
• Q1 and Q2 charges of the ions
• Opposite charges means (E)
• Endo or Exo? What does that mean about NRG in
  the system?

10
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.

11
Periodic Trends

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

N-3
O-2
F-1
B3
Li1
C4
Be2
12
Size of Isoelectronic ions

• Positive ions have more protons so they are
  smaller.
• A stronger charged nucleus pulls the electrons
  inward.

N-3
O-2
F-1
Ne
Na1
Al3
Mg2
13
Lattice Energy

• The energy release that occurs when separated
  gaseous ions are packed together to form an ionic
  solid
• Xx(g) Y-y(g) ? XyYx (s) energy
• Lattice NRG (Eel) k(Q1Q2)/r
• k constant
• Q1 and Q2 charges of ions
• r distance between ion centers

14
Example

• Which has the most exothermic lattice energy,
  NaCl or KCl?
• NaCl why?
• Since both have the same charges (1 and -1), the
  distance between the charges needs to be
  considered. Since Na is smaller than K, the
  distance between the centers of Na and Cl is less
  and therefore has a greater lattice NRG.

15
Example 2 and 3

• Arrange the following ionic compounds in order of
  increasing lattice energy NaF, CsI, and CaO
• CsI lt NaF lt CaO
• Which substance would you expect to have the
  greatest lattice energy, AgCl, CuO or CrN?
• CrN

16
Covalent Bonds

• Bond is a force which causes a group of atoms to
  behave as a single unit.
• Electrons are shared by atoms.
• Electron orbitals must overlap.

17
The Covalent Bond

• 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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Energy
0
Internuclear Distance
19
Energy
0
Internuclear Distance
20
Energy
0
Internuclear Distance
21
Energy
0
Internuclear Distance
22
Energy
0
Bond Length
(They reach a distance apart with the lowest
possible energy.)
Internuclear Distance
23
Energy
Bond Energy
0
Internuclear Distance
24
Features of Covalent Compounds (aka Molecular
Compounds)

1. Electrons are shared, (Single, Double or Triple
   Bonds are possible.)

2. Non-metals bonded to Non-metals.
3. Includes all diatomic molecules.
4. Relatively low melting/boiling points.
5. No repeating formula, particle is a single
unit called a molecule.
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How We Represent Covalent Compounds
1. Molecular Formulas
CH4
2. Structural Formulas
H H - C - H H
Show the bonds between the atoms.
26
Covalent vs. Ionic Bonding

• Covalent is sharing, ionic is stealing.
• Totally different from each other.
• Reality Check There are many compounds which
  exhibit both traits!
• These are called polar covalent bonds.
• The electrons are shared, but shared unequally.

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Polar Covalent Bonds

• The electrons are shared, but they are not shared
  evenly.
• One atom has the pair more often than the other.
• A polar molecule results One end is slightly
  positive, while the other is slightly negative.
• A partial charge is called a dipole.

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Example
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31
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How Do We Know Which Atom Has the Electron Pair
More Often?

• Answer Electronegativity Values
• Electronegativity - The ability of an atom to
  attract shared electrons to itself.
• The more electronegative an atom, the more often
  it has the shared pair.
• Greater electronegativity ? pole

33
Electronegativity

• E.N. values are assigned for almost every element
  (Figure 8.6, p. 285)
• Gives us relative electronegativities of all
  elements.
• Tends to increase left to right and decreases as
  you go down a group.
• Noble gases arent discussed.
• Difference in electronegativity between atoms
  tells us how polar.

34
Helpful Number Line

• Determine the Electronegativity Value difference
  between two atoms

Non-polar covalent
Polar covalent
ionic
0 0.4
2.0 3.4
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Electronegativity difference
Bond Type
Zero
Covalent
Covalent Character decreases Ionic Character
increases
Intermediate
Large
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Polar Covalent Bond-How it is drawn
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Reminder on Writing Formulas and Nomenclature

• Ionic Compounds
• Name the cation then name the anion
• When writing formulas, add subscripts to make
  sure that the charges balance
• Covalent Compounds
• When naming, use prefixes for the subscripts, the
  2nd atom will end in ide
• Write formulas, assign subscripts based on the
  prefixes

38
Lewis Dot Structures are Models to represent both
ionic and covalent

• What is a Model?
• Explains how nature operates.
• Derived from observations.
• It simplifies and categorizes the information.
• A model must be sensible, but it has limitations.

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Properties of a Model

• A human invention, not a blown up picture of
  nature.
• Models can be wrong, because they are based on
  speculations and oversimplification.
• You must understand the assumptions in the model,
  and look for weaknesses.
• We learn more when the model is wrong than when
  it is right.

