1. Copyright© by Houghton Mifflin Company. All rights reserved.
Chemical Bonding
Copyright© by Houghton Mifflin Company. All rights reserved.

2. Intramolecular Bonding Intermolecular Bonding
Within molecules
Intermolecular Bonding
Between molecule
Copyright© by Houghton Mifflin Company. All rights reserved.

3. Octet Rule
Atoms tend to gain, lose, or share electrons in order to acquire a full set of valence electrons
A very important rule
An Octet consist of eight electrons
This refers to the outermost eight electrons
Compounds form to satisfy the Octet Rule
Ionic or Covalent Bonds

4. Types of IntramolecularBonds
Ionic
Formed by transfer of electrons
Held together by electrostatic attraction
Covalent
Formed by sharing of electrons
Held together by shared electrons
Metallic
Formed when electron(s) become detached from metal atoms
Sea of electrons among + ions metal of atoms

6. Electronegativity values for selected elements.

7. Electronegativity Difference and Bond Type
(Approximate)
Bond Type
Example
0-0.4
Nonpolar Covalent
H-H
Moderately Polar Covalent
H-Cl
1.0-2
Very Polar Covalent
H-F
≥2.0
Ionic
NaCl

8. Ions as packed spheres.

9. Na+1 + Cl-1  NaCl (table salt)
Ionic Bonds
Ionic bonding occurs when one atom
transfers an electron to another atom
Both atoms involved become charged
One is negatively charged
One is positively charged
This occurs when a metal reacts with a nonmetal
Na+1 + Cl-1  NaCl (table salt)

11. Ions and Charges Cations Positive Ions Anions Negative Ions
Group 1, 2, 3
Form ions with charges equal to their group number
Group 5, 6, 7
Form ions with charges equal to group number minus eight

12. Ionic Compounds and Formulas
The formula of a compound describes what elements are in the compound and in what proportions.
Compounds that are held together by ionic bonds are called ionic compounds.

13. Common polyatomic ion names
Formula
Name
NH4+
Ammonium ion
CO32-
Carbonate ion
PO43-
Phosphate ion
SO42-
Sulfate ion
OH-
Hydroxide ion
NO3-
Nitrate ion

14. Formation of Ionic Compounds
Al SO42-
Al2(SO4)3
Ionic compounds forms such that total of ionic charges is zero.
For above case:
2 x (+3) = 6
3 x (-2) = -6
Total

15. O The Lewis dot structure for Oxygen
Oxygen is in group VIA so it has 6 valence electrons

16. Cl The Lewis dot structure for Chlorine
chlorine is in group VIIA so it has 7 valence electrons

17. Ca The Lewis dot structure for calcium
calcium is in group IIA so it has 2 valence electrons

18. Making calcium chloride+ Ca Cl
CaCl2

19. Covalent Bonds
A covalent bond is a chemical bond that is formed when two atoms share a pair of electrons.
H. + H.  H:H
Covalent Compounds and Formulas
In the above example, each hydrogen has a filled valence shell simulating the electron configuration of helium.
Compounds that are held together by covalent bonds are called covalent compounds.
Covalent compounds form from atoms on the right side of the periodic table

20. Multiple Bonds.
In electron dot notations, a pair of electrons can be represented by a pair of dots : .
This can be a bonding pair or a lone pair (non-bonding pair).
Bonding pairs can also be represented by lines connecting atoms.
H:H = H-H
When one pair of electrons is shared, it is called a single bond.
H-H

21. When two pairs of electrons are shared it is called a double bond.
When three pairs of electrons are shared it is called a triple bond.

22. Bond Length
Bond Length varies for covalent bonds between different elements
See p. 187
Two important trends
As one moves down a group, bond length between atoms increases
F—F nm
Cl—Cl nm
Br—Br nm
Multiple bonds are shorter than single bonds
Example: two carbon atoms bonded together
Single covalent bond nm
Double covalent bond nm
Triple covalent bond nm

24. Exceptions to the Octet Rule
Exceptions to the octet rule include those for atoms that cannot fit eight electrons, and for those that can fit more than eight electrons, into their outermost orbital.
Hydrogen forms bonds in which it is surrounded by only two electrons.
Boron has just three valence electrons, so it tends to form bonds in which it is surrounded by six electrons.
Main-group elements in Periods 3 and up can form bonds with expanded valence, involving more than eight electrons, e.g. PF5 and SF6

25. Electron-Dot Notation
To keep track of valence electrons, it is helpful to use electron-dot notation.
Electron-dot notation is an electron-configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol. The inner-shell electrons are not shown.

26. Lewis Structures Chapter 6
Electron-dot notation can also be used to represent molecules.
The pair of dots between the two symbols represents the shared electron pair of the hydrogen-hydrogen covalent bond.
For a molecule of fluorine, F2, the electron-dot notations of two fluorine atoms are combined.

27. Chapter 6 Lewis Structures
The pair of dots between the two symbols represents the shared pair of a covalent bond.
In addition, each fluorine atom is surrounded by three pairs of electrons that are not shared in bonds.
An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom.
See Example, p. 185

28. Multiple Covalent Bonds
Chapter 6
Multiple Covalent Bonds
Double and triple bonds are referred to as multiple bonds, or multiple covalent bonds. (See Table 2, p. 187)
In general, double bonds have greater bond energies and are shorter than single bonds.
Triple bonds are even stronger and shorter than double bonds.
When writing Lewis structures for molecules that contain carbon, nitrogen, or oxygen, remember that multiple bonds between pairs of these atoms are possible.

