Unit 7

1
Unit 7 Bonding Molecular Geometry
2
Definitions

• Chemical Bonds
• Force that holds atoms together
• Its all about the electrons (e-)
• Electrons available for bonding are called
  valence electrons!

3
Types of Chemical Bonds

• Ionic Bond
• Bond between metal and nonmetal due to
  electrostatic interactions
• Attraction between positively and negatively
  charged ions (cations and anions)
• Electrons are completely transferred from metal
  to nonmetal

4
Ionic bonds Result from a Transfer of Valence
Electrons

-
5
Types of Chemical Bonds

• Covalent Bond
• Bonds in which e- are shared
• Most common type

6
Shared Electrons Complete Shells
F
F
7
Hydrogen MoleculeEnergy Diagram
Note at .074 nm, attractive forces are balanced
with repulsive forces!
8
Definitions

• Octet rule (Rule of 8)
• Most atoms want 8 e- in the outer shell? very
  stable
• H2 and He want a duet (2 e-)
• Electron configuration for duet ns2
• Electron configuration for octet ns2 np6

9
Two other definitions you need to know

• Bonding pairs are electrons involved in bonding.
• Lone pairs are electrons NOT involved in bonding.
• They are only located on one atom. (a.k.a
  non-bonding pair)

10
Lewis Dot Diagrams

• A Lewis dot diagram depicts an atom as its symbol
  and its valence electrons.
• Ex Carbon

.
.
.
C
.
Carbon has four electrons in its valence shell
(carbon is in group 14), so we place four dots
representing those four valence electrons around
the symbol for carbon.
11
Drawing Lewis Dot Diagrams

• Electrons are placed one at a time in a clockwise
  manner around the symbol in the north, east,
  south and west positions, only doubling up if
  there are five or more valence electrons.
• Same group Same Lewis Dot structure
• Ex. F, Cl, Br, I, At
• Example Chlorine (7 valence electrons b/c it is
  in group 17)

.
.
.
.
Cl
.
.
.
12
Note In the final structure, the placement of
the dots around the element is not crucial
Maximum of valence electrons 8
13
Paired and Unpaired Electrons

• As we can see from the chlorine example, there
  are six electrons that are paired up and one that
  is unpaired.
• When it comes to bonding, atoms tend to pair up
  unpaired electrons.
• A bond that forms when one atom gives an unpaired
  electron to another atom is called an ionic bond.
• A bond that forms when atoms share unpaired
  electrons between each other is called a covalent
  bond.

14
Bonding in Ionic Compounds

• The ionic bond forms from attraction of cations
  for anions.

15
Review of Ionic Charge and Isoelectronic Ions

• Isoelectronic having same of e-
• (same e- configuration)
• Na ? Na e-
• Cl e- ? Cl1-
• What elements are Na and Cl- isoelectronic with?

16
Structure of Ionic Compounds

• Ionic compounds have
• formula unitsthese show ratio of ions in the
  crystal lattice.

17
Writing Lewis Dots Structures for Ions

• Uses either 0 or 8 dots, brackets and a
  superscript charge designate to ionic charge
• Ex.) Li, Be2, B3, C4, N-3, O-2, F-1

18
Writing Lewis Dots Structures(Ionic Compounds)

• Lewis Dot Diagrams of Ionic Compounds
• Ex. 1) NaCl
• Ex. 2) Li2O

19
Lewis Representations of Ionic Structures

• NaCl

MgO
Li2O
20
Covalent Compounds and Lewis Dot Diagrams

• Lewis structures for covalent molecules show
  sharing of e-
• HH OR H-H
• Bonding pair e- (shared e-) are counted as
  belonging to both atoms. (each atom has octet)
• Bonding pair can also be shown as a dash between
  atoms.

21
Drawing Electron Dot Diagrams for Molecules

• Chemists usually denote a shared pair of
  electrons as a straight line.

F F

• Sometimes the nonbonding pair of electrons are
  left off of the electron dot diagram for a
  molecule

22
Examples
H
CH4
C
H
H
H
H
H
N
NH3
H
23
Types of Covalent Bonds

• Single Bond
• 2 e- are shared in a bond (1 from each atom)
• Double Bond
• 2 pairs of e- are shared (4 e- total, 2 from each
  atom)
• Triple Bond
• 3 pairs of e- are shared (6 e- total, 3 from each
  atom)

24
Rules for Drawing Lewis Dot Diagrams

• Add up the total number of valence e- for each
  atom in the molecule.
• Each (-) sign counts as 1 e-, each () sign
  subtracts one e-
• Write the symbol for the central atom then use
  one pair of e- to form bonds between the central
  atom and the remaining atoms.
• Count the number of e- remaining and distribute
  according to octet rule (or the duet rule for
  hydrogen)
• If there are not enough pairs, make sure the most
  electronegative elements are satisfied. Then,
  start shifting pairs into double and triple bonds
  to satisfy the octet rule.
• If there are extra e-, stick them on the central
  atom.

