Post on 03-Jun-2018
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CM1502
Chapter 3
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Models ofchemical bonding
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Atomic properties and chemical bonds
Why do atoms bond at all?
Properties of an atom- Electronic config./ Zeff
Properties of a substance-Type/Strength of chemical bonds
less stable
more stable
P.E. = -K.E.
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Types of Bonding
Along the line IEs are about 8 eV
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Types of Bonding1. Metal with nonmetal:
ionic bonding
Metal loses electrons to form a positive ion (cation).
Non metal gains electrons to form a negative ion (anion).The electrostatic attraction between the ions draw them
into a three dimensional array to form an ionic solid.
Chemical formula is the empirical formula.
Movie (IVLE workbin-videos-Formation of ioniccompound)
2. Nonmetal with nonmetal:
covalent bonding
The atoms are drawn together as the nucleus of each
atom attracts the electrons of other.The electrons are shared.
Shared pair of electrons is localized between the atoms.
Chemical formula is the molecular formula
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Types of bonds
3. Metal with metal:
electron pooling and metallic bonding
Outer electrons of metals are losely
held due toshielding.These electrons move freely through
the entire piece of metal.
Hence electrons in metallic bonding
are delocalized.
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Lewis electron dot symbolSteps to write the Lewis symbol for main group elements:
1. Note the group number and is the same as their valence electrons.
2. Place one dot at a time on each of the four sides of the symbol.
3. Keep adding dots, pairing them until all are used up.
Octet rule for representative elements: When atoms bond they lose, gain or share
electrons to attain a filled outer level of eight electrons
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The Ionic Bonding Model
Lithium Fluoride: The three ways to depict the electron transfer
Electron configurations
Li 1s22s1
Orbital diagrams
Lewis electron-dot symbols
+ F 1s22s22p5 Li+ 1s2 + F 1s22s22p6
Li1s 2s 2p
F
1s 2s 2p
+
Li+
1s 2s 2p
F-
1s 2s 2p
+
.
+ F: ::
Li . Li+ + F::
:
:
Problem: Depict the formation of Potassium Oxide and Barium chlorideCM1502 Sem2-2013-14
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Ionic Bonding Begin with a neutral Cl
atom and a neutral Naatom.
Eionis the energy price
to pay to convert both intoits respective noble gasconfiguration.
The cation and anionattract each other, so atsome distance between
Na+and Cl-the potentialenergy of attraction winsover the positive valueEion.
The cation and anion repeleach other when close
because the e clouds ofboth ions can notinterpenetrate due to thePauli exclusion principle.
It is the strong Coulombattraction that binds allsalts
together.
Eion= IE(Na) EA(Cl)
Na Na++ e Cl- Cl + e
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Lattice energy
Li(s) Li(g) H =161 kJ
F2 F(g)
H = 79.5 kJLi(g) Li+
(g)+e- IE1= 520kJ
F(g)+e- F-(g) EA = -328 kJ
The electron transfer
process actually absorbsenergy !
Generally formation of ionic
solids releases energy.
Li+
(g)+ F-
(g) LiF(s) H0
= -1050kJ
The energy released when an
ionic solid is formed from its
ions is called Lattice energy.
Hence the overall reaction Li(s)+ F2 LIF(s)H = -617.5kJ; exothermic
Ionic solids exists only
because the lattice energy
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Periodic trends in Lattice energy
The lattice energy of L iF and MgO are in the ratio 1:4. Why?
Lattice energy = kQ1Q2
r
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k is the proportionality constant
Q1Q2are the charges on the ions
R is the shortest distance between thecenters of cation and anion
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Explaining the properties A typical ionic compound is hard and brittle.
This is because there is strong attractive forces that hold the ions
together.
They have high meltingand boiling points
They do not conduct electricity in solidstate but do so when melted or
dissolved.
This is because the ions can move in the molten or dissolved stateCompound mp (0C) bp (0C)
CsBr
661
1300
NaI
MgCl2
KBr
CaCl2
NaCl
LiF
KF
MgO
636
714
734
782
801
845
858
2852
1304
1412
1435
>1600
1413
1676
1505
3600
Solid ionic
compound
Molten ionic
compound
Ionic compound
dissolved in
water
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Why do ionic solids crack
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The covalent bonding model
The H2molecule
As the two hydrogen atoms are brought together,
1. The electrons in the two atoms repeleach other
because they have the same charge (E > 0);2. The protons in adjacent atoms repeleach other (E > 0);
3. The electron in one atom is attracted to the oppositely
charged proton in the other atom, and vice versa (E < 0);
At the observed bond distance the repulsive and attractiveinteractions are balanced
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Covalent bond formation in H2
Movie: IVLE-work bin-videos-covalent bond_2
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Properties of covalent bondBond order: is the number of electron pairs being shared by a
given pair of atoms.
