Transcript Chapter 8

Bonding: General Concepts
Chapter 8
Types of Chemical Bonds
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Bond Energy is _____________________
_________________________________.
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Compounds always strive to be at the lowest energy
orientation possible. They will bond or not bond
depending on whether or not it is favorable from an
energy perspective.
A higher value for bond energy indicates a stronger
bond.
Ionic Bonding
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An ionic bond forms when electrons are
_________ between atoms. When this
happens both atoms become _____; one
positive (____ of e-), the other negative
(_____ of e-).
• This type of bonding occurs between a ______
and a _____________.
• Elements bonded by an ionic bond form
ionic compounds.
Coulomb’s Law
E=
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Describes _____________ between two
particles.
_______ values indicate attraction, while
_______ values indicate repulsive forces.
Ion pairs have lower energy (more favorable)
when they are ______________________.
Bond Length
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Bond length is the distance between two ions
that maximizes the favorable attractions and
minimizes the amount of repulsive forces.
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Favorable attractive forces:
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Unfavorable attractive forces (repulsive forces):
Covalent and Polar Covalent Bonds
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A Covalent Bond occurs when two nuclei
_________ electrons.
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Occurs between two _____________ atoms.
Polar Covalent bonds occur when two
nuclei _________________________.
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The dipoles (+ and – centers) that are formed give the
covalent bond more ionic character than normal.
Lower case delta (d) is used to indicate partial charge.
Electronegativity
Follows the trend for electron affinity
(synonymous), which means that it
increases as we move ________ a group
and _________ a period on the periodic
table.
 Electronegativity is __________________
_________________________________.
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Electronegativity affects bonds
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If bonded atoms have a large difference in electronegativity
(2.0 or greater), they are considered to have an _________
bond because _____________________________________.
If bonded atoms have a moderate difference in electronegativity
(0.5-1.6), they are considered to have a _______-__________
bond because _____________________________________.
If bonded atoms have a negligible difference in electronegativity
(below 0.5), they will be considered to have a ___________
bond because ______________________________________.
If bonded atoms have a difference in electronegativity between
1.6 and 2.0, their identities have to be considered. If a metal is
involved, it will be deemed ionic. If 2 non-metals are bonded, it
will be considered polar-covalent.
Bond Polarity and Dipole Moments
Dipoles are formed in _________ bonds.
 If the dipoles (a form of force) are aligned to
point in exactly equal and opposite
directions, they will _________________.
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This makes the molecule non-polar.
The bond remains polar.
Molecules must be symmetrical for the cancellation to
occur. Shapes that cancel are (draw them):
1.
2.
3.
Assignment
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Arrange the bonds in each of the following
set in order of increasing polarity
 C-F,
O-F, Be-F
 O-Cl, S-Br, C-P
 C-S, B-F, N-O
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Using only the periodic table as a guide,
select
A) the most electronegative element in group 6A
B) The least electronegative element out of Al, Si, P
C) The element in the group K, C, Zn, F that is most
likely to form an ionic compound with Ba.
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Give three ions that are isoelectronic with
argon. Place these ions in order of
increasing size.
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What two requirements must be satisfied for
a molecule to be polar?
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Rank the following bonds in order of
increasing ionic character:
 N-O,
Ca-O, C-F, Br-Br, K-F
Section 5: Formation of Binary Ionic Compounds
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Lattice energy: the energy required to
combine elements to form an ionic
compound. (amount released when bond
forms)
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Lattice energy is negative.
Elements must start as a gas
The more exothermic (negative), the more likely the
substance is to form spontaneously.
Steps involved in forming ionic
bonds from elements.
Vaporization of elements. (endo)
Li(s)  Li(g)
F2(g)  2F(g)
 Ionization of elements. (exo)
Li(g)  Li+ (g) + eF(g) + e-  F-(g)
 Formation of solid by combination of ions.
(very exo)
Li+(g) + F-(g)  LiF(s)
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Energy Diagram
Partial Ionic Character of Covalent
Bonds
% ionic character 
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measured dipoleX Y
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calculateddipoleX Y
Any Compound That
Conducts Electricity
When Melted
Is IONIC
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x100
Models of Chemical Bonds
Models do not equal reality…they are
merely something to help us visualize a
concept near the truth.
 Models are often wrong because they
over-simplify.
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Covalent Bond Energy
Single bond = one pair shared electrons
 Double bond = two pair shared electrons
 Triple bond = three pair shared electrons
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As more pairs of electrons are shared, the bond length
shortens. (more orbitals have to overlap to allow the
sharing to happen)
Single bonds usually contain the least amount of energy,
while triple bonds usually contain the most…as bond
length shortens, bond energy increases.
Page 374 Tables 8.4 and 8.5
Covalent Bond Energy Calculation
DHbond = sum energies required to break
old bonds (positive signs) plus the sum of
energies required to form new bonds
(negative signs)
 DHbond = S bonds broken – S bonds formed
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Use values in the tables on pg 374.
Covalent Bond Energy Calculation
Using the bond energies listed in Table 8.4,
calculate the DH for the reaction of methane
with chlorine and fluorine to give Freon-12
(CF2Cl2).
CH4(g) + 2Cl2(g) + 2F2(g)  CF2Cl2(g) + 2HF(g) + 2HCl(g)
The VSEPR Model
Valence Shell Electron Pair Repulsion
 A model of molecular structure based on
the idea that ideal structures minimizes
electron pair repulsions.
 Draw and evaluate Lewis Structures
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Molecular Geometry Models
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We look at the molecular geometry of a single
atom, not of an entire molecule.
Constituent groups are the things bonded to the
atom under scrutiny.
