Bonding General Concepts

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Bonding General Concepts
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Types of Bonds
IONIC
COVALENT
e- are transferred from metal to nonmetal
e- are shared between two nonmetals
Bond Formation
Type of Structure
true molecules
crystal lattice
Physical State
liquid or gas
solid
Melting Point
low
high
Solubility in Water
yes
usually not
yes (solution or liquid)
Electrical Conductivity
no
Other Properties
odorous
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Types of Bonds
METALLIC
e- are delocalized among metal atoms
Bond Formation
Type of Structure
electron sea
Physical State
solid
Melting Point
very high
Solubility in Water
no
yes (any form)
Electrical Conductivity
malleable, ductile, lustrous
Other Properties
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Electronegativity The ability of anatom in a
molecule to attract shared electrons to itself.
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Bond Polarity

• Most bonds are a blend of ionic and covalent
  characteristics.
• Difference in electronegativity determines bond
  type.

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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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Coulombs Law
The energy of interaction between a pair of ions
is proportional to the product of their charges,
divided by the distance between their centers
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Table of Ion Sizes
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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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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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Formation of Binary Ionic Compounds

• Lattice energy
• Change in energy that takes place when separated
  gaseous ions are packed together to form an ionic
  solid.

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.

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Estimate ?Hf for Sodium Chloride
Na(s) ½ Cl2(g) ? NaCl(s)
Na(s) ? Na(g) 109 kJ
Na(g) ? Na(g) e- 495 kJ
½ Cl2(g) ? Cl(g) ½(239 kJ)
Cl(g) e- ? Cl-(g) - 349 kJ
Na(g) Cl-(g) ? NaCl(s) -786
kJ
Na(s) ½ Cl2(g) ? NaCl(s) -412 kJ/mol
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Definitions

• Bond energy
• Energy required to break bonds
• Bond length
• Distance that has the lowest energy available

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Bond Length Diagram
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Bond Length and Energy
Bonds between elements become shorter and
stronger as multiplicity increases.
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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 Polarity

• Nonpolar Covalent Bond
• e- are shared equally
• symmetrical e- density
• usually identical atoms

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Bond Polarity

• Polar Covalent Bond
• e- are shared unequally
• asymmetrical e- density
• results in partial charges (dipole)

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Bond Polarity

• Nonpolar
• Polar
• Ionic

View Bonding Animations.
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Dipole Moment

• Direction of the polar bond in a molecule.
• Arrow points toward the more e-neg atom.

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B. Determining Molecular Polarity

• Depends on
• Polarity of the bond
• molecular shape

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Determining Molecular Polarity

• Nonpolar Molecules
• Dipole moments are symmetrical and cancel out.

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Determining Molecular Polarity

• Polar Molecules
• Dipole moments are asymmetrical and dont cancel .

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Determining Molecular Polarity

• Therefore, polar molecules have...
• asymmetrical shape (lone pairs) or
• asymmetrical atoms

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Bond Energy and Enthalpy
Energy required
Energy released
D Bond energy per mole of bonds
Breaking bonds always requires energy
Breaking endothermic
Forming bonds always releases energy
Forming exothermic
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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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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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LE Model

• Description of the valence electron arrangement
  using Lewis structures.
• Predicts the geometry of the molecule using the
  VSEPR model.
• Description of the type of atomic orbitals used
  by the atoms to share electrons.

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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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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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Rules for Writing Lewis Structures

• Sum the valence electrons from all the atoms. Do
  not worry about keeping track of which electrons
  come from which atoms. It is the total number of
  electrons that is important.
• Use a pair of electrons to form a bond between
  each pair of bond atoms.
• Arrange the remaining electrons to satisfy the
  duet rule for hydrogen and the octet rule for the
  second-row elements.

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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.

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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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Molecular Structure

• VSEPR model
• Valence shell electron pair repulsion
• The structure around a given atom is determined
  principally by minimizing electron-pair
  repulsions.
• Lone pairs require more room than bonding pairs
  and tend to compress the angles between bonding
  pairs

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VSEPR Valence Shell Electron Pair
Repulsion
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VSEPR Valence Shell Electron Pair
Repulsion
A central atom
X atoms bonded to A
E nonbonding electron pairs on A
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VSEPR Linear
CO2
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VSEPR Trigonal Planar
BF3
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VSEPR Tetrahedral
CCl4
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VSEPR Trigonal Bi-pyramidal
SF4
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VSEPR Octahedral
SF6
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VSEPR Square Planar
XeF4
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What if I can make more than one Lewis structure?

