1. Objectives Chapter 6 Define chemical bond.
Introduction to Chemical Bonding
Chapter 6
Define chemical bond.
Explain why most atoms form chemical bonds.
Describe ionic and covalent bonding.
Explain why most chemical bonding is neither purely ionic nor purely covalent.
Classify bonding type according to electronegativity differences.

2. What is bonding?
Bonding is the “glue” that hold two or more elements together
This “glue” is most likely formed as a result of a chemical reaction
Bonding and molecular structure play a central role in determining the course of chemical reactions

3. What is a bond?
A bond can be thought of as a force that holds groups of two or more atoms together and makes them function as a unit
Example : water O
Bonds require energy to break
and release energy when made H H

4. Why do atoms bond?

5. Atoms bond because bonding lowers the atom’s potential energy.

7. Bond Energy is the energy required to break chemical bonds
-is an endothermic process
-requires energy
-endothermic process
-bond energy values are
postive

8. Bond Formation -is a energy releasing process -exothermic process
-the values for bond formation
are negative

9. Bond Length Is the distance between bonded atoms
Bond length and bond energy are indirectly related.

10. IONIC BONDING
Chemical bonding resulting from the electrical attraction between cations and anions.

11. Three Types of Bonds
Ionic
Covalent
Metallic

12. Visual Concepts
Chapter 6
Ionic Bonding

13. Ionic Bonds Na+ Cl- NaCl + Na and Cl The electrostatic interaction
Na is a metal and likes to lose one electron
Cl is a nonmetal and likes to gain one electron
the final ionic compounds is NaCl
Na+
Cl-
NaCl +
The electrostatic interaction
keeps them together!

14. Ionic Bonds Na looses an electron and chlorine gains it!
They do this to achieve an octet! Na Cl

15. Covalent Bonds Covalent Bonds exist between nonmetals bonded together
form when atoms of nonmetals share electrons
electrons can be shared equally or unequally

16. Types of Covalent
Polar
Nonpolar
Network Covalent

17. Visual Concepts
Chapter 6
Covalent Bonding

18. Metallic Bonds Metallic bonds exist between metals
Occur when two metals, usually the same metal, are bonded together

19. IONIC OR COVALENT?
How can you tell?

20. The type of bond can usually be calculated by finding the difference in electronegativity of the two atoms that are going together.

21. Using Electronegativity Difference to Classify Bonding
Visual Concepts
Chapter 6
Using Electronegativity Difference to Classify Bonding

22. Electronegativity Difference
If the difference in electronegativities is between:
1.7 to 4.0: Ionic
0.3 to 1.7: Polar Covalent
0.0 to 0.3: Non-Polar Covalent
Example: NaCl
Na = 0.8, Cl = 3.0
Difference is 2.2, so
this is an ionic bond!

23. Comparing Polar and Nonpolar Covalent Bonds
Visual Concepts
Chapter 6
Comparing Polar and Nonpolar Covalent Bonds

24. Ionic Vs. Covalent Bonding
Bonding and Ionic Compounds
Chapter 6
Ionic Vs. Covalent Bonding

25. Bonds Between Atoms Polyatomic Ions Ionic Covalent Metallic
Molecular Substance
Network Solids
Polar
Nonpolar
What are we going to learn about???

26. Marriage
Forming of a bond is like marriage
The breaking of a bond relates to a divorce.
More stable
exothermic
Divorce
Less stable
Endothermic

27. Ionic and Covalent Compounds
Compare and Contrast
Ionic and Covalent Compounds

28. PROPERTIES OF IONIC COMPOUNDS
Ionic compound – is composed of positive and negative ions that are equal in charge.
Formula Unit – is the simplest collection of atoms from which an ionic compound’s formula can be established.

29. Properties of Ionic Bonds
What is an Ionic Bond?
- An Ionic Bond is a chemical bond resulting from the TRANSFER of electrons from one bonding atom to another
When is an ionic bond formed?
- An ionic bond is formed when a cation (positive ion) transfers electrons to an anion (negative ion).

30. What are some characteristics of an ionic bond?
Exist as crystalline units at room temperatures
Brittle
Have higher melting points and boiling points compared to covalent compounds . Melting points are 1000 ˚C
Conduct electrical current in molten or solution state but not in the solid state
The smallest piece of an ionic compound is a formula unit.

