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Unit 6: Theories of Covalent Bonding and Intro. To Organic Chemistry
Lewis Structures VSEPR Theories of Covalent Bonding
Valence Bond Theory Molecular Orbital Theory
Organic Chemistry Functional Groups Nomenclature Simple Reactions
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Chemical Bonds
Octet Rule: Atoms tend to gain, lose or share electrons until they are surrounded by eight valence electrons.
When ionic compounds are formed, electrons are gained or lost.
When molecular compounds are formed, electrons are shared.
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Chemical Bonds
Chemical bond: strong attractive force that exists between atoms (or ions) in a compound ionic bonds covalent bonds metallic bonds
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Chemical Bonds
Ionic Bond: the electrostatic force of attraction between oppositely charged ions in an ionic compound metal cation (+) non-metal anion (-)
The Na+ and Cl- ions in a salt (NaCl) crystal are held together by electrostatic attraction.
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Chemical Bonds
Covalent Bonds: the attractive force between atoms in a molecule that results from sharing of one or more pairs of electrons non-metals
H2O :
Cl2 :
H-O and Cl-Cl bonds result
from sharing of electrons
OH H
Cl Cl
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Lewis Symbols
Valence electrons are involved in chemical bonding: electrons residing in the incomplete
outer shell of an atom
For main group elements, the number of valence electrons for an element = group number of the element N (group 5A) has 5 valence electrons Br (group 7A) has 7 valence electrons
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Lewis Symbols
Lewis symbols (electron-dot symbols) are used to depict valence electrons in an atom or ion chemical symbol for the element dot for each valence electron
dots are placed on all 4 sides of the chemical symbol
all four sides of the symbol are equivalent
up to 2 dots (electrons) per side
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Lewis Symbols
Example: Draw the Lewis symbol for oxygen.
Example: Draw the Lewis symbol for carbon.
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Covalent Bonding
Lewis structures (also called electron-dot structures) can be used to represent the covalent bonds that are present in a molecule.
The formation of H2:
H + H H H or H H
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Covalent Bonding
Components of Lewis (electron-dot) structures:
Elemental symbol for each atom Bond between atoms depicted using a
solid line Unshared electron pairs are shown
around the appropriate atom
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Covalent Bonding
Single bond: one pair of shared electrons
Double bond Two pairs of shared
electrons
Triple bond Three pairs of shared
electrons
Cl Cl
O C OO C O
N N
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Drawing Lewis Structures
To draw a Lewis structure:
Add up the valence electrons from all atoms For a cation (+), subtract 1 electron for
each positive chargeNH4
+ : 5 + 4 (1) -1 = 8 e-
For an anion (-), add 1 electron for each negative chargeCN- : 4 + 5 + 1 = 10 e-
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Drawing Lewis Structures
Write the chemical symbols for each atom showing which is attached to which using a single bond (-). Sometimes (but not always) the order in
which the formula is writtenHCN: H-CN
Central atom (often written first) surrounded by other atomsCommonly the least electronegative element (except H) will be the central atom
CCl4 : C with 4 Cl attached to it
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Drawing Lewis Structures
Add electron pairs to the atoms bonded to the central atom first until each has an octet of electrons.
Remember, H only gets 2 electrons
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Drawing Lewis Structures
Place any leftover electrons on the central atom. Sometimes results in more than an
octet on the central atom
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Drawing Lewis Structures
If there are not enough electrons to give the central atom an octet, try multiple bonds. Use one (or more) unshared pairs of
electrons from an outer atom to form double (or triple) bonds
H C N H C N
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Drawing Lewis Structures
Example: Draw the Lewis structure for COCl2.
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Drawing Lewis Structures
Example: Draw the Lewis structure for the carbonate ion.
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Drawing Lewis Structures
Example: Draw all possible resonance structure for CO2.
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Drawing Lewis Structures
Example: Draw all possible resonance structures for SCN-. Which resonance structure is the major contributor to the resonance hybrid?
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VSEPR
Lewis structures show the number and type of bonds between atoms in a molecule. All atoms are drawn in the same plane
(the paper). Do not show the shape of the molecule
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VSEPR
The valence-shell electron-pair repulsion model (VSEPR) can be used to predict the shape of an ABn molecule when A is a main group element.
ABn
where A = central atom, main group elementB = outer atomsn = # of “B” atoms
Examples: CO2, H2O, BF3, NH3, CCl4
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VSEPR
VSEPR counts the number of electron domains around the central atom where electrons are likely to be found and uses this number to predict the shape.
Electron domains: regions around the central atom where electrons are likely to be found.
Two types of electron domains are considered: bonding pairs of electrons nonbonding (lone) pairs of electrons
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VSEPR
Bonding pairs of electrons: electrons that are shared between two atoms
Cl
Cl C Cl
Cl
Bonding pairs
Bonding pairs
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VSEPR
Nonbonding (lone) pairs of electrons: electrons that are found principally
on one atom unshared electrons
H N H
H
Nonbonding pair
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Ammonia (NH3) has 4 electron domains:
H N H
H
VSEPR
3 bonding pairs
1 nonbonding pair
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VSEPR
Electron domains tend to repel each other regions of high electron density like charges repel each other
According to VSEPR, the best arrangement of a given number of electron domains is the one that minimizes repulsions between them.
