1. Chapter 9
Covalent Bonding

2. Why do Atoms Bond? Ionic Bond Get to a stable electron configuration
8 Valence Electrons
Called the “Noble Gas” configuration
Ionic Bond
One atom takes another atom gives electron(s) so each can be at 8 valence electrons
Covalent Bond
Atoms share electrons so that they are all at 8 valence electrons

3. What is a Covalent Bond? Covalent Bond:
The chemical bond that results from the sharing of electrons
The shared electrons are considered to be part of both energy levels of the elements that are bonded and sharing electrons
Occurs when elements are close to each other on the periodic table
Majority form between non-metallic elements

4. Formation of a Covalent Bond
Diatomic Molecules
Two of the same monatomic molecules sharing single or pairs of electrons to get to the stable noble gas configuration
Occur in nature because they are more stable than individual atoms F2 N2 O2
Cl2
Br2 I2

5. Formation of a Covalent Bond
Diatomic Molecules

6. Single Covalent Bonds
A covalent bond in which a single electron is shared
Bonding Pair = pair of shared electrons
Lewis Dot Structure
A line is DRAWN to represent the bonding pair
Group 7A = forms 1 covalent bond
Group 6A = forms 2 covalent bonds
Group 5A = forms 3 covalent bonds
Group 4A = forms 4 covalent bonds
Hydrogen = Forms 1 covalent bond

7. Single Covalent Bonds

8. Single Covalent Bonds Sigma Bond Sigma Bond AKA Single Covalent Bonds
Symbolized by the Greek Letter σ
Occurs when electron pair is centered between two atoms
If the electron orbitals overlap end-to-end
Concentrates the electrons in a bonding orbital between two atoms

9. Multiple Covalent Bonds
When atoms share more than 1 electron to attain noble gas configuration

10. Think-Pair-Share
Take 30 seconds to think about which diatomic molecules might form double covalent bonds.
Write these on your whiteboard/expo sheet
Take 30 seconds to discuss your finding with your partner. Add/subtract and change your molecules, as needed
Take 30 seconds and be prepared to share with the class

11. Multiple Covalent Bonds

12. Multiple Covalent Bonds
Pi Bond
A pi-bond is formed when parallel orbitals overlap to share electrons
Denoted by П
Double Covalent Bond has one sigma and one pi bond
Triple covalent bond has one sigma and two pi bonds

13. Multiple Covalent Bonds
Pi Bond

14. Strength of Covalent Bonds
Depends on how much distance separates the bonded nuclei
More shared electrons = stronger covalent bonds
Bond length decreases
Endothermic Reaction = occur when a greater amount of energy is required to break existing bonds in the reactants than released when the new bonds formed
Exothermic Reaction = occur when a more energy is released forming new bonds than in released in breaking apart the reactants

15. Naming Binary Molecular Compounds
The first element in the formula is always named first, using the entire element name
The second element in the formula is named using the root of the element and adding the suffix “-ide”
Prefixes are used to indicate the number of atoms of each type that are present within the compound
The first element will never use the word “Mono”

16. Covalent Bond Prefixes
Number of Atoms
Prefix 1
Mono- 2
Di- 3
Tri- 4
Tetra- 5
Penta- 6
Hexa- 7
Hepta- 8
Octa- 9
Nona- 10
Deca-

17. Name the Following Compounds
CCl4
As2O3 CO
SO2
NF3

18. Name the Following Compounds
CCl4 – Carbon Tetrachloride
As2O3 – Diarsenic Trioxide
CO – Carbon Monoxide
SO2 – Sulfur Dioxide
NF3 – Nitrogen Trifluoride

19. Common Names of Some Molecular Compounds
Formula
Common Name
Common Compound Name
H2O
Water
Dihydrogen monoxide
NH3
Ammonia
Nitrogen trihydride
N2H4
Hydrazine
Dinitrogen tetrahydride
N2O
Nitrous Oxide (Laughing Gas)
Dinitrogen monoxide NO
Nitric Oxide
Nitrogen monoxide

20. Naming Binary Acids
Binary acids contains hydrogen and one other element
Use the prefix “hydro-” to name the hydrogen part of the compound
The rest of the name is element root + “ic”
Place the word acid after the name
HBr = hydrobromic acid
What do you think HCN would be named?

