Chemistry I Chapter 9

Covalent Bonding

Why do atoms bond?

1. Atoms bond to become more stable.2. A full valence shell is the most stable configuration3. In an ionic bond, atoms gain or lose electrons.4. In covalent bonds, atoms share electrons so that each atom has a full valence shell.

Why do atoms bond?

• Covalent bonds usually happen between element s that are close to each other on the periodic table.

Why do atoms bond?

• Covalent bonds usually happen between element s that are close to each other on the periodic table.

• Most covalent bonds are between nonmetals.

Why do atoms bond?

• Covalent bonds usually happen between element s that are close to each other on the periodic table.

• Most covalent bonds are between nonmetals.

• A molecule is formed when two or more atoms bond covalently.

Why do atoms bond?

5. Covalent bonds usually happen between element s that are close to each other on the periodic table.6. Most covalent bonds are between nonmetals. 7. A molecule is formed when two or more atoms bond covalently.

Why do atoms bond?

8. There are 7 elements that always exist as diatomic molecules. Diatomic molecules do not exist as single atoms but as molecules of the atoms bonded together. 9. The 7 diatomic molecules are • H2 O2 N2

• F2 Cl2 Br2 I2

10. Draw the lewis structure of the 7 diatomic molecules.

Why do atoms bond?

11.These seven molecules exist as diatomic molecules because by sharing electrons each can have a full valence shell.

Why do atoms bond?

• The electrons that are shared or the shared pairs are called the bonding electrons.

Why do atoms bond?

12. The electrons that are shared or the shared pairs are called the bonding electrons.13.The electrons that belong to individual atoms are called unshared pairs.

14. Draw Lewis structures for the following molecules:

• H2O

• NH3

• CH4

• HBr• PH3

• CCl4

• SiH4

• HCl

Why do atoms bond?

15.Unshared pairs are shown with two dots. Shared pairs are shown with a dash.16.Usually the element furthest to the left on the periodic table is in the middle.

Why do atoms bond?

17. Hydrogen is always on an end because it only makes one bond.18.Single covalent bonds are formed when a single pair of electrons are shared.

Why do atoms bond?

19.Single bonds are also called sigma bonds. The greek letter sigma is written σ.20. In a sigma bond, the valence orbitals overlap each other. Sigma bonds happen with s and p orbitals.

Multiple Bonds

• 21. Double and triple bonds are multiple covalent bonds.

22. A double bond occurs whenever atoms need to share two pairs of electrons so that each atom has a full valence shell.

• 23. Draw the Lewis structures for the following molecules:

• O2

• C2H4

• CO2

Multiple Bonds

• 24. When atoms have multiple bonds a pi bond forms (π bonds). The pi bond is found above and below the line where the atoms are connected.

Multiple Bonds

25.A double bond has one sigma bond and one pi bond.26.A triple bond has one sigma bond and two pi bonds.

Strength of Covalent bonds

27. The strength of the bonds depends on how close the atoms nuclei are. 28.The shorter the bond the stronger the bond.

Strength of Covalent bonds

29.As the number of shared pairs increases the length decreases.

Strength of Covalent bonds

• Single bonds are the longest• Double bonds are in the middle• Triple bonds are the shortest.

Strength of Covalent bonds

30.The amount of energy needed to break a specific covalent bond is called the bond dissociation energy.

Strength of Covalent bonds

31.Bond dissociation energy is always positive.

Strength of Covalent bonds

• For exampleMolecule bond length dissociation energy

F21.43 X 10-10 m 159 kJ/molO2 1.21 x 10-10 m 498 KJ/molN2 1.10 X 10-10 m 945 KJ/mol

Multiple bonds

32.The chemical potential energy of a molecule is the sum of the bond dissociation energies for all bonds in a molecule

Multiple bonds

33.Endothermic reactions occur when it takes more energy to break the bond than is released when the new bonds form.• Endothermic – energy is

absorbed.

Multiple bonds

34.Exothermic reactions occur when more energy is released when the new bond forms than was needed to break the original bonds. • Exothermic – energy is released.

Section 2 Naming Binary Molecular Compounds.

35. Binary molecular compounds contain only two different elements both of which are nonmetals.

Section 2 36.Naming Binary Molecular Compounds.

• Naming rules1. The first element is named using

the entire first element name2.The second element is named

using the root word with the suffix – ide added.

Section 2 36.Naming Binary Molecular Compounds.

• Naming rules3. Prefixes are used to indicate the

number of atoms of each element present. The exception is that the first element in the formula never uses the prefix mono.

Section 2 37.Naming Binary Molecular Compounds.

• # ATOMS/ PREFIX• 1 - MONO

2 - DI 3 - Tri 4 - Tetra 5 - Penta

Section 2 Naming Binary Molecular Compounds.

• # ATOMS/ PREFIX• 6 - Hexa• 7 - Hepta• 8 - Octa• 9 -Nona• 10 - Deca

38. Name the following molecular compounds

• P2O5

• CCl4

• As2O3

• CO• SO2

• NF3

39. Common names of molecular compounds

H2O water Dihydrogen MonoxideNH3 ammonia Nitrogen TrihydrideN2O laughing gas Dinitrogen monoxide

Naming Acids40. All acids have hydrogen as the cation.41. Acids are molecules that separate into ions in water.42. The two types of acids are binary and oxyacids.

Binary Acids

43. All binary acids have hydrogen and one other element

ExamplesHClH2S

Naming Binary Acids44. Use the prefix hydro, the root name with the ending changed to –ic• HCl _____________• HBr _____________• HF ______________• H3P _____________

Naming Binary Acids• All acids that have a polyatomic ion that

does not contain oxygen will use the above rules• HCN _____________

Naming Oxyacids45. Any acid that contains hydrogen and an oxyanion is an oxyacid.