31. What are some characteristics of an ionic bond?
6. A formula unit is the smallest collection of atoms from which a formula can be established.
Usually dissolves in polar solvents such as water. This is called dissociation.
The best way to test for an ionic compound is electrical conductivity in solution.
Formula Units
of salt

32. Dissolving of salt in water - Dissociation

33. Properties of Ionic Compounds
Lattice energy- is the energy released when one mole of an ionic crystalline compound is formed from gaseous ions.

34. Properties of Ionic Compounds
Lattice energy values are negative.
The more negative the value the stronger
the ionic bond.
Compound
Calculated Lattice Energy
NaCl
−756 kJ/mol
LiF
−1007 kJ/mol
CaCl2
−2170 kJ/mol

36. Covalent Bonds What is an Covalent Bond?
- A covalent bond is a chemical bond resulting from SHARING of electrons between 2 bonding atoms.
What forms a covalent bond?
- A covalent bond is formed between two nonmetals.

37. There are five different categories associated with covalent bonds
There are five different categories associated with covalent bonds. What are the 5 different categories?
Covalent
Network Solids
Molecular Substance
Polar
Coordinate Covalent
Nonpolar

38. What are some characteristics of a covalent bond?
Very strong
Low melting and boiling points
Below 500 ˚C
Exist as solids, liquids, or gases at room temperature
Do not conduct electricity unless it is a molecular electrolyte (acid or base)
Molecular electrolytes undergo ionization in polar solvents acids and bases.

39. Ionization of Molecular Acids in Water.

40. Melting and Boiling Points of Compounds
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Melting and Boiling Points of Compounds

41. Covalent Bonds can have multiple bonds, so you should be familiar with the following…
Single Covalent Bond- chemical bond resulting from sharing of an electron pair between two atoms.
Double Covalent Bond- chemical bond resulting from sharing of two electron pairs between two atoms.
Triple Covalent Bond-chemical bond resulting from sharing of three electron pairs between two atoms.
Triple bonds are the strongest and shortest

42. First, we are going to look at Polar Covalent…
What is polar covalent?
-Polar covalent is a description of a bond that has an uneven distribution of charge due to an unequal sharing of bonding electrons.
The boy is not equally sharing with anyone else but rather taking all the food for himself.

43. Next, we are going to look at Non-Polar Covalent…
What is non-polar covalent?
-Non polar covalent is a covalent bond that has an even distribution of charge due to an equal sharing of bonding electrons.
This couple is non- polar because they are sharing the drink equally between them.

44. Now, we are going to look at Network Solids…
Diamond and graphite are examples of network covalent compounds.
What is a Network Solid?
-A network solid is a solid that has covalently bonded atoms linked in one big network or one big macromolecule.
Name 3 Characteristics of a Network Solid.
Poor conductors of heat and electricity
Hard / Strong
High melting and boiling points

45. Just as a summary to what each bond looks like…

46. Properties of Metallic Bond
The Metallic-Bond Model
The chemical bonding that results from the attraction
between metal atoms and the surrounding sea of
electrons is called metallic bonding.

47. Metallic Bonding Name 4 Characteristics of a Metallic Bond.
What is a Metallic Bond?
- A metallic bond occurs in metals. A metal consists of positive ions surrounded by a “sea” of mobile electrons.
Good conductors of heat and electricity
Great strength
Malleable and Ductile
Luster
This shows what a metallic bond might look like.

48. Properties of Metals: Malleability and Ductility

49. Metallic Bonding
Malleability is the ability of a substance to be hammered or beaten into thin sheets.
Ductility is the ability of a substance to be drawn, pulled, or extruded through a small opening to produce a wire.

50. Metallic Bonding
Chemical bonding is different in metals than it is in ionic, molecular, or covalent-network compounds.
The unique characteristics of metallic bonding gives metals their characteristic properties, listed below.
electrical conductivity
thermal conductivity
malleability
ductility
shiny appearance

51. Why:
a. are metals malleable?
are metals ductile?
are metals shiny?

52. Intermolecular Forces
An attractive force that operates between molecules
There are 3 kinds of intermolecular forces:
London dispersion force
Dipole-dipole force
Hydrogen-bonding force

53. London Dispersion Forces
Instantaneous dipole
A temporary dipole formed when the electrons in an atom or nonpolar molecule happen to be more on one side in an instant in time, causing it to be more negative than normal and the opposite side positive
Induced dipole
Positive end of the dipole exerts an attractive force on nearby electrons, causing an adjacent atom to develop into another temporary dipole

54. London Dispersion Forces
The attraction between temporary dipoles
Occurs between atoms and molecules
Only intermolecular force in nonpolar substances
Tend to be stronger the larger the atom or molecule (the more electrons are in the atom or molecule)
Relatively weak forces

