1. Chemical Bonding Bonding Theories are applied to design
Atoms interact with other atoms to form molecules, this is chemical bonding
Bonding theories – are models that predict how atoms bond together to form molecules
Bonding Theories are
applied to design
molecules that will
interfere with the active
site of HIV-protease.
This delays or inhibits
the onset of AIDS.

2. OUTLINE Chemical Bonds Ionic Bonds and Covalent Bonds
Electronegativity
Bond Polarity & Electronegativity
Lewis Structures
Resonance
Molecular Shapes
Molecular Polarity

3. CHEMICAL BOND
The nature and type of the chemical bond is directly responsible for many physical and chemical properties of a substance: (e.g. melting point, conductivity).
Most matter in nature is found in form of compounds: 2 or more elements held together through a chemical bond.
Elements combine together (bond) to fill their outer energy levels and achieve a stable structure (low energy).
Noble gases are un-reactive since their energy levels are complete.

4. CHEMICAL BOND
When the conductivity apparatus is placed in salt solution, the bulb will light.
But when it is placed in sugar solution, the bulb does not light.
This difference in conductivity between salt and sugar is due to the different types of bonds between their atoms.
Two common types of bonding are present: ionic & covalent.

5. Lewis is known for:
Covalent bond Lewis dot structures Valence bond theory Electronic theory of acids and bases Chemical thermodynamics Heavy water Named photon Explained phosphorescence
Gilbert Newton Lewis ( ) was a famous American physical chemist
known for the discovery of the covalent bond
(see his Lewis dot structures and his 1916 paper "The Atom and the Molecule")
Other major contributions were his theory of Lewis acids and bases and
Lewis coined the term "photon" for the smallest unit of radiant energy.

6. The Origin of Lewis Symbols of Atoms
Drawings of cubical atoms, the corners of the cube represented possible electron positions
Lewis later cited these notes in his classic 1916 paper on chemical bonding, as being the first expression of his ideas.

7. LEWIS SYMBOLS OF ATOMS
Lewis structures use Lewis symbols to show valence electrons in molecules and ions of compounds.
In Lewis symbols, valence electrons for each element are shown as a dot.
Lewis symbols for the first 3 periods of representative elements are shown below:

8. Lewis Bonding Theory
atoms bond because it results in a more stable electron configuration
atoms bond together by either transferring or sharing electrons so that all atoms obtain an outer shell with 8 electrons
Octet Rule
there are some exceptions to this rule – the key to remember is to try to get an electron configuration like a noble gas

9. Lewis Symbols of Ions
Cations have Lewis symbols without valence electrons
Lost in the cation formation
Anions have Lewis symbols with 8 valence electrons
Electrons gained in the formation of the anion
Li• Li+1 :F: [:F:]-1 • ••

10. Ionic Bonds metal to nonmetal metal loses electrons to form cation
nonmetal gains electrons to form anion
ionic bond results from + to - attraction
larger charge = stronger attraction
smaller ion = stronger attraction
Lewis Theory allow us to predict the correct formulas of ionic compounds

11. IONIC BOND Metal Nonmetal
Ionic bonds occur between metals and non-metals.
After bonding, each atom achieves a complete shell (noble gas configuration).
Ionic bonds occur when electrons are transferred between two atoms.
Metal
Nonmetal

12. IONIC BOND Anion Cation
The smallest particles of ionic compounds are ions (not atoms).
Atoms that lose electrons (metals) form positive ions (cations).
Atoms that gain electrons (non-metals) form negative ions (anions).
Anion
Cation

13. Using Lewis Theory to Predict Chemical Formulas of Ionic Compounds
Predict the formula of the compound that forms between
calcium and chlorine. Cl ∙ Ca ∙
Draw the Lewis dot symbols
of the elements Cl ∙ Cl ∙
Transfer all the valance electrons
from the metal to the nonmetal,
adding more of each atom as you
go, until all electrons are lost
from the metal atoms and all
nonmetal atoms have 8 electrons Ca ∙
Ca2+
CaCl2

14. Ionic Bonding
The transfer of electrons in an ionic bond is from an atom of low ionization energy (usually a metal) to an atom of high electron affinity (usually a nonmetal).
The electrostatic attraction between two oppositely charge ions constitutes an ionic bond.

15. Ionic Bonding & Lattice Structures
A lattice is a stable, ordered, solid three dimensional array of ions associated with ionic compounds.
Lattice energy, ∆Hlattice is the energy required to completely separate a mole of solid ionic compound into its gaseous ions.