29. Molecular Shape Arises because electrons repulse one another
Called VSEPR
Valence Shell Electron Pair Repulsion
VSEPR states that in a small molecule, the pairs of valence electrons are arranged as far apart from each other as possible

30. Summary of Molecular Shapes
Type
Bond
Angle
Unshared Pairs
Example
Linear
180°
Balanced
(2 each side)
CO2
Trigonal
Planar
120°
none
BCl3
Tetrahedral
109.5°
CH4
Pyramidal
107°
one
NH3
Bent
105°
two
H2O

31. Bent molecular structure of the water molecule.

32. A closer look at a water molecule (bent)

33. Molecular structure of methane (Tetrahedral).

34. Pyramidal

35. 8.3
Molecular Orbitals
When two atoms combine, the molecular orbital model assumes that their atomic orbitals overlap to produce molecular orbitals, or orbitals that apply to the entire molecule.
Just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a molecule as a whole.
A molecular orbital that can be occupied by two electrons of a covalent bond is called a bonding orbital.

36. Molecular Orbitals 8.3 Sigma Bonds
When two atomic orbitals combine to form a molecular orbital that is symmetrical around the axis connecting two atomic nuclei, a sigma bond is formed.

37. 8.3
Molecular Orbitals
When two fluorine atoms combine, the p orbitals overlap to produce a bonding molecular orbital. The F—F bond is a sigma bond.
Two p atomic orbitals can combine to form a sigma-bonding molecular orbital, as in the case of fluorine (F2). Notice that the sigma bond is symmetrical around the bond axis connecting the nuclei.

38. Molecular Orbitals 8.3 Pi Bonds
In a pi bond (symbolized by the Greek letter ), the bonding electrons are most likely to be found in sausage-shaped regions above and below the bond axis of the bonded atoms.

39. Polarity Bonds can be polar or nonpolar
Depends on electronegativity difference
Molecules can be polar or nonpolar
Polar molecules are called dipoles
One end of a polar molecule has “+” charge; other end has “-” charge
Will align in electric field
Will be attracted or deflected by magnetic field
Polarity of a molecule is determined by shape and the type of bonds between its atoms

40. The chlorine atom attracts the electron cloud more than the hydrogen atom does. Why?

41. Polar Molecules 8.4 A hydrogen chloride molecule is a dipole.
When polar molecules, such as HCl, are placed in an electric field, the slightly negative ends of the molecules become oriented toward the positively charged plate and the slightly positive ends of the molecules become oriented toward the negatively charged plate. Predicting What would happen if, instead, carbon dioxide molecules were placed between the plates? Why?

42. Intermolecular Forces
Attraction between molecules
Holds groups of molecules together
Intermolecular forces are known as van der Waals forces
Dipole-Dipole attraction
Dispersion Forces
Hydrogen Bonds

43. van der Waals Forces Dipole-Dipole attraction Dispersion Forces
Electrostatic attraction between polar molecules
Molecules line up like magnets
Dispersion Forces
Polarity arises due to electron imbalance
Weak electrostatic forces arise
Hydrogen Bonds
Strong dipoles arising from covalent bonds of hydrogen atom to a very electronegative atom
H electronegativity = 2.1 F=4.0
Results in liquids with high boiling points

44. Probability representations of the electron sharing in HF.

45. Charge distribution in the water molecule.

46. Water molecule behaves as if it had a positive and negative end.

47. Polarity Formaldehyde CH2O Carbon Dioxide CO2 Water H2O
Forms dipole due to high electronegativity of O
Carbon Dioxide CO2
Has polar bond but they cancel out
CO2 is nonpolar
Water H2O
Water is a bent molecule due to unshared pairs
unlike CO2 bond polarities do not cancel out
Water is polar and forms dipole
Why is water liquid and Carbon Dioxide a gas?

48. Large Molecules Examples: protein, Subunits linked together in a chain
Often bend and twist to form 3D shape
Chains, rings, balls

49. The three states of water:

51. Table 12.3

52. Relative sizes of some ions and their parent atoms.

53. The three possible types of bonds.

54. Figure 12.1: The formation of a bond between two atoms.

55. Tetrahedral arrangement of electron pairs.
109.5°

56. Tetrahedral arrangement of four electron pairs around oxygen.

58. Table 12.1

59. Figure 12.6: Polar water molecules are strongly attracted to negative ions by their positive ends.

60. Figure 12.6: Polar water molecules are strongly attracted to positive ions by their negative ends.

61. The NH3 molecule pyramidal structure.

62. Bond Energies and Bond Lengths for Single Bonds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Bond Energies and Bond Lengths for Single Bonds

63. Hybrid Orbitals When atoms bond, their outer orbitals become distorted
Orbitals do not look like those of unbonded atom
Hybrid orbitals are formed which are a cross between the bonding orbitals sp
sp2
sp3
Orbital hybridization contributes to shape of molecule

64. Hybrid Orbitals