Molecular structure and bonding
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Transcript of Molecular structure and bonding
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Molecular Structure and Bonding
Dr.Christoph Phayao University March 2015
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Part 1
What is a chemical bond ?
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Ionic Bond
Normally between a metal and a non-metal: They exchange electrons and become ions (charged atoms) which attract each other by electrostatic force.
A pair of ions does not stay alone but form crystals
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Covalent Bond
Two non-metals share (valence) electrons:
(Remark: Transition metals can form covalent bonds also !)
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Polar Covalent Bond
Two non-metals share electrons unevenly because of electronegativity difference. Electrons are closer to one atom than the other.
This results on partially negative and positive charges on the atoms
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Metallic Bond
Metal atoms share all their valence electrons, which freely move between all atoms which form a network.
Therefore all metals can conduct electricity and look shiny
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Bond Polarity
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Polar Bonds
Uneven sharing of electrons due to differences in Electronegativity
The “pull” an atom has for electrons
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Electronegativity Trends
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Common Electronegativites
Highest value, set to 4
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“Calculate” EN The first idea of EN came from L.Pauling – he estimated EN from bond energies
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Allred-Rochow EN
Example: Flourine (r = 72 pm) Carbon (r = 77 pm) Calculate the AR electronegativity
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Shielding of 2p electrons: Flourine: S = 6 * 0.35 + 2 * 0.85 = 3.8 => Z* = 9 – 3.8 = 5.2 EN = 3590 * 5.2/(72 2) + 0.744 = 4.35 Carbon: S = 3 * 0.35 + 2 * 0.85 = 2.75 => Z* = 6 – 2.75 = 3.25 EN = 3590 * 3.25 / (77 2) ) + 0.744 = 2.71
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Calculate Dipole moments
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Polar Molecules
Electrons are not equally shared in a bond, which can lead to a dipole moment of the whole molecule
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Polar Bonds and Geometry
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Which bond type ?
(exception: Transition metals !)
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Electron counting
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Formal Charge Split all bonds in the middle => “real” charge on atoms
(2) Octet Rule Count all bonding electrons for one atom => 8 is most stable
(3) Oxidation Number Give all bonding electrons to the more electronegative atom
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Special Cases
“Extended octet” Especially P and S can use d-orbitals to make more than 3 resp. 2 bonds !
6 VE: Especially common for B and Al !
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Part 2: Valence Bond Theory (VB)
“Valence Electrons are located in bonds and lone pairs”
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Sigma bonds
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Pi Bonds
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“Resonance”
Write the resonance formula for OZONE ! Does the molecule have a charge ?
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Important exception: Carbon Monoxide !
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Ho
mew
ork (2
)
Draw Lewis structure(s) and find formal charges (all atoms) and hybridization (central atom) in: 1. NO3 (-) 2. PO4 (3-) 3. CH3 Cl 4. CH2Cl2 5. SO2 6. SO3 7. CO3 (2-) 8. H2O2 9. N2O 10. Cl O2
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***** Break *****
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Prentice Hall © 2003 Chapter 9
• Atomic orbitals can mix or hybridize in order to adopt an appropriate geometry for bonding.
• Hybridization is determined by the electron domain geometry.
sp Hybrid Orbitals
• Consider the BeF2 molecule (experimentally known to exist):
Hybrid Orbitals
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Figure 11.2 The sp hybrid orbitals in gaseous BeCl2.
atomic
orbitals
hybrid
orbitals
orbital box diagrams
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Figure 11.2 The sp hybrid orbitals in gaseous BeCl2(continued).
orbital box diagrams with orbital contours
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Figure 11.3 The sp2 hybrid orbitals in BF3.
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sp2 and sp3
Hybrid Orbitals
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Figure 11.4 The sp3 hybrid orbitals in CH4.
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Figure 11.5 The sp3 hybrid orbitals in NH3.
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Figure 11.5 continued The sp3 hybrid orbitals in H2O.
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Including d-orbitals
3d orbitals can be filled as well => Al acts as Lewis acid => P and S have “hypervalence”
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Figure 11.6 The sp3d hybrid orbitals in PCl5.
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Figure 11.7 The sp3d2 hybrid orbitals in SF6.
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SOLUTION:
PROBLEM: Describe the types of bonds and orbitals in acetone, (CH3)2CO.
PLAN: Use the Lewis structures to ascertain the arrangement of groups and
shape at each central atom. Postulate the hybrid orbitals taking note of
the multiple bonds and their orbital overlaps.
