1. Chapter 6
Chemical Bonding

2. Intro to Chemical Bonding
Why do elements combine? (bond)
In purely ionic bonding, atoms completely give up electrons to other atoms.
Atoms joined by purely covalent bonding share electrons equally.
The electrons are “owned” equally by two atoms

3. Ionic or Covalent?
How can we tell whether or not a bond is ionic or covalent?
It has to do with electronegativity values.
It is not as simple as metal/nonmetal or nonmetal/nonmetal bonding
There are grey areas

4. Ionic or Covalent?
Bonding usually falls somewhere between these two extremes.
Depends on electronegativities of the atoms involved
Recall electronegativity is a measure of an atom’s ability to attract an electron
Electronegativity tends to increase going right and decrease going down the periodic table

5. Ionic or Covalent?
The degree to which bonding between atoms of two elements is ionic or covalent can be estimated by calculating the difference in electronegativity values
Elements are given electronegativity values from 0 to 4.
Ended pd 3
Ended pd 6

6. Ionic or Covalent?
Bonding between atoms with an electronegativity difference of 1.7 or less has a more covalent nature.
Classified as covalent compounds.
Ended pd4

7. Ionic or Covalent?
Bonding between two atoms of the same element is considered completely covalent or a nonpolar-covalent bond.
A nonpolar-covalent bond- a covalent bond in which the electrons are shared equally resulting in a balanced distribution of electrons.

8. Ionic or Covalent?
Bonding between atoms with an electronegativity difference of 0 to 0.3 are generally considered non-polar covalent.
Bonds have 0% to 5% ionic character.

9. Ionic or Covalent?
Bonding between atoms with an electronegativity difference of 0.3 to 1.7 are classified as polar covalent.
A polar-covalent bond is a bond in which the bonded atoms have an unequal attraction for the shared electrons.

10. Ionic or Covalent? Compare Chlorine and Hydrogen
Chlorine has a electronegativity value of 2.9
Hydrogen has electronegativity value of 2.1
What type of bond is this?

11. Characteristics of the Covalent Bond
Bond energy is the energy required to break a chemical bond and form neutral isolated atoms.
Bigger atoms make longer bonds
General trend -->Less energy to break bonds if bonds are long
H-F bond length of 92 pm bond energy 569 kj/mol
H-Cl 127 pm 432 kj/mol
H-Br 141 pm 366 kj/mol
H-I 161 pm 299 kj/mol

12. Characteristics of Covalent Bonds
Double and single bonds will make bonds shorter
Triple Bonds will have more bond energy
C-C 154 pm 346 kj/mol
C=C 134 pm 612 kj/mol
CΞC 120 pm 835 kj/mol

13. Octet Rule Atoms want to acquire eight valence electrons
Why? We don’t know

14. Exceptions to Octet Rule
Hydrogen will surround itself with 2 valence Ex/ H20
Boron will surround itself with 6 electrons
BF3
Phosphorus and Sulfur when bound to F,O,Cl
Ex/ PF5 SF6

15. How many bonds?
Think of how many electrons they are away from noble gas.
H should form 1 bond- always
O should form 2 bonds – if possible
N should form 3 bonds – if possible
C should form 4 bonds– if possible

16. Practice Draw electron dot diagrams for the following. PCl3 H2O2 CH2O

17. Resonance
Some molecules and ions cannot be represented by a single Lewis structure.
Ex/ 03 CO3-2 NO2-1
Chemists once speculated that molecules resonate between the two structures. New research suggests the bond are identical. Think of it’s correct structure as the average of both Lewis structure.
Ended 5th
Ended 6th

18. Properties of Ionic Compounds
Crystalline structure.
A regular repeating arrangement of ions in the solid.
Ions are strongly bonded.
Structure is rigid.
High melting points- because of strong forces between ions.

