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AnnouncementsChem 7 Final Exam Wednesday, Oct 10 1:30-3:30AMChapter 1-1270 or 75 multiple choice questions
Exam III (Chapter 7-10)Wednesday, October 3, 2012Time: 6:00PM - 7:30PMSEC A 214A and 215A
Please note the following topics will be excluded from assessment. The page numbers refer to the 2nd Edition of the textbook.
1. Numerical problems involving the Rydberg equation (Chapter 7)2. Spectral analysis in the laboratory (Chapter 7 p. 226-227)3. Trends among the transition elements (Chapter 8 p. 261)4. Trends in electron affinity (Chapter 8 pp. 265-266)5. Pseudo-noble gas configuration (Chapter 8 p. 269)6. Lattice energy (Chapter 9 pp. 283-285)7. IR spectroscopy (Chapter 9 p. 292)8. Numerical problems involving electronegativity (Chapter 9 p. 296)9. Electronegativity and oxidation number (Chapter 9 p. 297)10. Section 11.3: MO theory and electron delocalization (Chapter 11 p.343)11. All sections in chapter 12 except 12.3 (types of intermolecular forces).
Chapter 10The Shapes of Molecules
Chemical bonds and the chemistry of an element is related to the number of valence electrons are in the outer shell (highest value of n quantum number) of the atom.
1A 1ns1
2A 2ns2
3A 3ns2np1
4A 4ns2np2
5A 5ns2np3
6A 6ns2np4
7A 7ns2np5
Group # of valence e-e- configuration
LOOK FOR Group #
• For B group elements, the valence electrons are in the highest value ns orbital and the (n-1)d orbitals.
Lewis dot structures are used to depict valence electrons and bonding between atoms.
• A chemical symbol represents the nucleus and all core e-.
• A single dot around the symbol represents one valence e-.
+1 +2 +3 -1-2-3
1
GOAL: Draw and connect Lewis structures with geometry of a molecule (VSEPRT) and in Chapter 11 connect the geometry with how orbitals bond (VBT).
Molecular formula
Lewis structure
VSEPRTGeometry
Hybrid orbitals
1 2 3 4
2 3 4
VB Theory
Molecular formula
Lewis structure
VSEPRTGeometry
Hybrid orbitals
OctahedralLinearTrigonalPyramidal Tetrahedral Trigonal
Bipyramidal
sp3d2sp sp2 sp3 sp3dValance Bond Theory
Valence Bond theory explains how bonding occurs between atoms using “hybridized orbitals”.
Must Learn How To Draw Lewis Structures1) Put least electronegative element as the central atom. C,S,P and N are often central atoms. H and halogens are often bonded to central atoms.
2) Sum all valence e- from each atom in the molecule (careful with ions add or subtract e-). Use the Group numbers!
3) Place bonds to central atoms using 2 e- per bond.
4) Place an octet of electrons (octet rule) around bonded atoms remembering that H--bonds have no octet. Know the incomplete and expanded octet exceptions.
5) Place remaining electrons around central atom which should have an octet if period 2 or less, but could be more than octet if period 3 or higher.
6) Helpful Rules of Thumb: H forms 1-bond, C forms 4-bonds, N forms 3-bonds, O forms 2-bonds.
Electronegativity is an element’s inherent ability to draw electrons to itself when chemically bonded to another atom in a molecule (relative to Li).
F, O, N, Cl, Br, C are highly electronegative with F being the most electronegative
Write the Lewis structure of nitrogen trifluoride (NF3).
Step 1 – N is less electronegative than F --> N is central atom!
Step 3 - Write Lewis structrue with N central and three bonds and rest non-bonding octet electrons around the central atom.
Step 2 - Count valence electrons
F N F
F
octet
octet
octet
octet
N 5e-
F 7e- x 3 = 21e!Total 26e!
An electron group (domain) is either a pair of bonding electrons or a pair of non-bonding electrons surrounding a central atom. Multiple bonds only count as 1-group or domain.We count and “code” the bonding/non-bonding information into shorthand called AXnEm classification.
AX2E0 = AX2A = Central Atom
X = # of BondedDomains
E = # Non-Bonded Domains
shorthand
F N F
F
4 electron groups3 bonding1 non-bonding AX3E1
1) Incomplete Octet - Occurs with Be, B and Al as central atoms.