• Formal Charge
• Hypothetical charge obtained by assuming that
  bonding electrons are equally shared between the
  two atoms involved in the bond. Lone pair
  electrons belong only to the atom to which they
  are bound.

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Formal Charge Rules

• Sum the lone pair electrons and ½ the shared
  electrons.
• Subtract this number from the number of electrons
  in ground state.
• To determine the more likely resonance structure
• FC should be as close to zero as possible.
• Negative charge should reside on the most
  electronegative and positive charge on the least
  electronegative element.

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Try it!
POCl3
BF3
SO42-
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Hybridization - The Blending of Orbitals


Poodle

Cocker Spaniel

Cockapoo


sp orbital


s orbital
p orbital
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What Proof Exists for Hybridization?
We have studied electron configuration notation
and the sharing of electrons in the formation of
covalent bonds.
Lets look at a molecule of methane, CH4.
Methane is a simple natural gas. Its molecule has
a carbon atom at the center with four hydrogen
atoms covalently bonded around it.
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Carbon ground state configuration
What is the expected orbital notation of carbon
in its ground state?
Can you see a problem with this?
(Hint How many unpaired electrons does this
carbon atom have available for bonding?)
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Carbons Bonding Problem
You should conclude that carbon only has TWO
electrons available for bonding. That is not not
enough!
How does carbon overcome this problem so that it
may form four bonds?
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Carbons Empty Orbital
The first thought that chemists had was that
carbon promotes one of its 2s electrons
to the empty 2p orbital.
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However, they quickly recognized a problem with
such an arrangement
Three of the carbon-hydrogen bonds would involve
an electron pair in which the carbon electron
was a 2p, matched with the lone 1s electron from
a hydrogen atom.
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This would mean that three of the bonds in a
methane molecule would be identical, because
they would involve electron pairs of equal
energy.
But what about the fourth bond?
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The fourth bond is between a 2s electron from the
carbon and the lone 1s hydrogen electron.
Such a bond would have slightly less energy than
the other bonds in a methane molecule.
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This bond would be slightly different in
character than the other three bonds in methane.
This difference would be measurable to a chemist
by determining the bond length and bond energy.
But is this what they observe?
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The simple answer is, No.
Measurements show that all four bonds in methane
are equal. Thus, we need a new explanation for
the bonding in methane.
Chemists have proposed an explanation they call
it Hybridization.
Hybridization is the combining of two or more
orbitals of nearly equal energy within the same
atom into orbitals of equal energy.
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In the case of methane, they call the
hybridization sp3, meaning that an s orbital is
combined with three p orbitals to create four
equal hybrid orbitals.
These new orbitals have slightly MORE energy than
the 2s orbital
and slightly LESS energy than the 2p orbitals.
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sp3 Hybrid Orbitals
Here is another way to look at the sp3
hybridization and energy profile
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sp Hybrid Orbitals
While sp3 is the hybridization observed in
methane, there are other types of hybridization
that atoms undergo.
These include sp hybridization, in which one s
orbital combines with a single p orbital.
This produces two hybrid orbitals, while leaving
two normal p orbitals
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sp2 Hybrid Orbitals
Another hybrid is the sp2, which combines two
orbitals from a p sublevel with one orbital from
an s sublevel.
One p orbital remains unchanged.
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Hybridization Involving d Orbitals
Beginning with elements in the third row, d
orbitals may also hybridize
dsp3 five hybrid orbitals of equal energy
d 2sp3 six hybrid orbitals of equal energy
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Hybridization and Molecular Geometry
A central atom
X atoms bonded to A
E nonbonding electron pairs on A
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Sigma and Pi Bonds
Sigma (?) bonds exist in the region directly
between two bonded atoms.
Pi (?) bonds exist in the region above and below
a line drawn between two bonded atoms.
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Sigma and Pi BondsSingle Bonds
1 ? bond
Ethane
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Sigma and Pi BondsDouble bonds
1 ? bond
Ethene
1 ? bond
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Sigma and Pi BondsTriple Bonds
1 ? bond
Ethyne
1 ? bond
1 ? bond
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The De-Localized Electron Model
Pi bonds (?) contribute to the delocalized model
of electrons in bonding, and help explain
resonance
Electron density from ? bonds can be distributed
symmetrically all around the ring, above and
below the plane.