25
Hints

• H is NEVER a central atom!
• Halogens (Group 17) are usually not central
  atoms.
• If you only have 1 of a certain element, it is
  usually the central atom.

26
Checking Your Work!

• But Remember....
• The Structure MUST Have the right number of
  atoms for each element, the right number of
  electrons, the right overall charge, and 8
  electrons around each atom (ideally).

27
Examples

• F2
• H2O
• OCl-
• PO43-

28
Examples

• O2
• CH4
• HF
• NH3

29
Examples

• NH4
• SO32-
• N2
• CH3OH

30
Exceptions to the Octet Rule

• Reduced Octets electron deficient molecules
• (Be and B)
• Be 2 valence e-, doesnt form
• octet (BeH2 Be has 4 e-)
• B 3 valence e-, doesnt form octet
• (BF3 B has 6 e-)

31
Exceptions to the Octet Rule

• Expanded Octets
• (Examples P, S, Cl, As, Se, Br, Kr, Xe)
• How to recognize
• The central atom in PERIOD 3 or greater is
  surrounded by gt 4 atoms.
• You draw the Lewis diagram and the results dont
  make sense the central atom has gt 8 e-

32
Expanded Octets (P, S, Cl, As, Se, Br, Kr, Xe)

• Examples
• PF5 XeF4

33
Resonance Structures

• Definition
• When a single Lewis structure does not adequately
  represent a substance, the true structure is
  intermediate between two or more structures which
  are called resonance structures.
• Resonance Structures are created by moving
  electrons, NOT atoms.

34
Resonance Structure Example, SO2

• Central atom S
• This leads to the following structures
• These equivalent structures are called
• RESONANCE STRUCTURES. The true structure is a
  HYBRID of the two.
• Arrow means in resonance with

35
Resonance Structure Example, NO3-

• Draw the Lewis diagram for NO3- with all possible
  resonance structures.

36
Radicals

• When there is an odd of total electrons, there
  will be a single, unpaired electron in the
  structure!
• Example NO
• Radicals are extremely reactive they want to
  have paired electrons!!

37
Linus Pauling, 1901-1994

• The only person to receive two unshared Nobel
  prizes (for Peace and Chemistry).
• Chemistry areas bonding, electronegativity,
  protein structure

38
Electronegativity

• Definition
• A measure of the ability of an atom in a
  molecule or bond to attract electrons to itself.
• Scale proposed by Linus Pauling
• Greater E.N. means element more strongly
  attracts electrons.

39
Electronegativity

• Trends on periodic table
• Highest on upper right
• (F has highest with e/n 4.0)
• Lowest on lower left (Francium 0.7)
• Noble gases have ZERO E.N.

40
Electronegativity
41
Bond Polarity

• Polar Covalent Bond
• Covalent bond in which the electrons are
  unequally shared
• Ex. H2O
• Non-polar Covalent Bond
• Covalent bond in which the electrons are equally
  shared
• Ex. F2 or CH4
• Predicting Bond Polarity
• Use Electronegativity!! (see next slide)

42
Predicting Bond Polarity

• Calculate the difference between the Pauling
  electronegativity values for the 2 elements

Type of Bond IONIC (COVALENT) (COVALENT)
Type of Bond IONIC POLAR NON-POLAR
Types of Atoms 1 metal 1 nonmetal (ex. NaCl) (generally) 2 nonmetals Ex. NH3, H2O (generally) 2 nonmetals Ex. CCl4, O2
Electronegativity Difference 1.7 0.4 but lt 1.7 0.4
0 0.4 ? Non-polar covalent 0.4 1.7 ? Polar
covalent (more e/n element has greater pull) 1.7
and up ? Ionic (e- are transferred between atoms)
43
Using e/n to predict polarity of individual bonds

• A polar bond has a partial charge due to unequal
  sharing of electrons.
• A polar bond is shown using partial charges
  either with delta or cross/arrow.

Negative delta or arrow next to more E.N. atom.
44
Bond Polarity

• HCl is POLAR because it has a positive end and a
  negative end.

Cl has a greater share of bonding electrons than
H.
Cl has slight negative charge (d-) and H has
slight positive charge (d)
45
Bond Polarity

• What type of bonds are these?
• OH OF
• E.N. 3.5 - 2.1 3.5 - 4.0
• ? 1.4 0.5

46
Molecular Geometry

• Molecular Geometry describes the
• 3-D arrangement of atoms in a molecule.
• We will use VSEPR theory to predict these 3-D
  shapes!

