Covalent Bonds and Molecular Structure

1
Chapter 7

• Covalent Bonds and Molecular Structure

2
The Covalent Bond

• Covalent bond formed by the sharing of
  electrons between two nonmetal atoms
• Forces involved in the bond
• Electrostatic attraction between proton and
  electron
• Electrostatic repulsion between electron and
  electron
• When will a bond form?

3
Strengths of Bonds

• Bond dissociation energy the amount of energy
  required to break a bond
• Energy increases as the length of the bond gets
  shorter

4
Bond length and covalent radius.
5
Strengths of Bonds

• Single bonds gt double bonds gt triple bonds

6
Problem

• Arrange the following bonds in order of
  increasing bond strength. 
• A.  C-I lt C-Br lt C-Cl lt C-F
• B.  C-F lt C-Cl lt C-Br lt C-I
• C.  C-Br lt C-I lt C-Cl lt C-F
• D.  C-I lt C-Br lt C-Flt C-Cl
• E.  none of these orders is correct

7
Problem

• Select the strongest bond in the following
  group. 
• A.  C-S
• B.  C-O
• C.  CC
• D.  CN
• E.  C-F

8
Electron Dot Structures

• Electron dot symbols
• Aid in understanding the formation of bonds
  between atomic nuclei
• Elemental symbol represents the type of element
  and all core electrons the valence electrons are
  represented by dots around the symbol

9
Electron Dot Structures

• A metal in an ionic loses its electrons to
  achieve an octet or pseudo-octet (transition
  elements) in its outermost shell
• A nonmetal in an ionic compound gains electrons
  to achieve an octet in its outermost shell
• Period 1 and 2 elements of a covalent compound
  share enough electrons to achieve an octet

10
Electron-Dot Structures

• Ionic
• Covalent

11
Electron Dot Structures

• Using electron dot symbols build
• H2, H2O, CH4, O2, N2, HCN, CO2
• Least electronegative atom is often central
  (except H)

12
Naming Binary Molecular Compounds

• Electronegativity indicates how well an
  elements nuclei attract the electrons in a
  covalent bond

13
The Periodic Table and Electronegativity
14
Electron Dot Structures

• Using electron dot symbols build
• H2, H2O, CH4, O2, N2, HCN, CO2
• Least electronegative atom is often central
  (except H)

15
Electron Dot Structures

• Single bond A covalent bond formed by sharing
  one electron pair.
• Double bond A covalent bond formed by sharing
  two electron pairs.
• Triple bond A covalent bond formed by sharing
  three electron pairs.
• Single bonds are longer (weaker) than double
  bonds
• Double bonds are longer (weaker) than triple bonds

16
Electron-dot Structures

• Step 1 Count the total valence electrons.
• Step 2 Identify the central atom
• - Often least electronegative
• Step 3 Place all other atoms around the central
• atom
• Step 4 Draw a single bond between each
• external atom and the central atom
• subtracting 2 electrons for each
  bond
• drawn from the total valence electrons.

17
Electron-dot Structures

• Step 5 Distribute remaining valence
• electrons around the external
• atoms giving the external atoms
• an octet
• Step 6 If valence electrons still remain,
• place them on the central atom in pairs
• Step 7 Verify that each atom has an octet
• Hydrogen needs only 2 electrons
• Boron needs only 6 electrons

18
Electron-Dot Structures

• Step 8 If the central atom does not have an
• octet, form a multiple bond by bringing
  a
• pair of electrons in from the external
• atom
• Step 9 Calculate formal charge and minimize
• the formal charge if acceptable
• Period 3 elements and greater can have expanded
  octets if one is necessary to minimize formal
  charge
• Formal charge valence electrons for the atom
  1 for every dot on the atom 1 for every line
  around the atom

19
Problems

• BF3
• PF3
• C2H6
• I3
• NH4
• SO42-
• KClO3

20
Electron-dot Structures and Resonance

• How is the double bond formed in O3?
• The correct answer is that both are correct, but
  neither is correct by itself.

21
Electron-Dot Structures and Resonance

• When multiple structures can be drawn, the actual
  structure is an average of all possibilities.
• The average is called a resonance hybrid. A
  straight double-headed arrow indicates resonance.

22
Problem

• S3
• PO43-
• CO32-
• NO2

23
Molecular Shapes The VSEPR Theory

• The approximate shape of molecules is given by
  Valence-Shell Electron-Pair Repulsion (VSEPR).

