1. Covalent Bonding
Chapter 8

2. Covalent bonds form when atoms share electrons!!!
Chapter Main Idea
Covalent bonds form when atoms share electrons!!!

3. Section 1:The Covalent Bond

4. Section 1: Essential Questions & Vocabulary
How does the octet rule apply to atoms that form covalent bonds?
Why do atoms form single, double, and triple covalent bonds?
What are sigma and pi bonds and how do they contrast?
How are the strength of a covalent bond, its bond length, and its bond dissociation energy related?
Vocabulary
Chemical bond
Sigma bond
Covalent bond
Pi bond
Molecule
Endothermic reaction
Lewis structure
Exothermic reaction

5. Section 1: Main Idea
Atoms gain stability when they share electrons and form covalent bonds.

6. Why do atoms bond?
The stability of an atom, ion or compound is related to its energy:
lower energy states are more stable.
Elements can gain stability in two ways:
Transferring electrons – elements gain or lose electrons to form ions with noble-gas configurations. (Ionic Bonds)
Sharing electrons – elements share valence electrons with other atoms resulting in noble-gas configurations. (Covalent Bonds)

7. What is a covalent bond?
Covalent bond - chemical bond that results from sharing electrons
Molecule is formed when two or more atoms bond covalently.
Most covalent bonds form between nonmetal - nonmetal elements.
Diatomic molecules (H2, N2, F2, O2, I2, Cl2, Br2) exist because the two-atom molecules are more stable than the individual atoms.

8. Each chlorine atom wants to gain one electron to have a full octet.
Formation of Cl2 Cl
Each chlorine atom wants to gain one electron to have a full octet.

9. Chlorine atoms get close together until they share electrons.
Formation of Cl2 Cl Cl
Chlorine atoms get close together until they share electrons. Cl
Full Octet
Full Octet
The octet is achieved by each atom sharing the electron pair in the middle.

10. Single bonds are abbreviated
Chlorine Molecule Cl Cl Cl
Single bonds are abbreviated
with a dash
This is the bonding pair or a shared pair of electrons called a single bond.

11. What is a covalent bond?
The most stable arrangement of atoms exists at the point of maximum net attraction, where the atoms bond covalently and form a molecule.

12. Single Covalent Bonds
single covalent bond - one pair of electrons is shared between two atoms
Example: two hydrogen atoms forming a hydrogen molecule with a single covalent bond

13. Single Covalent Bonds by Group
Lewis structure – 2 dots or a line are used to symbolize a single covalent bond.
Group 14
Group 15
Group 16
Group 17
Each element has 4 valence electrons
Will form 4 single bonds with other non-metals
Each element has 5 valence electrons
Will form 3 single bond with other non-metals
Each element has 6 valence electrons
Will form 2 single bond with other non-metals
Each element has 7 valence electrons
Will form 1 single bond with other non-metals

14. Sigma Bonds Sigma bonds - single covalent bonds.
Sigma bonds occur when the pair of shared electrons is in an area centered between the two atoms.

15. Double Bonds & Triple Bonds
Double Bond – two pairs of shared electrons
1 sigma bond, 1 pi bond
Triple Bond – three pairs of shared electrons
1 sigma bond, 2 pi bonds

16. Formation of Oxygen Molecule
Oxygen atoms are highly electronegative. So both atoms want to gain two electrons.
Each atom has two unpaired electrons

17. Formation of Oxygen Molecule
Both electron pairs are shared.
6 valence electrons
plus 2 shared electrons
= full octet

18. Formation of Oxygen Molecule=
two bonding pairs, or two pairs of shared electrons called a double bond.
For convenience, the double bond
can be shown as two dashes.

19. Nitrogen Molecule

20. Lewis Dot Structures: Practice
PH3
CCl4
H2S
SiH4
HCl

21. Strength of Covalent Bonds
Bond Length – distance between the two bonded nuclei
Single Bonds are the longest and weakest bonds.
Double bonds are shorter but stronger than single bonds.
Triple bonds are shorter than double but stronger than double bonds.

22. Bond Length & Bond Strength

23. Bonds and Energy Exothermic
Endothermic Reaction – greater amount of energy is required to break the existing bonds in the reactants than is released when the new bonds form in the products.
Energy goes into the reaction.
Exothermic Reaction – more energy is released during product bond formation than is required to break the bonds in the reactants.
Energy is released from the reaction.

24. Section 2: Naming Molecules

25. Essential Questions and Vocabulary
What rules do you follow to name a binary molecular compound from its molecular formula?
How are acidic solutions named?
Vocabulary
Oxyanion
Oxyacid

26. Section 2: Main Idea
Specific rules are used when naming binary molecular compounds, binary acids, and oxyacids.

27. Naming Binary Molecular Compounds
Binary Molecule – Compound composed of only two nonmetal atoms
Use of prefixes to indicate the number of each atom in the molecule.

