Download - Chapter 8
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Chapter 8
Bonding
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What is a Bond? A force that holds atoms together. Why? We will look at it in terms of energy. Bond energy: the energy required to
break a bond. Why are compounds formed? Because it gives the system the lowest
energy.
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Ionic Bonding An atom with a low ionization energy reacts
with an atom with high electron affinity.– A metal + a non metal (Which atom is
which?)– The electron moves to another atom.
(Which atom loses an electron to the other atom?)
– Opposite charges hold the atoms together.
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Coulomb's Law E= 2.31 x 10-19 J · nm(Q1Q2)/r
– Q is the charge on the ion.– r is the distance between the centers of
the atoms. If charges are opposite, E is negative
(WHY?)– exothermic
Same charge = positive E - requires energy to bring them together.
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What about covalent compounds?
The electrons in each atom are attracted to the nucleus of the other.– The electrons of the atoms repel each other,– The nuclei repel each other.
Compromise: They reach a distance with the lowest possible energy.
The distance between the two nuclei is the bond length.
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Internuclear Distance
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Internuclear Distance
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Internuclear Distance
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Internuclear Distance
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0
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Internuclear Distance
Bond Length
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0
Ener
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Internuclear Distance
Bond Energy
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Covalent Bonding Electrons are shared by atoms. “Ionic” and “covalent” bonds are actually
two extremes.– In between are polar covalent bonds.– The electrons are not shared evenly.– One end is slightly positive, the other
negative.– Indicated using small delta
When you think of polar compounds or molecules, which one comes to mind?
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H - F+ -
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H - F+ -
H - F+
-H - F+-
H - F
+-
H - F +-
H - F+-
H - F+
-
H - F +-
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H - F+ -
H - F+
-H - F+-
H - F
+-
H - F +-
H - F+-
H - F+
-
H - F +-
+-
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Electronegativity Def: The ability of an electron to attract
shared electrons to itself. Pauling method
–Imaginary molecule HX–Expected H-X energy = H-H
energy + X-X energy 2 = (H-X) actual - (H-X)expected
is the diff b/t actual and expected bond energy.)
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Electronegativity is known for almost every element.
–Gives us relative electronegativities of all elements.
–Tends to increase left to right.–decreases as you go down a group.
Most noble gases aren’t discussed. Difference in electronegativity between
atoms tells us how polar the bond is.
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Electronegativity difference
Bond Type
Zero
Intermediate
Large
Covalent
Polar Covalent
Ionic
Covalent C
haracter decreasesIonic C
haracter increases
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Dipole Moments A molecule with a center of negative
charge and a center of positive charge is dipolar (two poles – like a magnet), or has a dipole moment.–Center of charge doesn’t have to be
on an atom.–Will line up by charge in the
presence of an electric field.
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H - F+ -
H - F+
-H - F+-
H - F
+-
H - F +-
H - F+-
H - F+
-
H - F +-
+-
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How It is drawn
H - F+ -
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Which Molecules Have Dipoles? Any two atom molecule with a polar
bond. With three or more atoms there are
two considerations.1) There must be a polar bond.2) Geometry of the molecule can’t
cancel out the charge of a single bond.
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Geometry and polarity Three shapes will cancel them out. Linear
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Geometry and polarity Three shapes will cancel them out. Planar triangles
120º
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Geometry and polarity Three shapes will cancel them out. Tetrahedral
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Geometry and polarity Others don’t cancel Bent
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Geometry and polarity Others don’t cancel Trigonal Pyramidal
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Ions Atoms tend to react to form noble gas
configuration. Metals lose electrons to form cations Nonmetals can share electrons in
covalent bonds. –When two non-metals react.(more later)
Or they can gain electrons to form anions.
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Ionic Compounds We mean the solid crystal. Ions align themselves to maximize
attractions between opposite charges, and to minimize repulsion between like
ions. Can stabilize ions that would be unstable
as a gas. React to achieve noble gas configuration
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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. First half are cations. Second half are anions.
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Periodic Trends Across the period nuclear charge
increases so they get smaller. Energy level changes between anions
and cations.
Li+1
Be+2
B+3
C+4
N-3O-2 F-1
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Size of Isoelectronic ions Positive ions have more protons so
they are smaller.
Al+3
Mg+2
Na+1 Ne F-1 O-2 N-3
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Forming Ionic Compounds Lattice energy - the energy associated
with making a solid ionic compound from its gaseous ions.
M+(g) + X-(g) MX(s) This is the energy that “pays” for
making ionic compounds. Energy is a state function so we can get
from reactants to products in a round about way.
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Na(s) + ½F2(g) NaF(s) First sublime Na Na(s) Na(g)H =
109 kJ/mol Ionize Na(g) Na(g) Na+(g) + e-
H = 495 kJ/mol Break F-F Bond ½F2(g) F(g)
H = 77 kJ/mol Add electron to F F(g) + e- F-(g)
H = -328 kJ/mol
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Na(s) + ½F2(g) NaF(s) Lattice energy
Na+(g) + F-(g) NaF(s)H = -928 kJ/mol
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Calculating Lattice Energy Lattice Energy = k(Q1Q2 / r) k is a constant that depends on the
structure of the crystal. Q’s are charges. r is internuclear distance. Lattice energy is with smaller ions Lattice energy is greater with more
highly charged ions.
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Calculating Lattice Energy This bigger lattice energy “pays” for
the extra ionization energy. Also “pays” for unfavorable electron
affinity. O2-(g) is unstable, but will form as part
of a crystal
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Bonding
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Partial Ionic CharacterThere are probably no totally ionic
bonds between individual atoms.Calculate % ionic character.Compare measured dipole of X-Y
bonds to the calculated dipole of X+Y- the completely ionic case.