19. Crystalline structure
3 dimensions

20. Do they Conduct? Conducting electricity is allowing charges to move.
In a solid, the ions are locked in place.
Ionic solids are insulators.
When melted, the ions can move around.
Melted ionic compounds conduct.
First get them to 800ºC.
Dissolved in water they conduct.

21. Metallic Bonds How atoms are held together in the solid.
Metals hold onto their valence electrons very weakly.
Think of them as positive ions floating in a sea of electrons.

22. Sea of Electrons + Electrons are free to move through the solid.
Metals conduct electricity. +

23. Metals are Malleable Hammered into shape (bend).
Ductile - drawn into wires.

24. Malleable +

25. Malleable
Electrons allow atoms to slide by. + + + + + + + + + + + +

26. Ionic solids are brittle+ -

27. Ionic solids are brittle
Strong Repulsion breaks crystal apart. + - + -
End pd 1 + - + -

28. Alloys
Solutions made by dissolving metal into other elements- usually metals.
Melt them together and cool them.
If the atoms of the metals are about the same size, they substitute for each other
Called a substitutional alloy

29. → + Substitutional alloy Metal B Metal A Bronze – Copper and Tin
Brass- 60 % Copper 39% Zinc and 1%Tin
18 carat gold- 75% gold, 25%Ag or Cu

30. Alloys
If they are different sizes the small one will fit into the spaces of the larger one
Called and interstitial alloy

31. → + Interstitial Alloy Metal A Metal B Steel – 99% iron 1 % C
Cast iron- 96% Iron, 4%C

32. Alloys Making an alloy is still just a mixture Blend the properties
Still held together with metallic bonding
Most of the metals we use daily are alloys.
Designed for a purpose

33. Crystal Structures
The repeating unit is called the unit cell

34. Cubic

35. Body-Centered Cubic

36. Face-Centered Cubic

37. VSEPR Valence Shell Electron Pair Repulsion.
Predicts three dimensional geometry of molecules.
Name tells you the theory.
Valence shell - outside electrons.
Electron Pair repulsion - electron pairs try to get as far away as possible.
Can determine the angles of bonds.
And the shape of molecules

38. VSEPR
Based on the number of pairs of valence electrons both bonded and unbonded.
Unbonded pair are called lone pair.
CH4 - draw the structural formula

39. H H C H H VSEPR Single bonds fill all atoms.
There are 4 pairs of electrons pushing away.
The furthest they can get away is º. H C H H

40. H C H H H 4 atoms bonded Basic shape is tetrahedral.
A pyramid with a triangular base.
Same basic shape for everything with 4 pairs. H
109.5º C H H H

41. N H N H H H H H 3 bonded - 1 lone pair
Still basic tetrahedral but you can’t see the electron pair.
Shape is called trigonal pyramidal. N H N H H H
<109.5º H H

42. O H O H H H 2 bonded - 2 lone pair
Still basic tetrahedral but you can’t see the 2 lone pair.
Shape is called bent. O H O H
<109.5º H H

43. H C H O O C H 3 atoms no lone pair
The farthest you can the electron pair apart is 120º.
Shape is flat and called trigonal planar.
Will require 1 double bond C H O
120º H O C H

44. 2 atoms no lone pair
With three atoms the farthest they can get apart is 180º.
Shape called linear.
Will require 2 double bonds or one triple bond C O
180º

45. Hybridization The process of orbital mixing
The reason we spread out electrons in Lewis Dot structures
The valence atomic orbitals in the molecule are different from those in the isolated atoms.
Valence electrons “spread out” to form equal orbitals

46. Hybridization Carbon is the central atom in methane CH4
It forms 4 identical orbitals to bind with Hydrogen
Carbon’s electron configuration is
1s22s22p2
Carbon hybridizes to form an sp3 configuration

47. Depends on Molecular Geometry
Arrangement Hybridization
linear sp
trigonal planar sp2
bent sp3
trigonal pyramidal sp3
tetrahedral sp3

48. Hybridization All the hybrid orbitals that form are the same.
sp3 -1 s and 3 p orbitals mix to form 4 sp3 orbitals.
sp2 -1 s and 2 p orbitals mix to form 3 sp2 orbitals leaving 1 p orbital.
sp -1 s and 1 p orbitals mix to form 2 sp orbitals leaving 2 p orbitals.