H3C
C
CH3
O
sp3 hybridized
sp3 hybridized
CC
C
O
H
H
HHH
H
sp2 hybridized
bonds bond
CC
C
O
sp3
sp3
sp3
sp3
sp3
sp3
sp3
sp3
sp2 sp2
sp2
sp2
sp2sp2
H
HH
HH
H
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Tasks
• Draw the Lewis Structures and the Hybrid Orbitals for Ethane, Ethene and Ethyne (mark the hybrid orbitals)
• Which hybridization has the central atom in: H2O, O2, NH3, NH4+, N in pyridine, O in THF, S in SOCl2, C in HCHO compared to CO
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Chemical Reactivity
From the hybrid orbitals we can estimate if a molecule acts as Lewis acid or base (if there is an electrophilic or nucleophilic center) Consider the “empty” pz orbital of C in HCHO vs. the “filled” sp orbital of C in CO -> in the first case, it acts as Lewis acid, in the second as base !
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Acid or Base ?
Compare AlCl3 and PCl3 ? Which acts as acid and which as base – and why ? Why is FeCl3 a strong Lewis acid ?
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***** Break *****
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VSEPR
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VSEPR Theory Intro: http://www.youtube.com/watch?v=nxebQZUVvTg
Practise: http://www.youtube.com/watch?v=xwgid9YuH58
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Rules to remember
Order of repulsions:
(1) Lone Pair – Lone Pair
(2) Lone Pair – Bond
(3) Bond - Bond
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Examples
XeF2
The lone pair needs most space, so they are in equatorial position (bond angle 120 deg) ClF3 lone pairs again in eq. position so have max. distance
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Estimate the structures of:
SF4
BrF5
IF5 2-
AsF5
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cis and trans
mer and fac
eq and ax
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MO Theory
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The Central Themes of MO
Theory
A molecule is viewed on a quantum mechanical level as a collection of nuclei
surrounded by delocalized molecular orbitals.
Atomic wave functions are summed to obtain molecular wave functions.
If wave functions reinforce each other, a bonding MO is formed (region of
high electron density exists between the nuclei).
If wave functions cancel each other, an antibonding MO is formed (a node of
zero electron density occurs between the nuclei).
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Amplitudes of wave
functions added
Figure 11.14
An analogy between light waves and atomic wave functions.
Amplitudes of
wave functions
subtracted.
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Prentice Hall © 2003 Chapter 9
Molecular Orbitals
• Molecular orbitals:
• each contain a maximum of two electrons
• have definite energies
• can be visualized with contour diagrams
• are distributed over the whole molecule (not only in between 2 atoms)
• When two AOs overlap, two MOs form.
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Prentice Hall © 2003 Chapter 9
Molecular Orbitals
The Hydrogen Molecule
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Prentice Hall © 2003 Chapter 9
Figure 11.15 The MO diagram for H2.
Energ
y
MO
of H2
*1s
1s
AO
of H
1s
AO
of H
1s
H2 bond order
= 1/2(2-0) = 1
Filling molecular orbitals with electrons follows the
same concept as filling atomic orbitals.
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Prentice Hall © 2003 Chapter 9
Electron Configurations and Molecular
Properties
• Two types of magnetic behavior:
• paramagnetism (unpaired electrons in molecule): strong attraction between magnetic field and molecule;
• diamagnetism (no unpaired electrons in molecule): weak repulsion between magnetic field and molecule.
• Magnetic behavior is detected by determining the mass of a sample in the presence and absence of magnetic field:
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Diatomic molecules
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Please try to draw the AO’s
and MO’s
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Naming of MO’s: example O2 molecule
“g” = symmetric to C axis “u” = anti-symmetric
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Diatomic molecules Consider the EN of each atom – the higher the EN, the lower is the energy of the orbitals ! The highest filled MO is called “HOMO”, the lowest unoccupied MO “LUMO” -> check example CO
http://firstyear.chem.usyd.edu.au/calculators/ mo_diagrams.shtml
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Example CO
HOMO
LUMO
“lone pair” on C
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Chemical Reactivity Important are the HOMO and LUMO (“frontier orbitals”)
http://www.meta-synthesis.com/webbook/12_lab/lab.html
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Gro
up
Orb
itals
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Construction of Group Orbitals – example H2O
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Interaction 1: in-phase H orbitals
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Interaction 2: out-of-phase H orbitals
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Indicate different MO types: (bonding, non-bonding. anti-bonding)
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Combination of 3 H orbitals to 3 group orbitals
BH3 molecule
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Compare HOMO/LUMO to BH3 ! => what is an acid / base ?
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Homework (3)
Number Molecule
1 CN
2 CN(-)
3 BC
4 BN
5 BO
6 BF
7 CF
8 NO
9 NO (+)
10 NO (-)
Number Molecule
11 NF
12 OF
13 CH4
14 BH3
15 SbF6
16 XeF2
17 XeF4
18 XeF6
http://firstyear.chem.usyd.edu.au/calculators/mo_diagrams.shtml