2) Expanded Octet (the largest class of octet exceptions)-occurs mostly with Period 3 non-metals like P, S and halogens.
3) Odd-number electrons highly reactive species called radicals that have an odd number of electrons (uneven).
There are three major exceptions to the octet rule.
BF3
BeH2
AlCl3
1. Incomplete Octet - no “octet” around central atom. Occurs with Be, B and Al as central atoms.
Draw Lewis structures for the following
Incomplete Octet: Occurs With Group 2A (Be) and 3A (Boron and Aluminum)
BF3B – 3e-
3F – 3x7e-
24e-
F B F
F
Be – 2e-
2H – 2x1e-
4e-
BeH2 H HBe
Cl Al Cl
Cl
AlCl3Al – 3e-
3Cl – 3x7e-
24e-
Draw Lewis structures for the following
AX3
AX3
AX2
SF6Phosphorous trichloride
[ICl4]-1PCl5
2. Expanded Octet (the largest class of octet exceptions)-occurs mostly with Period 3 non-metals like P, S and halogens.
Draw Lewis structures for the following
SF6
S – 6e-
6F – 42e-
48e- S
F
F
F
FF
F
••
•• ••
••
P
Cl
ClCl
••••
••••
••
••
Phosphorous trichloride PCl3
P
Cl
Cl
••••Cl
••••
••
••••
•• •• ••
••Cl
••••
••Cl••
Phosphorous pentachloride
[ICl4]-1P – 5e-
5Cl – 35e-
40e-PCl5
Expanded Octet (the largest class of octet exceptions)-occurs mostly with Period 3 non-metals like P, S and halogens.
Draw Lewis structures for PCl3 PCl5 and the carbonate anion. Determine the AXE designation for each.
Draw Lewis structues and determine the AXE designation for each.
P
Cl
Cl
••
••Cl
••••
••
••••
•• •• ••
••Cl
••••
••Cl••
5 electron groups5 bonding 0 non-bonding
O C O
O ][ 2- 3 electron groups3 bonding0 non-bonding
••
•• ••
••
P
Cl
ClCl
••••
••••
••
••
4 electron groups3 bonding1 non-bonding
AX3E1
AX5E0
AX3E0
]BrO3[–
Valence e- = 7 + 3(6) + 1 = 26
O Br O
O
MUST look to see if its an ion and add the necessary electron!
H C N
HCNValence e- = 1 + 4 + 5 = 10
Carbon is central atom, watch for hydrogen--1 bond
–
Write the Lewis structure of the carbonate ion (BrO3-)
and hydrogen cyanide, give AXE designation.
4 electron groups3 bonding1 non-bonding
AX3E1
Electron Domains
2 electron groups2 bonding0 non-bonding AX2
Write resonance structures for the carbonate ion, CO3
-.Write the Lewis structure of the carbonate ion (CO3
2-).Step 1 – C is less electronegative than O, put C in centerStep 2 – Count valence electrons (C and O)
Valence electrons = 4 + 6 + 6 + 6 + 2 = 24 valence electronsStep 3 - Arrange the atoms draw bonds between C and O atoms and complete octet on C and O atoms.
3 electron groups3 bonding0 non-bonding
AX3
Electron Domains “resonance structures”All equally good and plausible
Which structure is correct?
A concept called “resonance” is used when more than one plausible Lewis structure can be drawn.
Measured bond lengths show they are equal!
O O O O O O••••••
••
••••
••••
••••
••
••2 equally good Lewisstructures
Which structure is correct?
Example: Ozone, O3
O O O O O O••••••
••
••••
••••
••••
••
••
Both are!
O O O••••••••
••a resonance hybrid structure
Write resonance structures for the nitrate ion, NO3-.
Write resonance structures for the nitrate ion, NO3-.
PLAN:
Count valence e- of atoms = 5 + (3X6) + 1 = 24 e-Most electronegative atom in centerSurround and get an octet around N
3 electron groups3 bonding0 non-bonding
AX3
Electron Domains
O C O O C O
Both are two plausible structures for CO2
When more than one Lewis structure is plausible, we apply the concept of FORMAL CHARGE to figure out the best Lewis structure!