21. Naming Oxyacids Acids Oxyacids contains an oxyanion and a hydrogen
Determine the oxyanion present
If it ends with “ate” replace with “ic”
If it ends with “ite” replace with “ous”
Acid name = element root + proper ending then acid
HClO3 = Chloric Acid
HClO2 = Chlorous Acid

22. Structural Formulas
Structural Formula: Letter symbols and bonds to show relative positions of atoms
The Lewis Structure can be used to predict many structural formulas

23. Lewis Structure Rules Predict the location of certain atoms
Hydrogen is always the terminal (end) atom
Central atom is the atom with the least attraction for shared electrons
Often closest to the left on the periodic table
2) Determine the total number of electrons available for bonding (valence electrons in the molecule)
Determine number of bonding pairs by dividing the number of electrons available for bonding in two
Place one bonding pair between the central atom and each of the terminal atoms
Subtract number of pairs used in step 4 from total number of bonding pairs. The remaining electrons pairs include lone pairs and those used in double and triple bonds
Convert bonds to double and triple bonds until central atom has an octet.

24. Resonance Structures
Resonance is a condition that occurs when more than one valid Lewis Structure can be written for a molecule or ion

25. Resonance Structures Examples

26. Exceptions to the Octet Rule
Those elements with odd number of valence electrons cannot form a bonded octet
Certain compounds form with fewer than 8 valence electrons present
Coordinate covalent bond: When one atom donates a pair of electrons to be shared with an atom or ion that needs two electrons to be stable
The central atom has more than 8 valence electrons
Will form an expanded octet

27. Octet Rule Exceptions Examples

28. Coordinate Covalent Bond Examples

29. Molecular Shape VESPR Model Bond Angle
Valence Shell Electron Pair Repulsion Model
Based on the arrangement that minimizes the repulsion of shared and unshared pairs of electrons around the central atom
Bond Angle
The angle formed by any two terminal atoms and the central atom
Those predicted by VESPR model are supported by experimental evidence
Lone pairs of electrons will have a slightly larger orbit than the shared electrons

30. Molecular Shape VESPR Model Bond Angle
Valence Shell Electron Pair Repulsion Model
Based on the arrangement that minimizes the repulsion of shared and unshared pairs of electrons around the central atom
Bond Angle
The angle formed by any two terminal atoms and the central atom
Those predicted by VESPR model are supported by experimental evidence
Lone pairs of electrons will have a slightly larger orbit than the shared electrons

31. Vertical to Horizontal: 90o Horizontal to Horizontal 109.5o sp3d
Total Pairs
Shared Pairs
Lone Pairs
Molecular Shape Name
Molecular Shape
Bond Angle
Hybrid Orbital 2
Linear
180o sp 3
Trigonal Planar
120o
sp2 4
Tetrahedral
109.5o
sp3 1
Trigonal Pyramid
107.3o
Bent
104.5o 5
Trigonal Bipyramidal
Vertical to Horizontal: 90o
Horizontal to Horizontal 109.5o
sp3d 6
Octahedral
90o
sp3d2

32. Hybridization
A process in which atomic orbitals are mixed to form new, identical hybrid orbitals
Each hybrid orbital contains 1 electron to be shared with another atom
Lone pairs can also occupy hybrid orbitals

33. Electronegativity and Polarity
Polar Covalent Bonds
Formed when there is a large difference in electronegativity
Results in the unequal sharing of electrons
When formed, the shared electrons are pulled towards one of the atoms
Electrons spend more time around this atom, than the other in the bond
Forms a partial negative and partial positive charge
Polar bonds are often called dipoles

34. Molecular Polarity
Non-polar = Similar electronegativities and charges are balanced within the VESPR Shape a
Polar = Dissimilar electronegativities and charges are not- balanced within the VESPR Shape

35. Solubility
Polar molecules and Ionic Molecules are soluble in polar substances
Non-polar substances are only soluble by non-polar substances

36. Properties of Covalent Compounds
Differences in properties is result of the attractive forces
Covalent: Force between atoms in molecules is strong, but the attraction between individual molecules is weak (intermolecular forces or Van der Waals Forces)
Dispersion Force: The intermolecular force between two non- polar compounds
Dipole-Dipole Force: The intermolecular force between two polar compounds
Hydrogen Bond: Between a Hydrogen and a Flourine, Oxygen or Nitrogen Atom
Weak to Strong

37. Covalent Properties Consist of uncharged or neutral molecules
Usually liquids or gasses
If solid, soft and waxy
Usually a low melting point (exception giant molecular compounds)
Not a conductor in all states (think solids)
Solubility in water depends on polarity