Bond energy: Also called Bond enthalpy or bond
strength. It is defined as the standard enthalpy
change for breaking the bond in 1 mol of gaseousmolecules.
A-B (g) A(g)+ B(g) H = BEA-B(always >0)
Stronger bonds have higher bond energy and
weaker ones have lower bond energy
Bond length: The distance between the nuclei of the
bonded atoms
H H: H F: :::
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Relationship between bond order, bond length and bond energy.
For a given pair of atoms, a higher bond order results in shorter bond length
and higher bond energy
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Most covalent substances melt and boil at low temperatures
Pentane
The two forces in covalent molecules are
1. Strong bonding forces that hold the atoms together with the molecule and
2. Weak intermolecular forces that act between the molecules.
When the covalent liquid boils it is the weak forces between the molecules are
overcome and not the strong bonds within the molecule. Hence it needs less energy.
Explaining the Physical properties
Network covalent solids have exceptional behavior
Melts at 1550C Melts at 3550C
Most of the covalent substances are poor electrical conductors.
Reason: 1. electrons are localized
2. No ions are present
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Bond energy and enthalpy.
The heat released or absorbed during a chemical change is due
to differences between reactant and product bond energies.Horeaction = H
obonds broken+ H
obonds formed
Horeaction = reactant bonds broken- productbonds formed
H2(g) + F2(g) 2HF(g) Hreaction = -546kJ
Horeaction = [1x H-H + 1xF-F] [2xH-F]
=[432 + 159] [2x565]
= -539 kJ
Test your self:Calculate theHorxn for chlorination of methane to form chloroform.
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Bond energies and food/fuel
A fuel generally consist of C-H,C-C C-O and O-H bonds.
Of which C-H and C-C bonds are weakerand C-O and O-H bonds are stronger.
Generally a fuel reacts with O2, all the bonds break and form C=O and O-H bonds.
If a fuel consist of many C-C and C-H bonds and fewer C-O , O-H bonds, said to
release Higher energy and these are known as good fuels.
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Between the two extremes
EN is the relative ability of a bonded atom to attract shared electrons.
432kJ/mol
H-H + F-F 2 H-F
159kJ/mol Expected BE of H-F is 296kJ/mol
Actual BE of H-F is 565kJ/mol
The reason for increased BE is Electrostatic attraction
Arbitrary cutoff divides ionicfrom covalent bonds
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Polar covalent bonds
Figure 9.23
EN
3.0
2.0
0.0
Electrons are not transferred
completely.
Electrons are not shared
equally.
One atom has a stronger
attraction for the sharedelectron than the other atom.
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Bond polarity and dipole moment
A molecule that has a positive center of charge of
magnitude Q and a negative center of charge ofmagnitude Q separated by a distance R has a
dipole moment() of QR
Dipole moment = = QR
SI Unit = C m (coulomb meter)
Often used unit = debye
1debye(D) = 3.336 x 10-30Cm)
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Molecules with polar bonds and a dipole moment
Molecules with polar bonds and no resulting dipole moment
http://www.cengage.com/chemistry/book_content/9781111580650_zumdahl/images/ch13/13p604_f05c.htmlhttp://www.cengage.com/chemistry/book_content/9781111580650_zumdahl/images/ch13/13p604_f05b.html8/12/2019 CM1502 chapter 3 2013-14
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Metallic bonding
All the metal atoms in the sample contribute their valenceelectrons to form a delocalized electron sea.
The piece is held together by the mutual attractionof themetal cations and the mobile electrons.
The metal ion array is regular but not rigid.
Not rigid hencecannot come under ionic bonding
No localized sharing - hence cannot be covalent bond.
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Properties of metalsMelting points are only moderately high.
Reason: the cations can move without
breaking the attraction to thesurrounding
Boiling points are very high
Reason: Higher energy is needed to break
the cation from all the valence electrons.
Periodic trends:
M.Pt decrease down the group.
Reason: Larger metal ions have a weaker
attraction to the electron sea.
M.Pt increase across the period.
Reason: The charge on the cation increases
from left to right, hence stronger attraction
towards the electron sea.
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Why do metals dent and bend?
When hammered, the metalsslide past each other through
the electron sea and end up in
new positions.