Rules for naming oxyacids46.The name of the acid has the name of the polyatomic ion, a suffix and the word acid. • If the polyatomic ion ends in –ate , change

the ending to –ic.• If the polyatomic ion ends in – ite, change

the ending to -ous

Rules for naming oxyacids• H2SO4 _______________• H2SO3 _______________• HNO3 _______________• HNO2 _______________

Name the following acids ( some are binary and some are oxyacid)

• HI _______________• HClO3 _______________• HClO2_______________• H2SO4 _______________• H2S __________________

47. Flow Diagram for all molecular compounds

Section 3Molecular Structural Models

48. There are five different kinds of molecular structures use to show the 3 dimensional arrangement of molecules.

Section 3Molecular Structural Models

• A.) molecular formula• The symbols can be arranged to show

how the atoms are grouped.• i.e., HC2H3O2

Section 3Molecular Structural Models

• B.) Structural formulas-• Uses dashes but does not show

unshared pairs of electrons.

Section 3Molecular Structural Models

• C.) Lewis Structures- • Uses dashes to show bonds and dots to

show unshared pairs.

Section 3Molecular Structural Models

• D.) Space filling- different elements are represented by different colors. Atoms are stuck together like marshmallows glued together.

Section 3Molecular Structural Models

• E.) Ball and stick- • Balls represent the atoms and the sticks

represent the bonds.

Lewis structures for polyatomic ions

49.The technique is the same except that the charge of the ion has to be accounted for.

For example:PO4

-3

Lewis structures for polyatomic ions

The technique is the same except that the charge of the ion has to be accounted for.

For example:NH4

+

Lewis structures for polyatomic ions

The technique is the same except that the charge of the ion has to be accounted for.

For example:ClO4

-

Resonance Structures50. Occurs when there is more than one

valid Lewis Structure.51.These happen when there is a double

bond or triple bond.52.Resonance structures only differ in the

position of the electron pairs not the atom position.

Resonance StructuresNO3

-

Resonance StructuresSO3

-2

Resonance StructuresSO2

Exceptions to the octet rule55. Some molecules/ ions do not obey the octet rule

Exceptions to the octet rule• Exceptions• 1.) odd number of electrons

(NO2)

Exceptions to the octet rule• Exceptions• 2.) The central atom may have

less than an octet if the central atom has fewer than 4 valence electrons. ( BF3)

Exceptions to the octet rule• Exceptions• 3.) More than an octet.• (PCl 5)

56. The shape of a molecule determines many of its physical and chemical properties.

VSEPR Model

57. The model used to determine the molecular shape is the valence shell electron pair repulsion model (VSEPR model)

58. The VSEPR theory states that in a small molecule the pairs of valence electrons are arranged as far apart from each other as possible.

59. The angle formed by any two terminal atoms and the central atom is a bond angle.

60. Both the shared and unshared pairs of electrons are important in determining molecular shape.

61. Because lone pairs are not shared between two nuclei, they occupy a slightly larger orbital than shared electrons.

Hybridization

62. Atomic orbitals undergo hybridization during bonding.63. Hybridization is a process in which atomic orbitals are mixed to form new,identical hybrid orbitals.

64. Notice that the number of atomic orbitals mixed to form the hybrid orbitals equals the total number of pairs of electrons.

65.Lone pairs of electrons also occupy hybrid orbitals.

Determine the molecular geometry, bond angle and type of hybrid orbital for the following.

BF3

Determine the molecular geometry, bond angle and type of hybrid orbital for the following.

OCl2

Determine the molecular geometry, bond angle and type of hybrid orbital for the following.

NH4+

Determine the molecular geometry, bond angle and type of hybrid orbital for the following.

BeF2

Determine the molecular geometry, bond angle and type of hybrid orbital for the following.

CF4

Electronegativity and Polarity

66. Electronegativity is a measure of the ability of an atom to attract electrons in a chemical bond.

67. The type of chemical bond between atoms can be predicted using the electronegativity difference of the elements that are bonded.

68. When the electronegativity (EN) difference is less than or equal to .4 the bond is non-polar covalent.

69. When the EN difference is between .4 and 1.8 the bond is polar covalent.

70. When the EN difference is greater than or equal to 1.8 the bond is ionic.

71. In non-polar bonds the electrons are being shared equally.72. In a polar covalent bond the electrons spend more time around the more electronegative atom than the less electronegative atom.

73. The side of the molecule with a larger electronegative value will have a partial negative charge (δ-)74. The side of the molecule with a lower electronegative value will have a partial positive charge (δ+)

• For example HCl

76. Molecules are either nonpolar or polar.

77. Properties of Polar Molecules

• Tend to align in electric fields.• Tend to be soluable in other polar

substance ( Nonpolar substances tend to be soluble in nonpolar solvents; polar substances tend to dissolve in polar solvents)

78. Polar molecules are also called dipoles.

Some molecules contain polar bonds but are not polar molecules

• Compare H2O and CCl4

80. The attraction between individual molecules are known as intermolecular forces or Van Der Waals forces.81. For nonpolar substances the attraction between molecules is weak and is called dispersion forces. (induced dipole)

82. The force between polar molecules is stronger and is called dipole – dipole forces.

83. The more polar the bond the stronger the dipole-dipole force.84. The hydrogen bond is formed between the hydrogen end of one dipole and the fluorine, oxygen or nitrogen end of another dipole.