28. Prefixes in Covalent Compounds
Exception: The first element NEVER uses the prefix mono-

30. Naming Binary Molecules - RULES
1- The name of the first element does not get changed, only a prefix is added if it is more than one.
Ex: Trisulfur tetrafluoride S3F4
2- The second element gets a different ending, as well as a prefix.
Ex : phosphorus trichloride PCl3
(NOT: phosphorus trichlorine)

31. Naming Binary Molecules – Rules cont.
3- If the first element in the compound is single you do not write mono:
ex. CO2 - carbon dioxide
(NOT monocarbon dioxide)
4- If the element’s name starts with a vowel, and the prefix ends with a vowel, you can drop the last vowel on the prefix:
ex. S5O10 - pentasulfur decoxide
(NOT pentasulfur decaoxide)

32. Naming Binary Molecules – Rules cont.
5- Covalent compounds do not get simplified. Name the compound exactly as you see it:
Ex. H2F2 = dihydrogen difluoride
(NOT hydrogen monofluoride HF)

33. Naming Binary Molecules: Examples
What is the name of SO3?
STEP 1 : The first nonmetal is S sulfur.
STEP 2 : The second nonmetal is O, named oxide.
STEP 3 : The subscript 3 of O is shown as the prefix tri.
SO3 → sulfur trioxide
The subscript 1(for S) or mono is understood.

34. Naming Binary Molecules: Examples
Name P4S3
STEP 1 : The first nonmetal P is phosphorus.
STEP 2 : The second nonmetal S is sulfide.
STEP 3 : The subscript 4 of P is shown as tetra.
The subscript 3 of O is shown as tri.
P4S3 → tetraphosphorus trisulfide

35. Names of Binary Molecules
Remember Prefixes are used
in the names of covalent compounds
because two nonmetals can form two or more different compounds
Examples of compounds of N and O:
NO nitrogen oxide
NO2 nitrogen dioxide
N2O dinitrogen oxide
N2O4 dinitrogen tetroxide
N2O5 dinitrogen pentoxide

36. Naming Binary Molecules: Practice
Select the correct name for each compound.
A. SiCl4 1) silicon chloride
2) tetrasilicon chloride
3) silicon tetrachloride
B. P2O5 1) phosphorus oxide
2) phosphorus pentoxide
3) diphosphorus pentoxide
C. Cl2O7 1) dichlorine heptoxide
2) dichlorine oxide
3) chlorine heptoxide

37. Naming Binary Molecules: Practice Solutions
A. SiCl4 3) silicon tetrachloride
B. P2O5 3) diphosphorus pentoxide
C. Cl2O7 1) dichlorine heptoxide

38. Naming Binary Molecules: Practice
Write the correct formula for each of the following:
A. phosphorus pentachloride
B. dinitrogen trioxide
C. sulfur hexafluoride

39. Naming Binary Molecules: Practice Solutions
A. phosphorus pentachloride
1P penta = 5Cl PCl5
B. dinitrogen trioxide
di = 2N tri = 3 O N2O3
C. sulfur hexafluoride
1S hexa = 6F SF6

40. Naming Binary Molecules: Practice
Write the name of each covalent compound.
CO _____________________
CO2 _____________________
PCl3 _____________________
CCl4 _____________________
N2O _____________________

41. Naming Binary Molecules: Practice Solutions
CO carbon monoxide
CO2 carbon dioxide
PCl3 phosphorus trichloride
CCl4 carbon tetrachloride
N2O dinitrogen oxide

42. TRY THESE:
e) SiO2
f) N2S5
g) P4O10
h) BrCl6
Diphosphorus pentoxide
Nitrogen monoxide
Xenon hexafluoride
Tetrarsenic decoxide

43. Naming Acids Acids produce Hydrogen ions (H1+) ions in solution.
Two common types of acids: Binary Acids & Oxyacids
Binary Acids - An acid that contains hydrogen and one other element
Oxyacids – acid that contains hydrogen and an oxyanion.
Oxyanions – polyatomic ion containing one or more oxygens.

44. Naming Binary Acids Rules
The first word
Begins with the prefix Hydro- to name the hydrogen part of the acid compound.
Rest of the word consists of the root of the second element and the suffix –ic.
The second word is acid.
Example:
HCl - Hydrochloric acid
HF - Hydrofluoric acid

45. Naming Oxyacids Rules
Name of the oxyanion with the changes described below.
Identify the oxyanion present.
If the oxyanion ends in –ate, change the ending to –ic.
If the oxyanion ends in – ite, change the ending to –ous.
The second word is acid.
Examples:
HNO3  Oxyanion nitrate (NO3) becomes nitric acid.
H2SO3  Oxyanion sulfite (SO3) becomes sulfurous acid.
NO3– the nitrate ion, becomes nitric.
The second word of the name is always acid, HNO3 (hydrogen and nitrogen ion) becomes nitric acid.