16. Lattice Energy Example
For example, the lattice energy of potassium fluoride is 808 kJ/mol.
KF(s) → K+(g)+ F- (g) ∆Hlattice = +808 kJ/mol
Lattice energies can be found in a table as known values.
Lattice energy increases with smaller, more highly charged ions

17. Potential Energy in Bonds
The potential energy of two interacting charged particles is given by the equation
E = KQ1Q2/d
Where Q1 and Q2 are the charges on the ions and d is the distance between them.

18. Characteristics of Ionic Compounds
Ionic Compounds tend to have high lattice energies
High lattice energies make them hard and brittle
High lattice energies make them have relatively high melting points
Ionization energies increase rapidly for each successive electron removed

19. Transition Metals
Lose their valence level “s” electrons first and then one or two “d” orbital elctrons
Having two outermost “s” electons is the major reason why transition metals commonly form 2+ ions
For example:
Fe: (Ar) 3d64s2
Fe2+: (Ar) 3d64s0 or (Ar) 3d6
Fe3+: (Ar) 3d54s0 or (Ar) 3d5

20. Covalent Bonds
Formed when two atoms share one or more pairs of electrons
Sharing one: Single bond
Sharing two: Double bond
Sharing three: Triple bond
sharing pairs of electrons to attain octets
molecules generally weakly attracted to each other
observed physical properties of molecular substance due to these attractions

21. COVALENT BOND
Covalent bonds form when electrons are shared between two atoms.
Covalent bonds form between two non-metals.
Electrons shared

22. Single Covalent Bonds •• • •• • • • • •• •• •• •• •• •• F H O H F H O
two atoms share one pair of electrons
2 electrons
one atom may have more than one single bond F •• • •• H • • O • • H •• F •• •• H •• O •• H •• F

23. Double Covalent Bond •• • •• • ••
two atoms sharing two pairs of electrons
4 electrons
shorter and stronger than single bond O •• • O •• • O •• O

24. Triple Covalent Bond •• • •• • ••
two atoms sharing 3 pairs of electrons
6 electrons
shorter and stronger than single or double bond N •• • N •• • N •• N

25. POLAR & NON-POLAR BONDS
Electrons shared equally
Two types of covalent bonds exist:
Non-polar covalent bonds occur between similar atoms.
Polar
&
Nonpolar
In these bonds the electron pair is shared equally between the two protons.

26. POLAR & NON-POLAR BONDS
Polar covalent bonds occur between different atoms.
In these bonds the electron pair is shared unequally between the two atoms.
As a result there is a charge separation in the molecule, and partial charges on each atom.
 
H F

27. Dipole Moments
A dipole is a material with positively and negatively charged ends
Polar bonds or molecules have one end slightly positive, d+; and the other slightly negative, d-
not “full” charges
come from nonsymmetrical electron distribution
Dipole Moment, m, is a measure of the size of the polarity
measured in Debyes, D

28. Least electronegative
ELECTRONEGATIVITY
Linus Pauling derived a relative Electronegativity Scale based on Bond Energies.
Electronegativity (E.N.) is the ability of an atom involved in a covalent bond to attract the bonding electrons to itself.
F 4.0
Cs 0.7
Most electronegative
Least electronegative

29. ELECTRONEGATIVITY (THIS IS REVIEW)
Electronegativity increases

30. BOND POLARITY & ELECTRONEGATIVITY
Polarity is a measure of the inequality in the sharing of bonding electrons
The more different the electronegativity of the elements forming the bond
The larger the electronegativity difference (EN)
The more polar the bond formed

31. POLARITY & ELECTRONEGATIVITY
As difference in electronegativity increases
Bond polarity increases
Most polar
Least polar

32. POLARITY & ELECTRONEGATIVITY
SERIOUSLY NEED TO MEMORIZE THIS!!!
POLARITY & ELECTRONEGATIVITY
Electronegativity
difference
Bond Type
EN = 0
Non-polar covalent
0 < EN <1.7
Polar covalent
1.7 < EN
Ionic

33. The molecule is nonpolar covalent
POLARITY & ELECTRONEGATIVITY
EXAMPLES
The molecule is nonpolar covalent H
Hydrogen Molecule
Electronegativity
2.1
EN = 0

34. The molecule is polar covalent
POLARITY & ELECTRONEGATIVITY
EXAMPLES
The molecule is polar covalent H Cl
Hydrogen Chloride Molecule + -
Electronegativity
2.1
3.0
EN = 0.9

35. POLARITY & ELECTRONEGATIVITY
EXAMPLES
No molecule exists
The bond is ionic
Sodium Chloride
Na+
Cl-
Electronegativity
0.9
3.0
EN = 2.1

36. (similar electronegativities) (small to moderate EN)
SUMMARY OF BONDING
Ionic Bond
(large EN)
EN > 1.7
Non-polar
(similar electronegativities)
EN = 0
Polar
(moderate EN)
Covalent Bond
(small to moderate EN)
0 < EN < 1.7

37. Bonding & Lone Pair Electrons
Electrons that are shared by atoms are called bonding pairs
Electrons that are not shared by atoms but belong to a particular atom are called lone pairs
also known as nonbonding pairs

38. LEWIS STRUCTURES
In a Lewis structure, a shared electron pair is indicated by two dots between the atoms, or by a dash connecting them.
Unshared pairs of valence electrons (called lone pairs) are shown as belonging to individual atoms or ions.