55. Dipole-Dipole Forces Attraction between polar molecules
Occurs when the partially positive end of one molecule attracts the partially negative end of another molecule
Generally stronger than London dispersion forces

56. Figure 5.8: Polar sulfur dioxide molecules attracted to one another (oxygen atoms attract electrons more than do sulfur atoms).
Fig. 5-8, p. 113

57. Hydrogen Bonding Special type of dipole-dipole force
Only occurs in molecules that contain hydrogen bonded to a small, highly electronegative element (N, O, F)
Stronger than a regular dipole-dipole force

58. Figure 5.10: The four hydrogen bonds between one water molecule and its neighbors.
Fig. 5-10, p. 114

59. Type of Force Type of Interaction Occurrence
London dispersion force
A temporary dipole in one molecule induces the formation of a temporary dipole in a nearby molecule and is attracted to it.
All atoms and molecules
Dipole-Dipole Force
Polar molecules (permanent dipoles) attract one another
Polar molecules
Hydrogen-Bonding Force
Two dipoles, one containing hydrogen to an electronegative element and the other containing an electronegative element, attract one another.
Polar molecules containing unpaired molecules and a hydrogen bonded to nitrogen, oxygen, or fluorine

60. Electron Dot

61. How to Draw Lewis Structures

62. .. linear 180o BeCl2 valence e- = 2 + (2 x 7) = 16e- fewer than 8e- Cl
two
valence pairs on Be
bonding e-
linear molecule

63. .. .. linear 180o CO2 valence e- = 4 + (2 x 6) = 16e- C O C O two
valence pairs on C
ignore double bonds
single and double bonds same
molecular geometry
linear
molecular shape
linear

64. .. : .. : .. : 120o SO2 valence e- = 6+ (2 x 6) = 18e- S O S O S O
three
valence pairs on S
two bonding pairs
one lone pair
molecular shape
bent
< 120o

65. : 109.5o NH3 valence e- = 5+ (3 x 1) = 8e- N H four valence pairs on N
three bonding pairs
one lone pair
molecular shape
trigonal pyramid
< 109.5o

66. tetrahedral
109.5o
CH4
valence e- = 4+
(4 x 1)
= 8e- C H
four
valence pairs on C
109.5o
molecular geometry
tetrahedral
molecular shape
tetrahedral

67. What Proof Exists for Hybridization?
We have studied electron configuration notation and
the sharing of electrons in the formation of covalent
bonds.
Lets look at a molecule of methane, CH4.
Methane is a simple natural gas. Its molecule has a
carbon atom at the center with four hydrogen atoms covalently bonded around it.

68. Carbon ground state configuration
What is the expected orbital notation of carbon
in its ground state?
Can you see a problem with this?
(Hint: How many unpaired electrons does this
carbon atom have available for bonding?)

69. Carbon’s Bonding Problem
You should conclude that carbon only has TWO electrons available for bonding. That is not not enough!
How does carbon overcome this problem so that
it may form four bonds?

70. Carbon’s Empty Orbital
The first thought that chemists had was that carbon promotes one of its 2s electrons…
…to the empty 2p orbital.

71. A Problem Arises… However, they quickly recognized a problem with such
an arrangement…
Three of the carbon-hydrogen bonds would involve
an electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom.

72. Unequal bond energy
This would mean that three of the bonds in a methane
molecule would be identical, because they would involve
electron pairs of equal energy.
But what about the fourth bond…?

73. Unequal bond energy The fourth bond is between a 2s electron from the
carbon and the lone 1s hydrogen electron.
Such a bond would have slightly less energy than the other bonds in a methane molecule.

74. Unequal bond energy
This bond would be slightly different in character than
the other three bonds in methane.
This difference would be measurable to a chemist
by determining the bond length and bond energy.
But is this what they observe?

75. Enter Hybridization: The simple answer is, “No”.
Measurements show that
all four bonds in methane
are equal. Thus, we need
a new explanation for the
bonding in methane.
Chemists have proposed an explanation – they call it
Hybridization.
Hybridization is the combining of two or more orbitals
of nearly equal energy within the same atom into
orbitals of equal energy.

76. In the case of methane, they call the hybridization
sp3, meaning that an s orbital is combined with three
p orbitals to create four equal hybrid orbitals.
These new orbitals have slightly MORE energy than
the 2s orbital…
… and slightly LESS energy than the 2p orbitals.
sp3 Hybrid Orbitals