46. Naming Oxyacids Examples

47. Writing Formulas from molecular names.
The name of a molecular compound reveals its composition.
The flowchart can help you determine the name of a molecular covalent compound.

48. Section 3: Molecular Structures

49. Essential Questions and Vocabulary
What are the basic steps used to draw Lewis structures?
Why does resonance occur, and what are some resonant structures?
Which molecules are exceptions to the octet rule, and why do these exceptions occur?
Vocabulary
Ionic bond
Coordinate covalent bond
Structural formula
Resonance

50. Section 3: Main Idea
Structural formulas show the relative positions of atoms within a molecule.

51. Structural Formulas
Structural Formulas – use of letter symbols and bonds to show relative positions of atoms.

52. Lewis Structures Gilbert Lewis
Lewis developed the idea of the octet and coined the term Lewis dot structures
He was nominated 35 times for the Nobel Prize in chemistry, but never won.

53. Drawing Lewis Structures - Guidelines
Predict the location of each atom
The atom that has the least attraction for shared electrons will be the central atom
All other atoms become terminal atoms. Note: Hydrogen is always a terminal atom.
Determine the number of electrons available for bonding.
Total number of valence electrons for all atoms in molecule.

54. Drawing Lewis Structures - Guidelines
Determine the number of bonding pairs.
Divide the total number of electrons available for bonding by two.
Place the bonding pairs.
Place one bonding pair (single bond) between the central atom and each of the terminal atoms.

55. Drawing Lewis Structures - Guidelines
Determine the number of electron pairs remaining.
Subtract the number of pairs used in step 4 from total number of bonding pairs in step 3.
Determine whether the central atom satisfies the octet rule.
Central atom should be surrounded by 4 electron pairs.
If not, convert one or two lone pairs on the terminal atoms to form a double bond or triple bond.

56. Lewis Structure for Ammonia (NH3)
Nitrogen is the central atom and 3 terminal hydrogen atoms.
Total of 8 valence electrons available for bonding.
1 N atom × 5 valence electrons 1 N atom + 3 H atoms × 1 valence electron 1 H atom = 8 valence electrons
4 bonding pairs.
8 electrons 2 electrons/pair = 4 pairs
4. Place a single bond between the central atom and the terminal atoms.

57. Lewis Structure for Ammonia (NH3)
5. Number of electron pairs remaining. 6. The remaining lone pair must go on the central atom because hydrogen can only have one bond.
4 pairs total − 3 pairs used = 1 pair available

58. Lewis Structure for Carbon Dioxide (CO2)
Carbon is the central atom and 2 terminal oxygen atoms.
Total of 16 valence electrons available for bonding.
1 C atom × 4 valence electrons 1 C atom + 2 O atoms × 6 valence electrons 1 O atom = 16 valence electrons
8 bonding pairs.
16 electrons 2 electrons/pair = 8 pairs
4. Place a single bond between the central atom and the terminal atoms.

59. Lewis Structure for Carbon Dioxide (CO2)
5. Number of electron pairs remaining.
6. Add three lone pairs to each terminal oxygen atom.
Carbon atom does not meet the octet. Subtract a lone pair from each oxygen to form a double bond with oxygen.
8 pairs total − 2 pairs used = 6 pair available

60. Resonance Structures
Resonance – a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion.
This figure shows three correct ways to draw the structure for nitrate ion (NO3)-1.

61. Resonance Structures
The molecule behaves as though it has only one structure.
The bond lengths are identical to each other and intermediate between single and double covalent bonds.

62. Exceptions to the Octet Rule
Suboctet - A few compounds form stable configurations with less than 8 electrons around the atom
Coordinate covalent bond - forms when one atom donates both of the electrons to be shared with an atom or ion that needs two electrons.

63. Exceptions to the Octet Rule
Expanded Octet – Occurs when central atom has more than eight valence electrons
Can occur in period 3 or higher because of the presence of a d- orbital which allows for more than four covalent bonds.

64. Lewis Structures Practice & Solutions
NH3
H2O2
CH4
PCl3

65. Section 4: Molecular Shapes

66. Essential Questions & Vocabulary
What is VSEPR bonding theory?
How can you use the VSEPR model to predict the shape of, and the bond angles in a molecule?
What is hybridization?
Vocabulary
Atomic Orbitals
Hybridization
VSEPR model

67. The VSEPR model is used to determine molecular shape.
Section 4: Main Idea
The VSEPR model is used to determine molecular shape.

68. VSEPR Theory Valence-Shell-Electron-Pair-Repulsion theory
This theory helps us understand the 3D structure of molecules and their properties.
Bonding and unshared pairs of valence electrons become very important to uswithin VSEPR theory!
The shapes of molecules are determined because electron pairs want to be far apart from each other (repulsion).

