1. Chapter 6
Chemical Bonding

2. Introduction to chemical bonding
Section 1:
Introduction to chemical bonding

3. Introduction to chemical bonding
What is a chemical bond???
A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together

4. Introduction to chemical bonding
Why do atoms bond?
They are working to achieve more stable arrangements where the bonded atoms will have lower potential energy than they do when existing as individual atoms- increase stability.

5. Introduction to chemical bonding
Types of Chemical Bonding:
1. Ionic – an electrical attraction that forms between cations (+) and anions (-)
2. Covalent – are formed when electrons are shared between atoms
3. Metallic – formed by many atoms sharing many electrons

6. Introduction to chemical bonding
However….
Bonds are never purely covalent or purely ionic.
The degree of ionic-ness or covalent-ness depends on property of electronegativity.

7. Degree of Ionic/Covalent Character in Chemical Bonds
100%
50% 5% 0%
Ionic
Polar-Covalent
Nonpolar-Covalent

8. Introduction to chemical bonding
Recall what electronegativity is:
The ability or degree of attraction that an atom has to electrons that are within a bonded compound.
(see page 161)

9. Introduction to chemical bonding
To determine the degree of ionic-ness or covalent-ness you must take each of the electronegativities for the elements in the compound and subtract them.

10. Introduction to chemical bonding
If difference is = nonpolar covalent
If difference is 0.3 – 1.7 = polar covalent
Above 1.7 = Ionic

11. Ionic/Covalent Character Due to Electronegativity Differences
3.3
1.7
0.3
100%
50% 5% 0%
Ionic
Polar-Covalent
Nonpolar-Covalent

12. Introduction to chemical bonding
= 0.4
Polar Covalent
= 1.8
Ionic
2.5 – 3.0 = 0.5
Sulfur + Hydrogen
Sulfur + Cesium
Sulfur + Chlorine

13. Introduction to chemical bonding
In general however…
If bonding elements are on opposite sides of the periodic table (metal with a nonmetal) then they tend to be ionic.
If elements are close together (nonmetal to nonmetal), then they tend to be covalent.

14. Covalent Bonding & Molecular Compounds
Section 2:
Covalent Bonding & Molecular Compounds

15. Covalent Bonding What is a molecule?
A neutral group of atoms that are held together by covalent bonds.
May be different atoms such as H2O or C6H12O6
May be the same atoms such as O2

16. Covalent Bonding
Molecular compounds are made of molecules ….. Not ions!
We represent covalent or molecular compounds by chemical formulas that show numbers of atoms of each kind of element in the compound. CH4 - methane

17. Covalent Bonding
Diatomic molecules are those elements that exist in pairs of like atoms that are bonded together.
There are 7 diatomic molecules:
H N O F2 Cl2 I2 Br2
Big 7

19. Covalent Bonding Formation of a covalent bond:
When atoms are far apart they do not attract – potential energy is zero.
As they come closer the electrons are attracted to protons but electrons and electrons repel – but e- to p attraction is stronger!

20. Covalent Bonding
The electron clouds of the bonded atoms are overlapped and form a “bond length.”

21. Covalent Bonding
Energy is released when these atoms join together with a bond.
Energy must be added to separate these atoms into neutral isolated atoms – called bond energies.
Bond energy is expressed in kilojoules per mole.

22. Covalent Bonding
Octet Rule – Atoms will either gain, lose, or share electrons so that their outer energy levels will contain eight electrons (H is an exception since it can only have 2 in the outer level).
These electrons that are being gained, lost, or shared are represented by using the electron dot diagrams.

23. Examples of electron dot notations
1 valence electron
3 valence electrons
5 valence electrons
7 valance electrons X X X X

24. Covalent Bonding Shared electron pairs and unshared pairs:
Cl:Cl Shared pair
Unshared pairs

25. Covalent Bonding
These electron dot representations are called Lewis structures.
Dots represent the valence electrons

26. Cl - Cl Covalent Bonding
Lewis structures can also be represented using structural formulas.
Dashes indicate bonds of shared electrons (unshared e- are not shown
Cl - Cl
One pair (2 e-) is shared here.