Dashed lines represent a bond behind the plane
of the paper; wedged lines represent a bond
coming toward you (in front of the paper plane)
Planar Geometry
Linear
Trigonal Planar
Bent
1-2 Constituents
0 Lone Pair
3 Constituents
0 Lone Pair
2 Constituents
1 Lone Pair
Bond Angle: 180o
Bond Angle: 120o
Bond Angle: <120o
Tetrahedral and Derivatives
Tetrahedral
Trigonal Pyramidal
Bent
4Constituents
0 Lone Pair
3 Constituents
1 Lone Pair
2 Constituents
2 Lone Pair
Bond Angle: 109.5o
Bond Angle: 107.3o
Bond Angle: 104.5o
Trigonal Bipyramidal and Derivatives
Trigonal Bipyramidal
See-Saw
T-Shaped
Linear
5 Constituents
0 Lone Pair
4 Constituents
1 Lone Pair
3 Constituents
2 Lone Pair
2 Constituents
3 Lone Pair
Bond Angle:
90o, 120o
Bond Angle:
<90o, <,120o, <180o
Bond Angle:
<90o, <180o
Bond Angle:
180o
Octahedral and Derivatives
Octahedral
Square Pyramidal
Square Planar
6 Constituents
0 Lone Pair
5 Constituents
1 Lone Pair
4 Constituents
2 Lone Pair
Bond Angle: 90o
Bond Angle: <90o
Bond Angle: 90o
Hybridization of Orbitals
aka Localized Electron Model
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Natural orbitals overlap to form hybridized orbitals
during bonding.
2 types of bonding
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Sigma (s bonds = take equatorial positions, can be hybridized.
Pi (p) bonds = take axial positions, exist in non-hybridized orbitals
Possibilities for hybridization:
sp3
sp2
dsp3
d2sp3
sp
How to become Hybrid
Orbitals in which bonding electrons exist will
become degenerate (equal energy).
2p ___ ___ ___
__ __ __ __ sp3
2s ___
By becoming equal energy, they also acquire
same shape.
sp3 Hybridization
Atoms with 4 constituent groups
 Tetrahedral shapes
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sp2 hybridization
3 constituent groups
 Trigonal planar
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sp2 continues
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One p orbital remains un-hybridized. This
orbital has the ability to house lone pairs
as well as form a double bond (pi bond!)
sp hybridization
2 constituents
 Linear
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Now 2 extra p remain un-hybridized.
 2 double bonds, or 1 triple bond can form.
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dsp3 and d2sp3 hybridization
dsp3 allows for 5 constituent groups
 trigonal bipyramidal
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d2sp3 allows for 6 constituent groups
 octahedral
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Molecular Orbital Theory
This is an alternative to the idea of
hybridized orbitals.
 Essentially, when two atoms come
together their orbitals merge and form two
completely new orbitals.
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 Bonding
orbital (low energy, electrons fill first)
 Anti-bonding orbital (high energy, electrons fill last)
Homonuclear Diatomic
Molecular Orbitals
The number of
molecular orbitals
should equal the
number of orbitals you
start with.
 In the case of p orbitals,
one is a sigma orbital,
and two are pi orbitals.
 Ones with stars are
anti-bonding orbitals.
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Determining Bond Order
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You can use molecular orbitals to
determine bond order to a much more
specific degree than you can with a Lewis
Dot Diagram.
Delocalized Electrons
Alternating double and single bonds
around a molecule can lead to available
unused p orbitals that create a cloud of
available space for electrons to sit in.
 This is very stable.
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Complex Ions and
Coordination Compounds
Highly colored compounds
 Paramagnetic (unpaired electrons)
 Made of a Transition Metal and Ligands
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How do they do that?
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Transition metals have two sets of valence
electrons.
Primary valence electrons affect oxidation
state (charge in ions) and allow for trading of
electrons (ionic bonds).
 Secondary valence electrons affect
coordination number (number of ligands) and
allow for ligands to share electrons with the
transition metal.
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Ligands
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Ligands are molecules or ions that have a
lone pair of electrons available to
coordinate with the transition metal.
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Rule of Thumb: The coordination number
of a metal can generally be assumed as
twice its charge.
 Ex:
Co+3 can have a coordination number of 6.
Naming Complex Ions
Name cations first
 Change ending of ligands to –o
 Use prefixes to tell how many.
 List oxidation state of metal using Roman
Numerals
 List ligands in alphabetical order
 If complex ion is an anion, end metal
name with –ate.
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Examples
[Al(OH)4] [Co(NH3)4 ]2+
 Amminetetraaquachromium(II) sulfate
 Potassium hexacyanoferrate(III)
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Isomers
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A set of compounds with the same
chemical formula that exhibit decidedly
different properties.
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Structural Isomer = the same atoms are present in each
molecule, but they are bonded to different things.
Stereoisomer = bonds are between same atoms, but
have a different spatial relationship with each other.
Concept Map
ISOMERS
STRUCTURAL
Coordination
Isomers
Linkage
Isomers
STEREOISOMERS
Geometric
Optical
These are
the ones we
care about.
Geometric Isomers
Come in two forms: cis and trans
 cis- prefix represents that similar
constituent groups are on the SAME side.
 trans- prefix represents that similar
constituent groups are on OPPOSITE
sides.
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More Complicated cis-, trans-
Optical Isomers
Two forms of the molecule have different
effects on plane polarized light. (one form
creates destructive interference)
 These isomers are non-super imposable
mirror images.
 Each isomer is referred to as an
ENANTIOMER.
 The chemical formula is said to have
CHIRALITY.
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