39. LEWIS STRUCTURES
Covalent molecules are best represented with electron-dot or Lewis structures.
Structures must satisfy octet rule (8 electrons around each atom).
Hydrogen is one of the few exceptions and forms a doublet (2 electrons).

40. LEWIS STRUCTURES Non-bonding electrons must be displayed as dots.
Bonding electrons can be displayed by a dashed line.

41. Polyatomic Ions
The polyatomic ions are attracted to opposite ions by ionic bonds
Form crystal lattices
Atoms in the polyatomic ion are held together by covalent bonds

42. Lewis Formulas of Molecules
shows pattern of valence electron distribution in the molecule
useful for understanding the bonding in many compounds
allows us to predict shapes of molecules
allows us to predict properties of molecules and how they will interact together

43. LEWIS STRUCTURES
More complex Lewis structures can be drawn by following a stepwise method:
1. Count the number of electrons in the structure.
Draw a skeleton structure.
- most metallic element generally central
- halogens and hydrogen are generally
terminal
- many molecules tend to be symmetrical
- in oxyacids, the acid hydrogens are attached to an oxygen

44. LEWIS STRUCTURES
More complex Lewis structures can be drawn by following a stepwise method:
3. Connect atoms by bonds (dashes or dots).
4. Distribute electrons to achieve Octet rule.
5. Form multiple bonds if necessary.

45. Example 1: H O H Write Lewis structure for H2O Step 1: H2O
= 8 electrons
2 (1) + 6 = 8
Step 2:
 
H O H
Step 3:
Skeleton structure should be symmetrical
Hydrogen has doublet
4 electrons used 4 electrons remaining
Octet rule is satisfied
Step 4:

46. Example 2: O C O Write Lewis structure for CO2 Step 1: CO2
= 16 electrons
4 + 2(6) = 16
Step 2:
 
 
O C O  
Step 3:  
 
 
Skeleton structure should be symmetrical
Step 4:
Octet rule is NOT satisfied
10 electrons used 6 electrons remaining
4 electrons used 12 electrons remaining
Octet rule is satisfied
Step 5:

47. Writing Lewis Structures for Polyatomic Ions
the procedure is the same, the only difference is in counting the valence electrons
for polyatomic cations, take away one electron from the total for each positive charge
for polyatomic anions, add one electron to the total for each negative charge

48. Octet rule is satisfied Octet rule is NOT satisfied
Example 3:
Write Lewis structure for CO32-
Step 1:
CO32-
= 24 electrons
4+3(6)+2 = 24
Step 2:
6 electrons remaining
12 electrons remaining
0 electrons remaining
18 electrons remaining
 
 
O C O O
Step 3:
 
 
Step 4:
Step 5:
Octet rule is satisfied
Octet rule is NOT satisfied

49. Structure has 14 electrons Structure is incorrect
Example 4:
Determine if each of the following Lewis structures are correct or incorrect. If incorrect, rewrite the correct structure.
Structure has 14 electrons
Only 12 electrons shown
2(5) + 4(1) = 14 2 4 2
Structure is incorrect 2 2
Octets are complete

50. Exceptions to the Octet Rule
H & Li, lose one electron to form cation
Li now has electron configuration like He
H can also share or gain one electron to have configuration like He
Be shares 2 electrons to form two single bonds
B shares 3 electrons to form three single bonds
expanded octets for elements in Period 3 or below
using empty valence d orbitals
some molecules have odd numbers of electrons NO

51. Some molecules, such as SF6 and PCl5 have more than 8 electrons around a central atom in their Lewis structure.
SF6 and PCl5 can violate the octet rule through the use of empty d orbitals:
both S and P can utilize empty d orbitals to hold pairs of electrons that help
bond halogen atoms.

52. Resonance
we can often draw more than one valid Lewis structure for a molecule or ion
in other words, no one Lewis structure can adequately describe the actual structure of the molecule
the actual molecule will have some characteristics of all the valid Lewis structures we can draw

53. Resonance
Lewis structures often do not accurately represent the electron distribution in a molecule
Lewis structures imply that O3 has a single (147 pm) and double (121 pm) bond, but actual bond length is between, (128 pm)
Real molecule is a hybrid of all possible Lewis structures
Resonance stabilizes the molecule
maximum stabilization comes when resonance forms contribute equally to the hybrid

54. Resonance
we can often draw more than one valid Lewis structure for a molecule or ion
Real molecule is a hybrid
of all possible Lewis structures
represents resonance structures
The three oxygens are chemically equivalent, so it makes no
difference to the ion which oxygen assumes the double bond.