47
VSEPR Shapes of Molecules

• VSEPR Theory (definition)
• Valence Shell Electron Pair Repulsion
• Based on idea that e- pairs want to be as far
  apart as possible
• The molecule adopts the shape that minimizes the
  electron pair repulsions.
• Based on molecular shape of Lewis diagram

48
We define the electron pair geometry by the
positions in 3D space of ALL electron pairs
(bonding and non-bonding). The molecular
geometry only considers the positions of the
bonded electrons.
49

• To determine the electron pair geometry
• 1. Draw the Lewis structure.
• 2. Count the number of bonded (X) atoms and
  non-bonded or lone pairs (E) around the central
  atom.
• 3. Based on the total of X E, assign the
  electron pair geometry.
• 4. Multiple bonds count as one bonded pair!

50
Electron-pair geometry around a central atom

• Sum of X E Shapes
• 2 linear
• 3 trigonal
  planar
• 4 tetrahedral
• 5 trigonal
  bipyramidal
• 6 octahedral

51
Molecular geometry around a central atom

• A Central Atom
• X Bonded Atom
• E Non-bonded electron pair (Lone pair e-)
• BPs Bonding Pairs
• LPs Lone Pairs

52
Molecular Geometry AX2
Formula BPs LPs Shape Angles Examples
AX2 2 0 Linear 180 CO2
53
Molecular Geometry AX3
Formula BPs LPs Shape Angles Examples
AX3 3 0 Trigonal Planar 120 BF3
54
Molecular Geometry AX4
Formula BPs LPs Shape Angles Examples
AX4 4 0 Tetra-hedral 109.5o CH4
55
(No Transcript)
56
Molecular Geometry AX5
Formula BPs LPs Shape Angles Examples
AX5 5 0 Trig. Bipyramid 90,120 180 PF5
57
Molecular Geometry AX6
Formula BPs LPs Shape Angles Examples
AX6 6 0 Octa- hedral 90, 180 SF6
58
(No Transcript)
59
Lone Pairs on Central AtomAX2E
Formula BPs LPs Shape Angles Examples
AX2E 2 1 Bent lt 120 NO2-
60
(No Transcript)
61
Lone Pairs on Central AtomAX3E
Formula BPs LPs Shape Angles Examples
AX3E 3 1 Trigonal Pyramid lt 109.5o NH3
62
Lone Pairs on the central atom AX3E
Classification AX3E Bond Angles are lt
109.5o Electron Pair Geometry Tetrahedral Molecu
lar Geometry Trigonal Pyramidal
63
Lone Pairs on Central AtomAX2E2
Formula BPs LPs Shape Angles Examples
AX2E2 2 2 Bent lt 109.5o H2O
64
Unshared Pairs of e- on the central atom
Classification AX2E2 Bond Angles are lt
109.5o Electron Pair Geometry Tetrahedral Molecu
lar Geometry Bent
65
More Complicated Shapes

• These shapes result from expanded octets.
• How to recognize
• The sum of X E gt 4!

66
More Complicated ShapesAX4E
Formula BPs LPs Shape Angles Examples
AX4E 4 1 Distorted Tetra-hedron OR Seesaw lt 90, 120, 180o SeCl4
67
More Complicated ShapesAX3E2
Formula BPs LPs Shape Angles Examples
AX3E2 3 2 T-shaped lt 90, 120, 180o BrCl3
68
More Complicated ShapesAX2E3
Formula BPs LPs Shape Angles Examples
AX2E3 2 3 Linear lt 180o XeF2
69
Molecular Geometries for Five Electron Pairs
AX5 AX4E AX3E2 AX2E3
70
More Complicated ShapesAX5E
Formula BPs LPs Shape Angles Examples
AX5E 5 1 Square Pyramid lt 90, 180o BrF5
71
More Complicated ShapesAX4E2
Formula BPs LPs Shape Angles Examples
AX4E2 4 2 Square Planar lt 90 XeF4
72
Molecular Geometries for Six Electron Pairs
AX6 AX5E AX4E2

73
Molecules with More than One Central
Atom Determine geometry for each central atom
separately! Example In acetic acid, CH3COOH,
there are three central atoms C
C O
74
Molecules with only two atoms are always linear!

• ExamplesHCl N2

75
VSEPR Examples

• What shape would the following compounds have
  according to VSEPR theory?
• O2
• CFCl3

76
VSEPR Examples

• What shape would the following compounds have
  according to VSEPR theory?
• H2S
• PBr5

77
VSEPR Examples

• What shape would the following compounds have
  according to VSEPR theory?
• SeCl22-
• C2H2
• (hint classify each C separately)

78
Electronic Flashcards

• http//www.proprofs.com/flashcards/cards.php?id75
  21
• Flash cards on molecular geometry and
  hybridization

79
Polar Molecules and Dipole Moments

• Polar molecules are NOT the same as polar bonds!
• Cant use ? E.N. to calculate if something is a
  polar molecule!