Properties of metals
Metals are good electrical and thermal conductorsin both
solid and liquid states because of their mobile electrons.
Presence of foreign atoms disrupt the array and reduce the
conductivity.
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Summary
Atomic properties and type of bond
Features of ionic bonding
lattice energy
properties-ionic compounds are brittle, high melting, conduct
electricity only in molten/dissolved state.
Features of covalent bonding
non metals
bond order bond energy and bond length
Polar covalent compounds
Features of metallic bonding
electron sea metals bend, have high melting and boiling points and conduct
electricity.
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Role of molecular shape
Interactions of reactants during a chemical reaction isbased on the molecular shapes.
Predict the physical and chemical behavior of syntheticmaterials.
Molecular shape is a crucial property of living systems
Eg: Hormonal regulation and function of genes.
Sugar Enzyme 28CM1502 Sem2-2013-14
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What is Lewis Structure?
Two-dimensional structural formula consists
of electron-dot symbols.
It shows which are the atoms bonded to
each other but it does NOT indicate
the three-dimensional shape.
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Molecular formula to Lewis structure.
Molecularformula
Atom
placement
Sum of
valence e
Remaining
valence e
Lewis
structure
Place atom with lowest
EN in center
Draw single bonds. Subtract2e for each bond.
Give each atom 8e
(2efor H)
Step 1
Step 2
Step 3
Step 4
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Molecular
formula
Atom
placement
Sum of
valence e-
Remaining
valence e-
Lewis
structure
NF3
N
FF
F
N 5e-
F 7e- X 3 = 21e-
Total 26e-
:
: :
::: :
:
:
:
31
Electrons involved in bonding are called bond
pairs . These are shared between atoms.
Unshared electrons are called lone pairs .They
belong to only one atom.
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L i St t 1 0
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Lewis Structures v1.0
An atoms valency should be satisfied.
N2must be a triple bond in order to complete theoctet.
O2a double bond in order to complete the octet. F2a single bond in order to complete the octet.
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More Lewis Structures
Note the first period nonmetal (H) completedshell is 2 electrons [He].
CO2can be readily explained. CO, however, can satisfy the octets, but the
valencyof C?
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Valency Rule CAN be Broken
Octets all completed,
valencies all satisfied
Octets all completed,
but C and O valency
broken.
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Alert! Issue Formal ChargePatch.
Lewis Structures v1.1
For nonmetals, when an atoms valency is broken, and theoctet rule is satisfied, the atom must possess either feweror greater electrons than it has valence electrons.
The difference is accounted for by placing a formal chargeon the offending atoms.
So in
This is because C owns 5 electrons in the above
structure but should own 4, and O owns 5 electrons, butshould own 6.
movie (IVLE /workbin/videos/Formal charge calculation)
- +
35
Formal charge of an atom =
No. of valence e- no. of lone pair e- no. of bonding e-
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Octets Not Always Completed
If a molecule has an odd number of electrons,there is no way the octet rule can hold, or thebond order (BO) can be an integer.
NO (nitric oxide) is an example.
Here the BO is 2 and is consistent
with the bond dissociation energy (De)of 6.52 eV
H2+also exists, and must
have a BO of .We will also see later that
He2+also exists and has a
BO also of !
Octet satisfiedOctet notsatisfied
Octets satisfied
BO is 2 But need weird
formal charges
-+
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Lewis Structures with Multiple Bonds.
Step 5: Itfollows the other steps in Lewis structure
construction. If a central atom does not have 8e-, an octet,then a lone pair of e-can be moved to form a multiple bond.
CCH
H H
H
37
CCH
H H
H
:
H C C H
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E di th O t t
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Expanding the Octet Unless required, or stated otherwise, we always stick to the
octet rule.
We are forced in molecules like SO2to violate the octetrule in order to be consistent with the experimentalobservation of equal bond lengths.
The octet on S expands.
S is in group 6, so up to 6 valence electrons can beinvolved in the bonding.
This is what occurs here. Note that there are normal octet structures one can write for SO2,
but these structure do not account for the experimental fact thatboth S-O bond lengths are identical.
The expansion occurs for third and higher period elements.
Other examples include SF6, SF4, PCl5.38
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Alert! Issue Expanded OctetPatch.
We Now Have Lewis Structures v1.2
For nonmetalsin the third and higher periodsit may not bepossible to satisfy the octet rule.
In these cases, we are permitted to expand the octet.