% dipole = Measured X-Y x 100 Calculated X+Y-
In the gas phase.
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% Io
nic
Cha
ract
er
Electronegativity difference
25%
50%
75%
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How do we deal with it? If bonds can’t be ionic, what are ionic
compounds?And what about polyatomic ions?An ionic compound will be defined as
any substance that conducts electricity when melted.
Also use the generic term salt.
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The Covalent BondThe forces that causes a group of atoms to
behave as a unit.Why?Due to the tendency of atoms to achieve the
lowest energy state. It takes 1652 kJ to dissociate a mole of CH4
into its ionsSince each hydrogen is hooked to the carbon,
we get the average energy = 413 kJ/mol
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CH3Cl has 3 C-H, and 1 C - Cl the C-Cl bond is 339 kJ/molThe bond is a human invention. It is a method of explaining the energy
change associated with forming molecules.
Bonds don’t exist in nature, but are useful.
We have a model of a bond.
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What is a Model?Explains how nature operates.Derived from observations. It simplifies them and categorizes the
information.A model must be sensible, but it has
limitations.
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Properties of a ModelA human inventions, not a blown up picture
of nature.Models can be wrong, because they are
based on speculations and oversimplification.
Become more complicated with age.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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Covalent Bond EnergiesWe made some simplifications in
describing the bond energy of CH4 Each C-H bond has a different energy.CH4 CH3 + H H = 435 kJ/molCH3 CH2 + H H = 453 kJ/molCH2 CH + H H = 425 kJ/molCH C + H H = 339 kJ/molEach bond is sensitive to its
environment.
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AveragesThere is a table of the averages of
different types of bonds pg. 365single 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, more bond energy, but
shorter bond length.
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Using Bond EnergiesWe can estimate H for a reaction. It takes energy to break bonds, and end
up with atoms (+).We get energy when we use atoms to
form bonds (-). If we add up the energy it took to break
the bonds, and subtract the energy we get from forming the bonds we get the H.
Energy and Enthalpy are state functions.
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Find the energy for this
2 CH2 = CHCH3
+
2NH3 O2+
2 CH2 = CHC N
+
6 H2O
C-H 413 kJ/molC=C 614kJ/molN-H 391 kJ/mol
O-H 467 kJ/molO=O 495 kJ/molCN 891 kJ/mol
C-C 347 kJ/mol
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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 Parts1) Valence electrons using Lewis structures2) Prediction of geometry using VSEPR3) Description of the types of orbitals
(Chapt 9)
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Lewis StructureShows 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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RulesSum the valence electrons.Use a pair to form a bond between
each pair of atoms.Arrange the rest to fulfill the octet rule
(except for H and the duet).H2OA line can be used instead of a pair.
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A useful equation(happy-have) / 2 = bondsCO2 C is central atomPOCl3 P is central atom
SO42-
S is central atom
SO32-
S is central atom
PO43-
P is central atom
SCl2 S is central atom
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Exceptions to the octetBH3 Be and B often do not achieve octetHave less than an octet, for electron
deficient molecules.SF6 Third row and larger elements can
exceed the octetUse 3d orbitals? I3-
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Exceptions to the octetWhen we must exceed the octet, extra
electrons go on central atom. (Happy – have)/2 won’t workClF3
XeO3
ICl4-
BeCl2
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ResonanceSometimes there is more than one valid
structure for an molecule or ion.NO3
- Use double arrows to indicate it is the
“average” of the structures. It doesn’t switch between them.NO2
- Localized electron model is based on pairs
of electrons, doesn’t deal with odd numbers.
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Formal ChargeFor molecules and polyatomic ions
that exceed the octet there are several different structures.
Use charges on atoms to help decide which.
Trying to use the oxidation numbers to put charges on atoms in molecules doesn’t work.
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Formal ChargeThe difference between the number of
valence electrons on the free atom and that assigned in the molecule or ion.
We count half the electrons in each bond as “belonging” to the atom.
SO4-2
Molecules try to achieve as low a formal charge as possible.
Negative formal charges should be on electronegative elements.
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ExamplesXeO3
NO43-
SO2Cl2
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VSEPRLewis structures tell us how the atoms
are connected to each other.They don’t 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
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VSEPRMolecules 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.
bondingnonbonding lone pairLone pair take 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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VSEPRElectron
pairsBond
AnglesUnderlyingShape
2 180° Linear3 120° Trigonal Planar
4 109.5° Tetrahedral
5 90° &120°
Trigonal Bipyramidal
6 90° Octagonal
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Actual shape
ElectronPairs
BondingPairs
Non-Bonding
Pairs Shape
2 2 0 linear3 3 0 trigonal planar3 2 1 bent4 4 0 tetrahedral4 3 1 trigonal pyramidal4 2 2 bent
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Actual Shape
ElectronPairs
BondingPairs
Non-Bonding
Pairs Shape
5 5 0 trigonal bipyrimidal5 4 1 See-saw5 3 2 T-shaped5 2 3 linear
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Actual Shape
ElectronPairs
BondingPairs
Non-Bonding
Pairs Shape
6 6 0 Octahedral6 5 1 Square Pyramidal6 4 2 Square Planar6 3 3 T-shaped6 2 1 linear
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Examples SiF4
SeF4
KrF4
BF3
PF3
BrF3
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No central atom Can predict the geometry of each
angle. build it piece by piece.
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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. Doesn’t deal with odd electrons
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Polar molecules Must have polar bonds Must not be symmetrical Symmetrical shapes include
–Linear–Trigonal planar–Tetrahedral–Trigonal bipyrimidal–Octahedral–Square planar