VS
Which is the best one?
USE FORMAL CHARGE
1. The best structure is one that minimizes total formal charge. Net charge of ion or molecule must equal total formal charge.
2.! Also, the best Lewis structure places negative charge on the most electronegative atom.
Assigned Atoms = all from lone pair e! + ! ( bonded e! )
# Valence e- 6 4 6 6 4 6 # of Assinged e- 6 4 6 5 4 7 Formal Charge 0 0 0 +1 0 !1
O C O O C O
This structure wins!
To use the concept of formal charge, we determine the formal charge for each atom.
Atom Formal charge = # valence e- - Assigned e- to Atom
Example: Write 3 plausible Lewis structures for the thiocyanate ion [SCN]–
Example: Write 3 plausible Lewis structures for the thiocyanate ion [SCN]–
S C N[ ] –S C N[ ] –
S C N[ ] –
1 2 3
3-plausible Lewis structures which one is best?
# of Valence = 6 e- + 4 e- + 5 e- + 1 e- = 16 e-S C N
it’s a -1 ion1. Count the Valence e-
2. Draw the Lewis Structures With Least Electronegative Atom as central atom.
N ] –-2S C N[ ] –
FCS = 6 - 4 -2 = 0FCC = 4 - 0 - 4 = 0FCN = 5 - 6 - 2 = -1
S C[FCS = 6 - 2 -3 = 1FCC = 4 - 0 - 4 = 0FCN = 5 - 6 - 1 = -2
S C N[ ] –
FCS = 6 - 6 -1 = -1FCC = 4 - 0 - 4 = 0FCN = 5 - 2 - 3 = 0
0 0 -1 -1 0 0 0+1
1. Formal charge must sum to charge of ion or molecule.2. N is more electronegative than C or S, it should have a the most negative charge in the “best structure”.3. The most plausible structure has the least amount of formal charge.
Structure on the left is “best” structure!S C N[ ] –0 0 -1
Example: Write 3 plausible Lewis structures for the thiocyanate ion [SCN]–
Write resonance structures for the nitrate ion, NCO-
and determine the most plausible Lewis structure.EXAMPLE: NCO- has 3 possible resonance forms -
A B C
formal charges
-2 0 +1 -1 0 0 0 0 -1
Forms B and C have negative formal charges on N and O; this makes them more preferred than form A.Form C has a negative charge on O which is the more electronegative element, therefore C contributes the most to the resonance hybrid.
Write resonance structures for the nitrate ion, NCO-
and determine the most plausible Lewis structure.
Chemists use Valence Shell Electron Pair Repulsion Theory to predict the shapes of molecules using these five electron group geometries.
1. Draw Lewis Structure from chemical formula.
2. Count all electron domains to get AXE code.
4. Match the number of bonding and non-bonding domains to the proper VSEPRT geometry.
3. Group domains into bonding and non-bonding pairs of electrons.
VSEPRT explains the geometry of molecules but NOT how covalent bonds are formed with that geometry.
Molecular formula
Lewis structure
VSEPRTGeometry
Hybrid orbitals
VSEPRT Valence BondTheoryVSEPRTLewis Structure
The electron geometry is the geometry of all electron groups, whereas the “molecular geometry” describes the geometry of only the atoms bonded to the central atom.
AX3E1 = Tetrahedral electron geometery with 109.5˚ bond angles.
Molecular Geometry is trigonal pyramidal bond angles <109.5˚
Molecular Geometry
Electron GroupGeometry
Molecular formula
Lewis structure
VSEPRTGeometry
Hybrid orbitals
The goal is to understand geometry (via VSEPRT) and to relate it to a picture of covalent bonding in molecules.
VB Theory
OctahedralLinearTrigonalPyramidal Tetrahedral Trigonal
Bipyramidal
sp3d2sp sp2 sp3 sp3d
The 3-D geometry of a molecule is one of five basic arrangements of electron groups (domains).
Trigonal Planar
Trigonal Bipyramidal Octahedral
Linear Tetrahedral
The total number of electron groups (domains) defines one of the five basic geometries.
2 EG
3 EG 4 EG
5 EG 6 EG
The electron geometry is the geometry of all electron domains, whereas the “molecular geometry” describes the geometry of only the atoms bonded to the central atom.