Electron domain geometry: The arrangement of electron domains
around the central atom
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Electron Domain Geometries
You must be able to draw these!
You must know these!
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Electron Doman Geometries
AB
B
B
A
B
BB
B
Trigonal planar
Tetrahedral
A
B
BBB
B
A
B
BB
B
B
B
Trigonal bipyramidal octahedra
l
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VSEPR
In order to determine the electron domain geometry: draw the Lewis structure count the total # of electron domains
multiple bonds = 1 electron domain determine the electron-domain
geometry
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VSEPR
Example: Predict the electron domain geometry of IF5.
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VSEPR
The electron domain geometry does NOT tell you the actual shape of the molecule.
Molecular geometry: the arrangement of the atoms in space
Molecular geometry is a consequence of electron-domain geometry.
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VSEPR
Example: Identify and draw the molecular geometry of IF5.
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VSEPR
Example: Identify and draw the electron domain and molecular geometries for I3
-.
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Valence Bond Theory
Covalent bonds form when atoms share electrons Electron density is concentrated
between the nuclei
Two common theories are used to explain the properties of molecules in terms of the bonding that exists between atoms. Valence bond theory Molecular orbital theory
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Valence Bond Theory
According to valence bond theory, an electron pair bond is formed between two atoms when a valence atomic orbital on one atom overlaps with a valence atomic orbital on another atom.
Overlap: share a region of space 2 electrons of opposite spin share
common space between the nuclei
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Valence Bond Theory
Overlap can occur between two s orbitals, two p orbitals, or one s and one p orbital:
Overlap between two s orbitals
overlap
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Valence Bond Theory
Overlap between one s and one p orbital
overlap
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Valence Bond Theory
Overlap between two p orbitals
overlap
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Valence Bond Theory
The previous bonds are called bonds. electron density is concentrated
symmetrically along an imaginary line connecting the two nuclei (internuclear axis)
Internuclear axis
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Valence Bond Theory
bonds are formed when sideways overlap occurs between two p orbitals that are oriented perpendicular to the internuclear axis. Overlap regions lie both above and
below the internuclear axis
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Valence Bond Theory
bonds are involved in the formation of double and triple bonds
Single bond:one bond
Double bond:one bond and one bond
Triple bond:one bond and two bonds
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Valence Bond Theory
Sometimes, simple overlap between s and/or p orbitals can’t explain the actual shape or properties of compounds.
Consider a BeF2 molecule:
F Be F
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Valence Bond Theory
F
1s22s22p5
F
1s22s22p5
Be
1s22s2
One unpaired e- One unpaired e-
No unpaired electrons
How can Be form covalent bonds with F if it doesn’t have unpaired electrons???
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Valence Bond Theory
Be
1s22s12p1
2 unpaired electrons
Promote one of the 2s e- to a 2p orbital
Be
1s22s2
no unpaired electrons
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Valence Bond Theory
F
1s22s22p5
F
1s22s22p5
Be
1s22s12p1
One unpaired e- One unpaired e-
2 unpaired electrons
Predicted overlaps:2s (Be) – 2p (F)2p (Be) – 2 p (F)
Implies two different kinds of Be – F bonds!
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Valence Bond Theory
Both of the Be – F bonds in BeF2 are identical!
The solution: Mix the Be 2s orbital with one of the Be
2p orbitals to form two hybrid orbitalsatomic orbitals formed by mixing 2 or more atomic orbitals on an atom
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Valence Bond Theory
Two sp hybrid orbitals are formed when one s and one p orbital are hybridized.
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Valence Bond Theory
When Be forms covalent bonds with two F, each sp hybrid orbital on the Be atom overlaps with a p orbital located on a F atom.
Be
1s sp 2p
F
1s22s22p5
F
1s22s22p5
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Valence Bond Theory
The BeF2 molecule is linear.
sp hybridization implies that the electron domain geometry around the central atom is linear.
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Valence Bond Theory
Other hybrid orbitals involving s, p, and d orbitals are possible.
s, p two sp hybrid orbitals s, p, p three sp2 hybrid
orbitals s, p, p, p four sp3 hybrid orbitals s, p, p, p, d five sp3d hybrid
orbitals s, p, p, p, d, d six sp3d2 hybrid
orbitals Know these!!!
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Valence Bond Theory
Each type of hybrid orbital is associated with a particular type of electron domain geometry. the same geometry that would be
predicted by VSEPR
You must know the electron domain geometry associated with each type of hybrid orbital.
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Valence Bond Theory
Hybrid Electron DomainOrbital Set Geometry
sp linearsp2 trigonal planarsp3 tetrahedral
sp3d trigonal bipyramidalsp3d2 octahedral
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Valence Bond Theory
Given the formula for an ABn compound, you must be able to: Draw the Lewis structure Determine the electron domain
geometry Sketch the electron domain geometry Identify the types of hybrid orbitals
present on the central atom
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Valence Bond Theory
Example: What is the electron domain geometry of AlH4
-? What type of hybrid orbitals are present?
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Valence Bond Theory
Example: What is the hybridization of the I atom in ICl2
-?