55. MOLECULAR SHAPES
A very simple model , VSEPR (Valence Shell Electron Pair Repulsion) Theory, has been developed by chemists to predict the shape of large molecules based on their Lewis structures.
The three-dimensional shape of the molecules is an important feature in understanding their properties and interactions.
All binary molecules have a linear shape since they only contain two atoms.
More complex molecules can have various shapes (linear, bent, etc.) and need to be predicted based on their Lewis structures.

56. MOLECULAR SHAPES
Based on VSEPR, the electron pair groups in a molecule will repel one another and seek to minimize their repulsion by arranging themselves around the central atom as far apart as possible.
Electron pair groups can be defined as any one of the following:
bonding pairs
non-bonding pairs
multiple bonds

57. SUMMARY OF VSEPR SHAPES
Number of electron pair groups around central atom
Molecular
Shape
Bond
Angle
Examples
Bonding
Non-bonding 2
Linear
180
CO2 3
Trigonal planar
120
BF3 1
Bent
SO2 4
Tetrahedral
109.5
CH4
Pyramidal
NH3
H2O

58. 2 electron pairs around the central atom
MOLECULAR SHAPES
Molecules with 2 electron pair groups around the central atom form a linear shape.
Linear molecules have polar bonds, but are usually non-polar.
Bond angle is 180
Shape is linear
2 electron pairs around the central atom

59. 3 electron pairs around the central atom
MOLECULAR SHAPES
Molecules with 3 electron pair groups around the central atom form a trigonal planar shape.
Trigonal planar molecules have polar bonds, but are usually non-polar.
Bond angle is 120
3 electron pairs around the central atom
Shape is trigonal planar

60. Bent molecules have polar bonds, and are polar.
MOLECULAR SHAPES
Molecules with 2 bonding pairs and 1 non-bonding pair groups around the central atom form a bent shape.
Bent molecules have polar bonds, and are polar.
1 Non-bonding pair
2 bonding pairs around the central atom
Shape is bent
Bond angle is 120

61. 4 bonding pairs around the central atom
MOLECULAR SHAPES
Molecules with 4 electron pairs groups around the central atom form a tetrahedral shape.
Tetrahedral molecules have polar bonds, and are usually non-polar.
4 bonding pairs around the central atom
Shape is tetrahedral
Bond angle is 109.5

62. 3 bonding pairs around the central atom
MOLECULAR SHAPES
Molecules with 3 bonding pairs and 1 non-bonding pair groups around the central atom form a pyramidal shape.
Pyramidal molecules have polar bonds, and are polar.
1 Non-bonding pair
Shape is pyramidal
3 bonding pairs around the central atom
Bond angle is 109.5

63. 2 bonding pairs around the central atom
MOLECULAR SHAPES
Molecules with 2 bonding pairs and 2 non-bonding pair groups around the central atom form a bent shape.
Bent molecules have polar bonds, and are polar.
2 Non-bonding pair
Shape is bent
2 bonding pairs around the central atom
Bond angle is 109.5

64. SUMMARY OF MOLECULAR SHAPES
Linear
Bent
Trigonal planar
Pyramidal
Tetrahedral
Symmetrical shapes
Polar bonds
Non-polar molecules
Unsymmetrical shapes
Polar bonds
Polar molecules

65. Polarity of Molecules For a molecule to be polar it must
1) have polar bonds, symmetrical shape, and
different terminal atoms
2) have polar bonds
electronegativity difference - theory
bond dipole moments – measured
3) have an unsymmetrical shape
using vector addition
polarity effects the intermolecular forces of attraction

66. and unsymmetrical shape causes molecule
Dipole moment is the measured polarity of a polar covalent bond.
It is defined as the magnitude of charge (electrons) on the atoms and the
distance between the two bonded atoms.
polar bonds,
and unsymmetrical shape causes molecule
to be polar
polar bonds,
but nonpolar molecule
because pulls cancel

67. CH2Cl2
m = 2.0 D
CCl4
m = 0.0 D

68. Adding Dipole Moments

69. COMPARING PROPERTIES OF IONIC & COVALENT COMPOUNDS
Structural Unit
Ions
Molecules
Melting Point
High
Low
Boiling Point
Solubility in H2O
Low or None
Electrical Cond.
None
Examples
NaCl, AgBr
H2, H2O