AX3E1 = Tetrahedral electron geometery with 109.5˚ bond angles.
Molecular Geometry is trigonal pyramidal bond angles <109.5˚
How Predict Geometry Using VSEPRT1.Draw a plausible Lewis structure for the molecule.
2.Determine the total number of electron domains and identify them as bonding or lone pairs.
3.Use the total number of electron domains to establish the electron geometry from one of the five possible geometric shapes.
4.Establish the AXnEm designation to establish the molecular geometry (or do both electron and molecular geometry together simultaneously)
5. Remember bond angles in molecules are altered by lone pairs of electrons (repulsion forces reduce angles).
6. Molecules with more than one central atom can be handled individually.
2 Electron Groups = Linear Electron Geometry and 1-Possible Molecular Geometry
Other Examples:CS2, HCN, BeF2
Bond Angle
AX2E0 = AX2
Cl ClBe
S C N
O C O
A = Central AtomX = # of BondedDomains
E = # Non-Bonded Domains
Examples:SO2, O3, PbCl2, SnBr2 A
Examples:SO3, BF3, NO3
-, CO32-
AX3A
3-Electron Domain
3 Electron Groups = Trigonal Planar Electron Geometry and 2-Possible Molecular Geometries
AX2E1
Examples:
CH4, SiCl4, SO4
2-, ClO4-
NH3
PF3
ClO3
H3O+
H2O
OF2
SCl2
AX4
AX3E1 AX2E2
4 Electron Groups = Tetrahedral Electron Geometry and 3-Possible Molecular Geometries
Bond Angle
SF4
XeO2F2
IF4+
IO2F2-
ClF3
BrF3
XeF2
I3-
IF2-
PF5
AsF5
SOF4
AX5 AX4E1
AX2E3
AX3E2
AxialPosition
EquatorialPosition
5 Electron Groups = Trigonal Bipyramial Electron Geometry and 4-Possible Molecular Geometries
SF6
IOF5
BrF5
TeF5-
XeOF4
XeF4
ICl4-AX4E2
AX6
AX5E1
6 Electron Groups = Octahedral Electron Geometry and 3-Possible Molecular Geometries
Predicting Molecular ShapesDraw the molecular shape and predict the bond angles (relative to the ideal bond angles) of (a) PF3 and (b) COCl2.
<109.50
Predicting Molecular ShapesDraw the molecular shape and predict the bond angles (relative to the ideal bond angles) of (a) PF3 and (b) COCl2.
2. Count the electron domains and find electron geometry and molecular from core 5 electron domain shapes (using AXE designation and sub-shapes)
5. The F-P-F bond angles should be <109.50 due to the repulsion of the nonbonding electron pair.
3. There are 4 electron domains so the electron geometry is tetrahedral
4. The designation is AX3E1 so the molecular geometry is trigonal pyramidal.
1. Count the valence electrons and draw Lewis structure for PF3: VE = 5 + 3(7) = 26 e-
Predicting Molecular Shapes with Two, Three, or Four Electron Groups
(b) For COCl2, C has the lowest EN and will be the center atom.There are 24 valence e-, 3 atoms attached to the center atom.
124.50
1110
Type AX3
5. The Cl-C-Cl bond angle will be less than 1200 due to the electron density of the C=O.
2. Count the electron domains and establish electron geometry from 5 shapes3. There are 3 electron domains so the electron geometry is trigonal planar4. The molecular geometry designation is AX3E0 so the molecular geometry is also trigonal planar (no lone pairs).
1. Draw the Lewis structure
Determine the molecular shape and predict the bond angles (relative to the ideal bond angles) of (a) SbF5 and (b) BrF5.
Determine the molecular shape and predict the bond angles (relative to the ideal bond angles) of (a) SbF5 and (b) BrF5.
(a) SbF5 - 40 valence e-; all electrons around central atom will be in bonding pairs; shape is AX5 - trigonal bipyramidal.
(b) BrF5 - 42 valence e-; 5 bonding pairs and 1 nonbonding pair on central atom. Shape is AX5E, square pyramidal.
More Than One Central Atom
• In acetic acid, CH3COOH, there are three central atoms.• We assign the geometry about each central atom
separately.
What is the geometryaround these atoms?