24
Molecular Shapes The VSEPR Theory
Molecular formula
Step 1
Count all e- groups around central atom (A)
Single, Double and Triple bonds are all counted
as 1 e- group
Lewis structure
Step 2
Electron-group arrangement
Note lone pairs and double bonds
Step 3
Count bonding and nonbonding e- groups separately.
Bond angles
Step 4
Molecular shape (AXmEn)
25
Problem

• Determine the shape of the molecules for which
  Lewis Structures have been developed.

26
Valence Bond Theory

• If, in order for a bond to form, a pair of
  electrons must be shared, then how does C form
  molecules with 4 bonds?
• Valence Bond Theory hybrid orbitals

27
Valence Bond Theory
Basic Principle
A covalent bond forms when the orbtials of two
atoms overlap and are occupied by a pair of
electrons that have the highest probability of
being located between the nuclei.
Themes
A set of overlapping orbitals has a maximum of
two electrons that must have opposite spins.
The greater the orbital overlap, the stronger
(more stable) the bond.
The valence atomic orbitals in a molecule are
different from those in isolated atoms.
28
Valence Bond Theory
The number of hybrid orbitals obtained equals the
number of atomic orbitals mixed.
The type of hybrid orbitals obtained varies with
the types of atomic orbitals mixed.
29
Valence Bond Theory
The conceptual steps from molecular formula to
the hybrid orbitals used in bonding.
Molecular shape and e- group arrangement
Molecular formula
Lewis structure
Hybrid orbitals
30
Problems

• Carbon uses ______ hybrid orbitals in ClCN. 
• A.  sp
• B.  sp2
• C.  sp3
• D.  sp3d
• E.  sp3d2

31
The s bonds in ethane.
32
The s and p bonds in ethylene (C2H4)
33
The s and p bonds in acetylene (C2H2)
34
Polar Covalent Bonds Electronegativity

• Electronegativity represents the ability of an
  atom to attract a shared pair of electrons
• Higher the EN the more the electrons in a bond
  will be pulled toward the atom
• Most electronegative atom is F
• EN ? down a group
• EN? across a period from left to right w/ few
  exceptions

35
Polar Covalent Bonds Electronegativity
36
Problem

• Which of the following elements is the most
  electronegative? 
• A.  S
• B.  Ru
• C.  Si
• D.  Te
• E.  Cs

37
Problem

• Arrange calcium, rubidium, sulfur, and arsenic in
  order of decreasing electronegativity. 
• A.  S gt As gt Rb gt Ca
• B.  S gt As gt Ca gt Rb
• C.  As gt S gt Rb gt Ca
• D.  As gt S gt Ca gt Rb
• E.  None of these orders is correct.

38
Polar Covalent Bonds Electronegativity

• Ionic Character As a general rule for two
  atoms in a bond, we can calculate an
  electronegativity difference (?EN ) ?EN EN(Y)
  EN(X) for XY bond.
• If ?EN lt 0.5 the bond is covalent.
• If ?EN 0.5 - lt 2.0 the bond is polar covalent.
• If ?EN gt 2.0 the bond is ionic.

39
Problem

• Select the most polar bond amongst the
  following. 
• A.  C-O
• B.  Si-F
• C.  Cl-F
• D.  C-F
• E.  C-I

40
Molecular Orbital Theory

• The molecular orbital (MO) model provides a
  better explanation of chemical and physical
  properties than the valence bond (VB) model.
• Atomic Orbital Probability of finding the
  electron within a given region of space in an
  atom.
• Molecular Orbital Probability of finding the
  electron within a given region of space in a
  molecule.

41
Molecular Orbital Theory

• Additive combination of orbitals (s) is lower in
  energy than two isolated 1s orbitals and is
  called a bonding molecular orbital.

42
Molecular Orbital Theory

• Subtractive combination of orbitals (s) is
  higher in energy than two isolated 1s orbitals
  and is called an antibonding molecular orbital.

43
Molecular Orbital Theory

• Molecular Orbital Diagram for H2

44
Molecular Orbital Theory

• Molecular Orbital Diagrams for H2 and He2

45
Molecular Orbital Theory

• Additive and subtractive combination of p
  orbitals leads to the formation of both sigma and
  pi orbitals.

46
Molecular Orbital Theory

• Second-Row MO Energy Level Diagrams

47
Molecular Orbital Theory

• Bond Order is the number of electron pairs shared
  between atoms.
• Bond Order is obtained by subtracting the number
  of antibonding electrons from the number of
  bonding electrons and dividing by 2.