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Lewis Structures

• Show 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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Electron Dot Placement
2
6
X
3
5
1
7
4
8
The X represents the symbol for the element.
The dots are placed around the symbol in the
order shown above.
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Rules

1. Sum the of ALL the valence electrons.
2. Determine the central atom. The least
   electronegative element is central (H never
   central, C nearly always central)
3. Write symbols for the atoms to show which atoms
   are attached and connect them with a single bond
   (a dash)
4. Complete the octets of the atoms attached to the
   central atom (except for H, follows a duet rule).
5. Place any leftover electrons on the central atom.
   Not enough electrons - consider a double or
   triple bond.

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A useful equation

• ( happy - have ) ? 2 bonds
• (what they want - what they have) ? 2
  bonds

H2O (12 - 8) ? 2 2 bonds
? ? H ? O ? H ? ?
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Practice Structures

• PCl3
• ICl
• CH4
• CH2Cl2
• HCN
• NO
• CO2
• H2

• NH3
• C2H4
• BrO3-1
• O2
• ClO2-1
• OH-1
• PO4-3

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Partial Ionic Character

• There are probably no totally ionic bonds between
  individual atoms.

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75
Ionic Character
50
25
Electronegativity difference
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How do we deal with it?

• If bonds cant be ionic, what are ionic
  compounds?
• An ionic compound will be defined as any
  substance that conducts electricity when melted.
• Also use the generic term salt.
• As it turns out, most compounds fall somewhere
  between ionic and covalent.

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• The bond is a human invention.
• It is a method of explaining the energy change
  associated with forming molecules.
• Bonds dont exist in nature, but are useful.
  
• We have a model of a bond.

H - Cl
49
Exceptions to the octet

• Less than an Octet
• BH3
• Be and B often do not achieve octet, but form
  highly reactive compounds
• More than an Octet
• SF6 and I3-
• Third row and larger elements can exceed the
  octet.
• How? Use 3d orbitals
• Odd Number of Electrons
• NO, NO2, and ClO2

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Exceptions to the octet

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

51
Resonance Structures

• When more than one valid Lewis structure can be
  written for a particular molecule.
• Actual structure is an average of the depicted
  resonance structures
• Drawn by writing the variant structures connected
  by a double-headed arrow

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What to do when more than one Lewis structure
works? Use Formal Charge

• Assign formal charges on atoms to help decide
  which is best.
• Trying to use the oxidation numbers to put
  charges on atoms in molecules doesnt work.
• Molecules try to achieve as low a formal charge
  as possible.
• Negative formal charges should be on the most
  electronegative elements.

53
Formal Charge

• Number of valence electrons on the free atom
  minus number of valence electrons assigned to the
  atom in the molecule
• Lone pair e- belong to atom in question
• Shared e- are divided equally between the sharing
  atoms
• The sum of the formal charges of all atoms in a
  given molecule or ion must equal the overall
  charge on that species
• If the charge on an ion is -2, the sum of the
  formal charges must be -2.

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Resonance Examples

• NO3-1
• SO3
• HCO2-1

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Bond Lengths are Averages

• Have made a table of the averages of different
  types of bonds pg. 305
• single bond one pair of electrons is shared.
• double bond two pair of electrons are shared.
• triple bond three pair of electrons are shared.
• More bonds, shorter bond length.

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How do they share electrons?Localized Electron
Model

• Simple model, easily applied.
• 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

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VSEPR the 3-D shape

• 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

58
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.
• count bonding pairs
• count nonbonding (lone) pairs
• Lone pair take up 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
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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
63
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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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.
• Well spend some time in this class drawing
  structures and making models to better understand
  the different VSEPR models

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Hybridization

• The mixing of two or more atomic orbitals of
  similar energies on the same atom to produce new
  orbitals of equal energies.
• Creates Hybrid Orbitals
• Methane goes from 2s2 2p2 to 2sp3
• Draw the sub orbitals according to Hunds rule

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Types of Hybrid Orbitals

• Two orbital sets sp
• Three orbital sets sp2
• Four orbital sets sp3
• Five orbital sets sp3d
• Six orbital sets sp3d2

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Bond Types

• Sigma Bonds (s)
• Bond in which the electron pair is shared in an
  area centered on a line running between the atoms
  
• Lobes of bonding orbital point toward each other
• Pi Bonds (p)
• Electron pair above and below the s bond
• Created by overlapping of non-hybridized p
  orbitals

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• Single bonds consist of one s bond
• Double bonds consist of one s bond and one p bond
• Triple bonds consist of one s bond and 2 p bonds