Take one atom at a time and apply the rules of electron domains.
ethaneCH3CH3
ethanolCH3CH2OH
More Than One Central Atom
Determine the shape around each of the central atoms in acetone, (CH3)2C=O.
Find the shape of one atom at a time after writing the Lewis structure.
tetrahedral tetrahedral
trigonal planar
>1200
<1200
Predicting the Molecular Shape With Multiple Central Atoms Electronegativity is an element’s inherent property to draw electrons to itself when chemically bonded to another atom in a molecule. The units are dimensionless (all relative measurements to Li).
RankFONClBr
Differences in elements electronegativity between bonding atoms result in the formation of polar-covalent bonds and net dipole moments in molecules.
Net Dipole MomentNo Net Dipole Moment
Polar BondPolar Bond
Polar Bond Polar Bond
Think of the dipole moment as a molecule with separated charges + and -.
Draw a Lewis structure, show the AXE designation, determine electron and molecular geometry and whether polar or non-polar of:
CCl3H
CCl4
CH4
Draw a Lewis structure, show the AXE designation, determine electron and molecular geometry and whether polar or non-polar of:
CCl4AX4Tetrahedral EGTetrahedra MG
CH3ClPolar bondPolar MoleculeHas Dipole Moment
CCl4Polar bondsNot Polar MoleculeNo Dipole Moment
CH3Cl
For a poly-atomic molecule we must consider the vector sum of polar bonds in the molecule to see if there is a net dipole moment.
No NetDipoleMoment
DipoleMoment
DipoleMoment
DipoleMoment
No NetDipoleMoment
From electronegativity (EN) values (button) and their periodic trends, predict whether each of the following molecules is polar and show the direction of bond dipoles and the overall molecular dipole when applicable:
(a) Ammonia, NH3
(b) Boron trifluoride, BF3
(c) Carbonyl sulfide, COS (atom sequence SCO)
10-
Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Predicting the Polarity of Molecules
(a) Ammonia, NH3 (b) Boron trifluoride, BF3
(c) Carbonyl sulfide, COS (atom sequence SCO)
PROBLEM: From electronegativity (EN) values (button) and their periodic trends, predict whether each of the following molecules is polar and show the direction of bond dipoles and the overall molecular dipole when applicable:
PLAN: Draw the shape, find the EN values and combine the concepts to determine the polarity.
SOLUTION: (a) NH3
ENN = 3.0
ENH = 2.1
bond dipoles molecular dipole
The dipoles reinforce each other, so the overall molecule is definitely polar.
10-
Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Sample Problem 10.9 Predicting the Polarity of Molecules
(b) BF3 has 24 valence e! and all electrons around the B will be involved in bonds. The shape is AX3, trigonal planar.
F (EN 4.0) is more electronegative than B (EN 2.0) and all of the dipoles will be directed from B to F. Because all are at the same angle and of the same magnitude, the molecule is nonpolar.
1200
(c) COS is linear. C and S have the same EN (2.0) but the C=O bond is quite polar ("EN), so the molecule is polar overall.
11-
Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Chapter 11
Theories of Covalent Bonding
11-
Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Theories of Covalent Bonding
11.1 Valence Bond (VB) Theory and Orbital Hybridization
11.2 The Mode of Orbital Overlap and the Types of Covalent Bonds
11.3 Molecular Orbital (MO)Theory and Electron Delocalization
This only!
Valence Bond Theory explains covalent bonding by the spatial overlap of atomic orbitals on bonding atoms and the sharing of electron pairs.
Electrons that must have opposite spins.
1s1 + 1s1
Bonding in F2
1s1 + 2p5
Bonding in H2
Bonding in HF
2p5 + 2p5
It’s natural to think of using “pure atomic orbitals” to describe bonding in some molecules. It works well for some....but fails with carbon.
Bonding in F2
1s1 + 1s22s22p5
Bonding in HF
2p5 + 2p5
1s22s22p5 + 1s22s22p5
1s1 + 2p5
sp3
hybridized orbitals
hybridization fixes the problem!
Bonding in carbon presents a problem as combining atomics orbitals fails. Valance Bond Theory solves this by allowing the blending or mixing of pure atomic orbitals in a process called hybridization.