We are allowed to expand the octet in these elementsbecause, allegedly, dAOin the same valence shell (nquantum number) are not so high in energyfor theseelements and can thus participate in bonding.
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Al t! I R P t h
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O
O O
O
O O
O
O O
A
B
C
O
O O
A
B
C
Alert! Issue ResonancePatch.
Lewis Structures v1.3
The experimental fact was the bond lengths in O3are identical.
Hence the correct description of O2is not given by any one ofthe two Lewis structures individually but by the superposition of
the two, called a resonance hybrid.
The BO is 1 for each bond. 40
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Resonance Structures
They have the SAMErelative placement
of atoms BUT different locations ofbonding and lone electron pairs.
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Benzene C6H6
H
H
H
H
H
H
H
H
H
H
H
H
Experimentally, the bonds in benzene are all of equal length,
between a single and double bond.
The structure of benzene can also be written as......
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Benzene C6H6
Experimentally, the bonds in benzene are all of
equal length, between a single and double bond.
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Criteria for Choosing the More Important Resonance
structure:
Smaller formal charges are preferred over larger ones
(eg, 0 is preferred over -1 or +1)
The same nonzero formal charges on adjacent atoms are
not preferred
A more negative formal charge should reside on a moreelectronegative atom
Predicting Structures/shapes of molecules
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Predicting Structures/shapes of molecules
Lewis structures determine the bond pairs and lone
pairs of electrons among the atoms.
Combining this information with VSEPR(Valence Shell
Electron Pair Repulsiontheory) enables us to predict
the shapes of the molecules.
The basis of VSEPR is that the repulsions between
electrons in bonds and lone pairsdetermines the
overall shape of a molecule
VSEPR assumes that core electrons make no
significant impact on the shape of a molecule, so can
be ignored.45CM1502 Sem2-2013-14
VSEPR R l
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VSEPR Rules1. Draw Lewis structures to determine the bonding
around atoms.
2. Assign an electron group arrangement:- To do this we count the number of atoms
around the central atom and add it to the
number of lone pairs it possesses.
3. Predict the overall geometry around the central atom.
O in water has two hydrogen atomsbonded to it, and two lone pairs. So
there is 2+2=4 different directionselectrons are more localized in.
S in SO2has two oxygen atoms bondedto it, and a single lone pair. So there is2+1=3 different directions electrons are
more localized in.46
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VSEPR Rules
4. The final predicted structure is governed by the idea
that electrons in lp repel electrons found in other lpthe most,and electrons found in bonds repel other
electrons found in bonds the least,with lp electrons
repelling bonding electrons intermediatebetween the
above two.
5. Draw and name the molecular shape by counting
bonding groups and nonbonding groups separately.
6. Predict the bond angle.
movie (IVLE /workbin/videos/VSEPR theory47
VSEPR Structure Predictions
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VSEPR Structure Predictions
Electrons in 2 directions
Electrons in 3 directions
Electrons in 4 directions
Electrons in 5 directions
Electrons in 6 directions
Geometry about
central atom
Possible structures
of the molecule
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AXmEn Geometry structure examplesAX2 lilnear linear CS2,HCN, BeF2
AX3 Trigonal planar Trigonal planar SO3BF3NO3-CO32-
AX2E Trigonal planar V shaped SO2O3PbCl2SnBr2AX4 Tetrahedral Tetrahedral CH4SiCl4So4
2-ClO4-
AX3E Tetrahedral Trigonal pyramidal NH3PF3ClO3-H3O
+
AX2E2 Tetrahedral Bent H2O OF2SCl2
AX5 TBP TBP PF5AsF5SOF4
AX4E TBP Seesaw SF4XeO2F2IF4+IO2F2-
AX3E2 TBP Tshaped ClF3BrF3
AX2E3 TBP linear XeF2I3- IF2
-
AX6 Octahedral Octahedral SF6IOF5
AX5E Octahedral Squarepyramidal BrF5TeF5- XeOF4
AX4E2 Octahedral Square planar XeF4ICl4-
A - central atom, X - surrounding atom, E -lone pairs, m, n - integers 49CM1502 Sem2-2013-14
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Molecular shapes with more than one central atom
ethane
CH3CH3
ethanol
CH3CH2OH
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Conclusions
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Conclusions
VSEPR Each group of electrons around a
central atom remains as far away from the othersas possible.
5 common geometric shapesresult when 2, 3, 4,
5 or 6 electron groups surround a central atom.
Lone pairs and double bonds exert greater
repulsions.
Bond polarity and molecular shape determine
molecular polarity.
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