27. Steps To Drawing Lewis Structures
Calculate the number of valance electrons.
Arrange atoms.
Compare number of electrons used with number of electrons available.
Check octet rule.
Change dots to dashes where appropriate.

28. Covalent Bonding
Lewis structure for ammonia (NH3)

29. Covalent Bonding Practice: Draw Lewis structure for methane CH4
Ammonia NH3
Hydrogen Sulfide H2S
Phosphorus trifluoride PF3

30. More Guidelines
H and halogen atoms usually bond to only one other atom in a molecule and are usually on the outside or end of a molecule (each only need 1 electron to form stable octet and electronegativity)

31. More Guidelines
The atom with the smallest electro-negativity is often the central atom
When a molecule contains more atoms of 1 element than the other, these atoms often surround the central atom

32. Covalent Bonding
Some atoms can form multiple bonds – especially C, O, & N.
Double bonds are bonds that share 2 pair of electrons
C=C means C::C
Triple bonds share 3 pair
C≡C means C:::C

33. Covalent Bonding Resonance:
Some substances cannot be drawn correctly with Lewis structure diagrams
Some electrons share time with other atoms – ex. Ozone – O3

34. Covalent Bonding Electrons in ozone may be represented as: O = O–O
Other times it may be represented as O–O=O
Actually these structures are shared – electrons “resonate” (go back & forth) between them

35. Ionic Bonding and Ionic Compounds
Section 3:
Ionic Bonding and Ionic Compounds

36. Section 3: Ionic Bonding & Compounds
Ionic compounds are formed of positive and negative ions
When combined these charges equal zero
Ex: Na = 1+
Cl = 1-
0 charge

37. Section 3: Ionic Bonding & Compounds
Ionic substances are usually solids
Ionic solids are generally crystalline in shape
An ionic compound is a 3-D network of + and – ions that are attracted to each other

38. Section 3: Ionic Bonding & Compounds
Crystals in ionic compounds exist in orderly arrangements known as a crystal lattice.

39. Section 3: Ionic Bonding & Compounds
Ionic substances are not referred to as “molecules”
Ionic substances are referred to as “formula units”
A formula unit is the simplest ratio of the ions that are bonded together.

40. Section 3: Ionic Bonding & Compounds
The ratio of ions depends on the charges.
What would result when F- combines with Ca2+?
CaF2

41. Section 3: Ionic Bonding & Compounds
When ions are written using electron dot structures the dots are written and symbols for their charges.
Na.  Na+
Cl 

42. Compared to molecular compounds, ionic compounds:
Have very strong attractions
Are hard, but brittle
Have higher melting points and boiling points
When dissolved or in the molten state they will conduct electricity

43. Polyatomic Ions:
A group of atoms covalently bonded together but with a charge.
Sulfate SO42-
Carbonate CO32-
Nitrate NO3-
Ammonium NH4+

44. Section 4:
Metallic Bonding

45. Metallic Bonding
Metals are excellent electrical conductors in the solid state.
This is due to highly mobile valence electrons that travel from atom to atom. e-

46. Metallic Bonding Generally metals have either 1 or 2 s electrons
p orbitals are vacant
Many are filling in the d level
Electrons become delocalized and move between atoms (sea of electrons)

47. Metallic Bonding
A metallic bond is the mutual sharing of many electrons among many atoms.

48. Metallic Properties High electrical conductivity
High thermal conductivity
High luster
Malleable (can be hammered or pressed into shape)
Ductile (capable of being drawn or extruded through small openings to produce a wire)

49. Metallic Bond Strength
Varies with nuclear charge and number of electrons shared.
High bond strengths result in high heats of vaporization (when metals are changed into gaseous phase)

50. Section 5:
Molecular Geometry

51. Molecular geometry…
A molecule’s properties depend on bonding of atoms, but also the molecular geometry.

52. Is the three dimensional arrangement of a molecule’s atoms in space.
Molecular geometry…
Is the three dimensional arrangement of a molecule’s atoms in space.