Pure atomic orbitals (valence orbitals)
only two bond are possible with use of atomic orbitals only....we don’t observe CH2
By hybridizing 4 bonds are possible.
Hybridization combines or mixes different numbers of pure atomic orbitals that match one of the VSEPRT geometries. For example 1 pure s orbital + 1 p-orbital combine to give and 2 “sp hybrids” that when superimposed form a linear geometry for bonding.
s-orbital + p-orbital --> 2 sp hybrid orbitals -->
s-orbital + Two p-orbital --> 3 sp2 hybrids = Trig Planar
2-superimposed sphybrid orbitals
s-orbital + Three p-orbitals -> Four sp3 hybrids = Tetrahedral
sp3 hybrid orbitals
The process of combining pure atomic orbitals to form “hybrid orbitals” on central bonding atoms in a molecule is called hybridization.
1. The number of hybrid orbitals obtained equals the number of atomic orbitals mixed.
2. The name of and shape of a “hybrid orbital” varies with the types of atomic orbitals mixed. (s + p vs s + two p)3. Each hybrid orbital has a specific geometry that matches one of five VSEPRT shapes (show below).
sp3d2
Octahedral
sp sp2 sp3 sp3d
Linear TrigonalPlanar
Tetrahedral Trigonal Bipyramidal
Some generalized rules and comments on VBT and the formation of hybridized orbitals.
sp3d2
Octahedral
sp sp2 sp3 sp3d
Linear Trigonalplanar
Tetrahedral Trigonal Bipyramidal
Molecular formula
Lewis structure
VSEPRTGeometry
Hybrid orbitals
Connect the dots and it becomes easy to see and understand.
Valence Bond Theoryexplains how bonds are made
ElectronGeometry
Molecular Geometry AXnEm Hybridization
Linear Linear AX2 sp
Trigonal planar
Trigonal planar V-shaped bent
AX3
AX2E1sp2
Tetrahedral
Tetrahedral Trigonal pyramidal
V-shaped bent
AX4
AX3E1 AX2E2
sp3
Trigonal bipyramidal
Trigonal bipyramidalSeesaw
T-shaped Linear
AX5
AX4E1
AX3E2
AX2E3
sp3d
OctahedralOctahedral
Square pyramidal Square planar
AX6
AX5E1
AX4E2
sp3d2
Determine the VSEPRT geometry, the bond angles and the hybridization of each indicated atom in the following molecule? How many sigma and pi bonds are in the molecule?
Determine the electron domain, molecular geometry, the bond angles and the hybridization of each indicated atom in the following molecule? How many sigma and pi bonds are in the molecule?
tetrahedral, 180, sp3
sp3
sp2
sp
sp2
bent, <109.5, sp3
trig planar 120˚, sp2
linear 180˚, sp
Atomic OrbitalsMixed
# Hybrid OrbitalsFormed
HybridShape
Linear AX2
Trig Planar AX3
Tetrahedral AX4
Trig Bypyr AX5
Octahedral AX6
s + p s + 2 p s + 3 p s + 3 p + d s + 3 p + 2d
Two sp Three sp2 Four sp3 Five sp3d Six sp3d2
Orbitals Leftover for Pi bonds
Two p one p none Four d Three d
Linking VSEPRT To Valence Bond Theory Hybrids
2s
--The number of hybrid orbitals formed is equal to the number of “pure orbitals” combined!
--When superimposed the “sp-hybrid” give us bonding orbitals for a linear molecules.
An sp hybrid is formed from the combination of a one pure 1s orbital and a one 2p orbital from a central bonding atom producing two new orbitals called sp orbitals.
s-orbital p-orbitalTwo sp hybrid orbitals
sp hybrid orbitals superimposed
Hybridization
s + p Hybridization = 2 sp
Example sp hybrid: Show the bonding scheme and hybridized orbitals used in BeCl2
2 unhybridized unoccupied p-orbitals
After hybridization we have on the central atom, 2 pure p-orbitals and two sp hybrids.
2 “left-over” p-orbitals
hybridization
Isolated Be Atom
Hybridized Be Atom
Show the bonding scheme and hybridized orbitals in BeCl2
two sp hybrids on Be
two lone p-orbitals
sp2 = Triginal planar geometry, 120˚ bond angle
3-atomic orbitals, s and two p’s combine to form 3-sp2 hybrid orbitals
An sp2 hybrid is formed from the combination of a one pure 1s orbital and a two 2p orbitals from a central bonding atom producing two new orbitals called sp2 orbitals.