80
Polar Bond

• Definition When electrons are unequally shared
  between two atoms in a BOND
• Example C-O bond in CO2

81
Polar Molecules

• Polar Molecules (a.k.a. Dipoles)
• Molecule with separate centers of () and (-)
  charge
• In other words, molecules are polar if the pull
  in any one direction is not balanced out by an
  equal opposite pull in the opposite direction

82
Dipole Moment
Polar molecules have a DIPOLE MOMENT will align
with an electric field.
83
Determining Molecular Polarity
Nonpolar Polar
84
Nonpolar Molecules Bond Polarity Cancels if
Structure is Symmetrical
85
Simple Molecules Rules for Determining Polarity

• A molecule is polar if
• It has only 2 atoms in it and both are different.
• It has 3 or more atoms and has lone pairs on the
  central atom (i.e., it is classified as an AXE)
• Exception where the lone pairs are
  symmetrical to the axis of bonded atoms in AX4E2
  or AX2E3.
• 3. It has 3 or more atoms in an AXn
  classification and all of the Xs are not the
  same atom.

86
Example

• Determine which of the following are dipoles
• BF3 BF2Cl CO2
• H2O H3O NCl3
• PF5 XeF4

87
Hybridization of Atomic Orbitals
The solutions of the Schrodinger equation led to
these atomic orbitals 1s, 2s, 2p, 3s, 3p, 3d,
4s, 4p, 4d, 4f, etc. However, overlap of these
orbitals does not give a satisfactory
explanation. In order to explain bonding, these
orbitals are combined to form new sets of
orbitals this method is called hybridization.
88
Hybridization of Atomic Orbitals
sp 2 sp hybrid orbitals from mixing of a s and a
p orbital sp 2 3 sp2 hybrid orbitals from mixing
of a s and 2 p orbitals sp3 4 sp3 hybrid
orbitals from mixing of a s and 3 p orbitals sp3d
5 sp3d hybrid orbitals from mixing of a s and 3
p and a d orbital sp3d 2 6 sp3d2 hybrid
orbitals from mixing of a s and 3 p and 2 d
orbitals
89
Hybridization of Atomic Orbitals
sp 2 sp hybrid orbitals from mixing of a s and a
p orbital sp 2 3 sp2 hybrid orbitals from mixing
of a s and 2 p orbitals sp3 4 sp3 hybrid
orbitals from mixing of a s and 3 p orbitals sp3d
5 sp3d hybrid orbitals from mixing of a s and 3
p and a d orbital sp3d 2 6 sp3d2 hybrid
orbitals from mixing of a s and 3 p and 2 d
orbitals Notice that there are five hybrid
orbital types they match up with the five
electron pair geometries!!!
90
Hybridization of Atomic Orbitals
sp 2 sp hybrid orbitals from mixing of a s and a
p orbital sp 2 3 sp2 hybrid orbitals from mixing
of a s and 2 p orbitals sp3 4 sp3 hybrid
orbitals from mixing of a s and 3 p orbitals sp3d
5 sp3d hybrid orbitals from mixing of a s and 3
p and a d orbital sp3d 2 6 sp3d2 hybrid
orbitals from mixing of a s and 3 p and 2 d
orbitals Superscripts on s, p, d added
together Sum of X E in the designation for
the electron pair geometry (sum of bonded atoms
and lone pair e-)
91
Hybridized Orbitals in bonding, sp
The sp hybrid orbitals formation of two sp
hybrid orbitals - -
- - hybridization of s and p
orbitals 2 sp hybrid orbitals Two sp hybrid
orbitals gt
92
The sp2 Hybrid Orbitals
The hybridization of a s and two p orbitals led
to 3 sp2 hybrid orbitals for bonding.
93
The sp3 Hybridized Orbitals
The hybridization of a s and three p orbitals led
to 4 sp3 hybrid orbitals for bonding. Compounds
involving sp3 hybrid orbitals CF4, CH4, NH3,
H2O, SiO44, SO42, ClO4, etc
94
How to determine

• ALWAYS results from tetrahedral electron pair
  geometry
• ONLY hybridization explains why C forms four
  equal bonds!
• Link for animations
• http//highered.mcgraw-hill.com/sites/0072512644/
  student_view0/chapter10/animations_center.html

95
Sample Problem

• Predict the molecular geometry and hybridization
  of the central atom in the following compounds
• OF2 NH4 CO2
• COCl2 XeF4

96
Counting Sigma and Pi Bonds

• Single bond 1 sigma bond
• Double bond1 sigma and 1 pi bond
• Triple bond1 sigma and 2 pi bonds

97
Sigma/Pi Bonds

• Link for animations
• http//www.mhhe.com/physsci/chemistry/animations/c
  hang_7e_esp/bom5s2_6.swf