49. Molecular Orbitals
The overlap of atomic orbitals from separate atoms makes molecular orbitals
Each molecular orbital has room for two electrons
Two types of MO
Sigma ( σ ) between atoms
Pi ( π ) above and below atoms

50. Two types of Bonds All single bonds are σ bonds
Double bond is 1 σ and 1 π bond
Triple bond is 1 σ and 2 π bonds

51. Sigma bonding orbitals
From s orbitals on separate atoms +
+ + + + +
Sigma bonding molecular orbital
s orbital
s orbital

52. Sigma bonding orbitals
From p orbitals on separate atoms ⊕ ⊕
p orbital
p orbital ⊕ ⊕ ⊕ ⊕
Sigma bonding molecular orbital

53. Pi bonding molecular orbital
Pi bonding orbitals
P orbitals on separate atoms ⊕ ⊕ ⊕ ⊕ ⊕ ⊕ ⊕ ⊕
Pi bonding molecular orbital

54. Sigma and pi bonds All single bonds are sigma bonds
A double bond is one sigma and one pi bond
A triple bond is one sigma and two pi bonds.

55. CO2
C can make two σ and two π
O can make one σ and one π O C O

56. How to show a bond is polar
Isn’t a whole charge just a partial charge
δ+ means a partially positive
δ− means a partially negative
The Cl pulls harder on the electrons
The electrons spend more time near the Cl δ+ δ− H Cl

57. Polar Molecules
Molecules with an unequal distribution of electrons

58. Polar Molecules
Molecules with a partially positive end and a partially negative end
Requires two things to be true
The molecule must contain polar bonds
This can be determined from differences in electronegativity.
Symmetry can not cancel out the effects of the polar bonds.
Must determine geometry first.

59. Polar Molecules
Symmetrical shapes are those without lone pair on central atom
Tetrahedral
Trigonal planar
Linear
Will be nonpolar if all the atoms are the same
Shapes with lone pair on central atom are not symmetrical
Can be polar even with the same atom

60. Is it polar? HF
H2O
NH3
CCl4
CO2
CH3Cl

61. - + H - F H - F H - F H - F H - F H - F H - F H - F δ+ δ- δ+ δ- δ+ δ-

62. Intermolecular Forces
What holds molecules to each other

63. Intermolecular Forces
They are what make solid and liquid molecular compounds possible.
The weakest are called van der Waal’s forces - there are two kinds
Dispersion forces
Dipole Interactions

64. Dispersion Force F2 is a gas Br2 is a liquid I2 is a solid
Depends only on the number of electrons in the molecule
Bigger molecules more electrons
More electrons stronger forces
F2 is a gas
Br2 is a liquid
I2 is a solid

65. Dipole interactions
Occur when polar molecules are attracted to each other.
Slightly stronger than dispersion forces.
Opposites attract but not completely hooked like in ionic solids.

66. Dispersion force H δ+ δ- H δ+ δ- δ− H

67. Hydrogen bonding
Are the attractive force caused by hydrogen bonded to F, O, or N.
F, O, and N are very electronegative so it is a very strong dipole.
They are small, so molecules can get close together
The hydrogen partially share with the lone pair in the molecule next to it.
The strongest of the intermolecular forces.

68. Hydrogen Bonding H O δ+ δ- H O δ+ δ-

69. Hydrogen bonding H O H O H O H O H O H O H O

70. Properties of Molecular Compounds
Made of nonmetals
Poor or nonconducting as solid, liquid or aqueous solution
Low melting point
Two kinds of crystals
Molecular solids held together by IMF
Network solids- held together by bonds
One big molecule (diamond, graphite)