Superimposed Hybrid orbitals form a triginal planar geometry
Example 2: sp2 hybridizaton scheme BF3.
Boron Orbital Box Diagram
Boron Hybrid Box Diagram
Bonding of pure p-orbital in F with sp2 hybridized orbitals in BF3
Tetrahedral geometry = sp3 hybrid orbitals
sp3 = Tetrahedral geometry = 109.5˚ bond angle
Note the number of hybrids formed is the number of atomic orbitals combined!
combine to generatefour sp3 orbitals
which are representedcollectively as: sp3
Example: sp3 orbital hybridization: CH4.
the four sp3 hybrid orbitals form a tetrahedral shape
sp3 hybridization mixes one 2s orbital with three 2p orbitals to produce four sp3 orbitals on each carbon atom. End to end overlap with a 1s orbital from H gives four sigma bond in CH4.
CH4
This is the ground stateconfiguration of valence atomic orbitals
Example 3: sp3 hybrid orbitals in H2O.
What is the electronic geometry?What is the molecular geometry?What orbitals contribute to bonding?
Note the lone pairs occupy 2-of the sp3 orbitals
sp3 hybridization mixes one 2s orbital with three 2p orbitals to produce four sp3 orbitals. The e- are distributed throughout the hybrids ready for bonding. End to end overlap with a 1s orbital from H gives four sigma bond in CH4.
sp3 is tetrahedral shape. In water we have AX2E2
What is the electron geometry, the molecular geometry at each carbon atom? Use that information to determine the hybridization around each carbon atom in nicotinic acid? How many sigma and pi bonds are in nicotinic acid?
Example 2: sp3 hybridization in NH3.
Tetrahedral Electron Geometry AX3E1Trigonal Pyramidal Molecular Geometry
sp3d hybridization in PCl5.
Isolated P atom
Trigonal Bipyramidal Electron Geometry AX5E0Trigonal BiPyramidal Molecular Geometry
The sp3d2 hybrid orbitals in SF6Octahedral Electron Geometry AX6E0Octahedral Molecular Geometry
ElectronGeometry
Molecular Geometry AXnEm Hybridizaton
Linear Linear AX2 spTrigonal planar
Trigonal planar V-shaped bent
AX3
AX2E1sp2
Tetrahedral
Tetrahedral Trigonal pyramidal
V-shaped bent
AX4
AX3E1 AX2E2
sp3
Trigonal bipyramidal
Trigonal bipyramidalSeesaw
T-shaped Linear
AX5
AX4E1
AX3E2
AX2E3
sp3d
OctahedralOctahedral
Square pyramidal Square planar
AX6
AX5E1
AX4E2
sp3d2
Describe the types of bonds and orbitals in acetone, (CH3)2CO and in CO2 and in HCN?
Molecular formula
Lewis structure
VSEPRTGeometry
Hybrid orbitals
Step 1 Step 2 Step 3
Describe the types of bonds and orbitals in acetone, (CH3)2CO.PLAN:
Draw the Lewis structures to ascertain the arrangement of groups and shape at each central atom. Postulate the hybrid orbitals taking note of geometries predicted from VSEPRT. Draw the orbitals and show overlap.
SOLUTION:
sp3 hybridized
sp3 hybridized
sp2 hybridized
# bonds$ bond
sp hybrid:Ethylyne: HC!CH:Linear
sp hybrid orbitals
Lone p orbitals that are not hybridized
Sigma bonds (" bonds) and Pi bonds (# bonds)are two different types of covalent chemical bonds that form as a result of end to end spatial overlap of atomic orbitals or hybridized orbitals (" bonds) or side to side overlap on bonding atoms (# bonds)
Lone p orbitals that were not hybridized on each carbon atom are able to form Pi bonds in a “side to side” overlap. A pair of electrons is shared in this region of space.
sp2 hybrid orbitals on each carbon atom use end to end overlap to form a sigma bond.
# bonds overlap side to side
sp hybrid:Ethylyne: HC!CH:Linear