69. AXE Notation AXE – Method to represent compounds
A represents the central atom X represents the bonding atoms E represents the lone pairs on the central atom

70. Linear Draw or build CO2 (AX2) Meaning 1 central atom, 2 bonded atoms
It has a linear shape
No lone pairs
A bond angle of 180o
Bonding pairs are far apart from each other

71. Trigonal Planar Draw or build BF3 This has a trigonal planar AX3
Meaning 1 central, 3 bonded
Bond angle: 120o
Bonds pointing to the corners of a triangle

72. Tetrahedral Draw or build CH4 This has a tetrahedral AX4
Meaning 1 central, 4 bonded
Bond angle: 109.5o

73. VSEPR Practice Use VSEPR theory to predict the shape of:
1. carbon tetrachloride, CCl4
2. carbon tetrabromide, CBr4
3. dichloromethane, CH2Cl2

74. Bent Draw or build H2O This has a bent shape AX2E2
Meaning 1 central, 2 bonded, 2 lone pairs
Bond angle: 105o
Similar angles to the tetrahedral bond angles

75. Trigonal pyramidal Draw or build NH3 This has a trigonal pyramidal
AX3E
Meaning 1 central, 3 bonded, 1 lone pair
Bond angle: 107o
Similar angles to the tetrahedral bond angles

76. VSEPR Theory – Practice & Solutions
BeCl2
OF2
AlCl3
PCl3
CF4
Linear
Bent
Trigonal Planar
Trigonal pyramidal
Tetrahedral

77. Hybridization
Hybridization – a process in which atomic orbitals mix and form new, identical hybrid orbitals.
Carbon often undergoes hybridization, which forms an sp3 orbital formed from one s orbital and three p orbitals.
Lone pairs also occupy hybrid orbitals.
Single, double, and triple bonds occupy only one hybrid orbital (CO2 with two double bonds forms an sp hybrid orbital).

78. Section 5: Electronegativity and Polarity

79. Essential Questions & Vocabulary
How is electronegativity used to determine bond type?
How do polar and nonpolar covalent bonds and polar and nonpolar molecules compare?
What are the characteristics of covalently bonded compounds?
Vocabulary
Electronegativity
Polar covalent bond

80. Section 5: Main Idea
A chemical bond’s character is related to each atom’s attraction for the electrons in the bond.

81. Electronegativity
Electronegativity – the ability of an atom to attract electrons in a chemical bond.

82. Electronegativity and Bond Character
Bond character can be determined based on the difference in electronegativity.

84. Polar Covalent Bond
Polar Covalent Bond – bond with unequal sharing of electrons
Bonding is often not clearly ionic or covalent.

85. Polar Covalent Bonds
Polar covalent bonds form when atoms pull on electrons in a molecule unequally.
Dipole – partial charges at the ends of the bond resulting from electrons spending more time around one atom than another

86. Polarity Practice & Solutions
N-H
F-F
Ca- Cl
Al- Cl
0.9 slightly polar
0 Nonpolar
2.0 ionic
1.5 very polar

87. Molecular Polarity
Covalently bonded molecules can be either polar or non- polar depending on the nature of the bonds and molecular geometry.
Non-polar molecules are not attracted by an electric field.
Polar molecules align with an electric field.

88. Polarity of Water Molecule
Water has two polar covalent bonds.
Water is polar molecule because its bent shape is asymmetrical.

89. Polar and Non-Polar Molecules
Molecules with nonpolar bonds are nonpolar.
Molecules with polar bonds can be both polar and non-polar depending on their geometry.
Note: If bonds are polar, asymmetrical molecules are polar and symmetrical molecules are nonpolar.
If the molecule is symmetrical, then the molecule is non-polar.
If the molecule is asymmetrical, then the molecule is polar.
Note: If bonds are polar, asymmetrical molecules are polar and symmetrical molecules are nonpolar.

90. Properties of molecular compounds
Composed of nonmetals
Can be a solid, liquid, or gas at room temperature
Low melting point and boiling points (compared to ionic compounds)
Poor to non-conductors of heat and electricity

92. Sodium chloride vs sugar
C12H22O11
NaCl
(Covalent Molecule)
(Ionic Compound)
Sodium is a metal
Chloride is a nonmetal
Melting point ~800oC
C, H, and O are nonmetals
Melting point ~185oC

93. Properties of Covalent Bonds: Intermolecular Forces
Covalent bonds between atoms are strong, but attraction forces between molecules are weak.
Van der Waals Forces – the weak interaction forces between molecules
Dipole-Dipole Force – the force between two oppositely charged ends of two polar molecules
Hydrogen bond - an especially strong dipole-dipole force between a hydrogen end of one dipole and a fluorine, oxygen, or nitrogen atom on another dipole.

94. Hydrogen Bonding