53. VSEPR Theory Valence Shell Electron Pair Repulsion
Electrons around a nucleus repel each other to be as far away from each other as possible.

54. Lone pairs repel more strongly than bonding pairs!!!
VSEPR Theory
Types of e- Pairs
Bonding pairs - form bonds
Lone pairs - nonbonding e-
Lone pairs repel more strongly than bonding pairs!!!

55. Know the common shapes & their bond angles! Draw the Lewis Diagram.
Tally up e- pairs on central atom.
double/triple bonds = ONE pair
Shape is determined by the # of bonding pairs and lone pairs.
Know the common shapes
& their bond angles!

56. Common Molecular Shapes
2 total
2 bond
0 lone
BeH2
LINEAR
180°

57. Common Molecular Shapes
3 total
3 bond
0 lone
BF3
TRIGONAL PLANAR
120°

58. Common Molecular Shapes
4 total
4 bond
0 lone
CH4
TETRAHEDRAL
109.5°

59. Common Molecular Shapes
4 total
3 bond
1 lone
NH3
TRIGONAL PYRAMIDAL
107°

60. Common Molecular Shapes
4 total
2 bond
2 lone
H2O
BENT
104.5°

61. Examples
F P F F
PF3
4 total
3 bond
1 lone
TRIGONAL PYRAMIDAL
107°

62. Examples
O C O
CO2
2 total
2 bond
0 lone
LINEAR
180°

63. Hybridization
Explains how atom’s orbitals become rearranged to form covalent bonds.
Hybridization is the mixing of 2 or more orbitals of similar energies on the same atom to produce new orbitals of equal energies.

64. Hybridization Methane (CH4) is an example of hybridization:
Carbon’s normal configuration is 2s22p2
In methane all the electrons in the 2nd energy level become equal in energy and is referred to as sp3

65. Intermolecular Forces
Intermolecular forces are attractive forces between molecules.
Intramolecular forces hold atoms together in a molecule.
Intermolecular vs Intramolecular
41 kJ to vaporize 1 mole of water (inter)
930 kJ to break all O-H bonds in 1 mole of water (intra)
“Measure” of intermolecular force
boiling point
melting point
Generally, intermolecular forces are much weaker than intramolecular forces.
11.2

66. Intermolecular Forces:
Strong IM forces exist in polar molecules.
Polar molecules act as tiny “dipoles” (equal & opposite charges separated by short distances)

67. Dipole Forces Types of Intermolecular Forces
Attractive forces between polar molecules
Orientation of Polar Molecules in a Solid
11.2

68. Types of IMF
Dipole-Dipole Forces + -
View animation online.

69. Dipole Forces Types of Intermolecular Forces
Attractive forces between an ion and a polar molecule
Ion-Dipole Interaction
11.2

70. Intermolecular Forces:
Another IM force is Hydrogen bonding.
Is the strongest type of dipole-dipole force
Explains high boiling points of H-containing substances such as water and ammonia

71. Intermolecular Forces:
In hydrogen bonding, a hydrogen atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule.

72. Types of Intermolecular Forces
Hydrogen Bond (strongest)
The hydrogen bond is a special dipole-dipole interaction between the hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. IT IS NOT A BOND. A H … B or
A & B are N, O, or F
11.2

73. Types of IMF
Hydrogen Bonding

74. Intermolecular Forces:
London/ Dispersion/VanDerWaals forces:
Are very weak bonds
Occur due to the fact that since electrons are in constant motion that briefly there are moments where electrons are unevenly distributed and thus the molecule briefly has a charged area.

75. Types of IMF
London Dispersion Forces
View animation online.