1. Chapter 8 – Covalent Bonding
Ms. Wang
Lawndale High School

2. Chapter 8.1 – Molecular Compounds
In Chapter 7, we learned about electrons being “given up” or “stolen away”
This type of “tug of war” bond between a metal and nonmetal is called an ionic bond
In this chapter, you will learn about another type of bond
Covalent Bond – atoms held together by sharing electrons between two nonmetals

3. Molecules
We know that a metal cation and nonmetal anion are joined together by an ionic bond and called a SALT
A neutral group of atoms joined together by a covalent bond is called a MOLECULE

4. Monatomic vs. Diatomic Molecules
Molecules can be monatomic or diatomic
Diatomic Molecule – a molecule consisting of two atoms
There are 7 diatomic molecules on the periodic table (SUPER 7) – N2, O2, F2, Cl2, Br2, I2, H2

5. Properties of Molecular Compounds
Lower Melting and Boiling Points than Ionic Compounds
Gases or liquids at room temperature

6. Molecular Formulas
Molecular Formula – the chemical formula of a molecular compound
It shows how many atoms of each element a molecule contains
Example
H2O contains 3 atoms (2 atoms of H, 1 atom of O)
C2H6 contains 8 atoms (2 atoms of C, 6 atoms of H)

7. Practice
How many atoms total and of each do the following molecular compounds contain? H2 CO
CO2
NH3
C2H6O

8. Chapter 8.2 – Covalent Bonding
Remember that ionic compounds either gain or lose electrons in order to attain a noble gas electron configuration
Covalent compounds form by sharing electrons to attain a noble gas electron configuration
So the Octet Rule still applies to covalent bonds

9. Single Covalent Bond
Single Covalent Bond - two atoms held together by sharing one pair of electrons
Unshared Pair / Lone Pair / Nonbonding Pair – a pair of valence electrons that is not shared between atoms
Let’s Practice (you can use dots or lines)
F2, H2O, CH4

10. Double and Triple Covalent Bonds
Double Covalent Bond – a bond that involves two shared pairs of electrons
Triple Covalent Bond – a bond that involves three shared pairs of electrons
Let’s Practice O2 N2

11. Polyatomic Ion
Polyatomic Ion - tightly bound group of atoms that has a positive or negative charge and behaves as one unit
Some examples are NH4+ and SO32-

12. Bond Dissociation Energy
Bond Dissociation Energy - the energy required to break the bond between two covalently bonded atoms
A large bond dissociation energy corresponds to a strong covalent bond
For example, carbon-carbon has a strong bond dissociation energy so it is not very reactive

13. Chapter 8.3 - Bonding Theories
So far, the orbitals we have been discussing are atomic orbitals (s, p, d, f) for each atom
When two atoms combine, their atomic orbitals overlap and make molecular orbitals
Molecular Orbitals – orbitals that apply to the entire molecule instead of just one atom

14. Molecular Orbitals
Just as atomic orbitals belong to a particular atom, a molecular orbital belongs to a molecule as a whole
Each orbital is filled with 2 electrons
Bonding Orbital – a molecular orbital that can be occupied by two electrons of a covalent bond

15. S orbitals overlapping P orbitals overlapping end-to-end
Sigma Bond ()
Sigma Bond - when 2 atomic orbitals combine to form a molecular orbital that is symmetrical around the axis
S orbitals overlapping
P orbitals overlapping end-to-end

16. P orbitals overlapping side-by-side
Pi Bond ()
Bonding electrons likely to be found in a sausage-shape above and below the axis
Pi bonds tend to be weaker than sigma bonds because pi bonds overlap less than sigma bonds
P orbitals overlapping side-by-side

17. VSEPR Theory Explains the 3D shape of molecules
According to VSEPR theory, the repulsion between electron pairs causes molecular shapes to adjust so that the valence electron pairs stay as far apart as possible

18. A Few VSEPR Shapes

19. Nine possible molecular shapes

20. VSEPR Theory
Unshared pairs of electrons (lone pairs) are very important in predicting the shapes of molecules
Examples
Methane (CH4) - tetrahedral
Ammonia (NH3) - pyramidal
Water (H2O) – bent
Carbon Dioxide (CO2) - linear

21. Hybrid Orbitals
VSEPR is good at describing the molecular shapes, but not the types of bonds formed
Orbital hybridization provides information about both molecular bonding and molecular shape
In hybridization, several atomic orbitals mix to form the same total number of hybrid orbitals

22. Bond Hybridization Hybridization Involving Single Bonds – sp3 orbital
Ethane (C2H6)
Hybridization Involving Double Bonds – sp2 orbital
Ethene (C2H4)
Hybridization Involving Triple Bonds – sp orbital
Ethyne (C2H2)

23. Chapter 7.4 – Polar Bonds and Molecules
There are two types of covalent bonds
Polar Covalent Bonds
Nonpolar Covalent Bonds
Polar Covalent Bond – unequal sharing of electrons where one atom has a slightly negative charge and the other atom has a slightly positive charge (HCl, H2O)
Nonpolar Covalent Bond – equal sharing of electrons between two atoms (Cl2, N2, O2)

24. Classification of Bonds
You can determine the type of bond artificially by calculating the difference in electronegativity between elements
Type of Bond
Electronegativity Difference
Nonpolar Covalent
0  0.4
Polar Covalent
0.5  1.9
Ionic
2.0  4.0

25. Let’s Practice Together
What type of bond is HCl? (H = 2.1, Cl = 3.1)
Difference = 3.1 – 2.1 = 1.0
Therefore it is polar covalent bond.
Your Turn To Practice
N(3.0) and H(2.1)
H(2.1) and H(2.1)
Ca(1.0) and Cl(3.0)
Al(1.5) and Cl(3.0)
Mg(1.2) and O(3.5)
H(2.1) and F(4.0)

26. Dipole
No bond is purely ionic or covalent … they have a little bit of both characters
When there is unequal sharing of electrons a dipole exists
Dipole is a molecule that has two poles or regions with opposite charges
Represented by a dipole arrow pointing towards the more negative end.

27. Practice Drawing Dipoles
P- Br
P = 2.1
Br = 2.8
P –Br
 -
Practice
H(2.1) – S(2.5)
C(2.5) – F(4.0)
Si(1.8) – C(2.5)
N(3.0) – O(3.5)

28. Attractions Between Molecules
Intermolecular attractions are weaker than either ionic or covalent bonds
Van der Waals forces – consists of the two weakest attractions between molecules
dipole interactions – polar molecules attracted to one another
dispersion forces – caused by motion of electrons (weakest of all forces)

29. Hydrogen Bond
Hydrogen Bonds - attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom

30. Hydrogen Bond
This other atom may be in the same molecule or in a nearby molecule, but always has to include hydrogen
Hydrogen Bonds have about 5% of the strength of an average covalent bond
Hydrogen Bond is the strongest of all intermolecular forces

31. Intermolecular Attractions
A few solids that consist of molecules do not melt until the temperature reaches 1000ºC or higher called network solids (Example: diamond, silicon carbide)
Network Solid – solids in which all of the atoms are covalently bonded to each other
Melting a network solid would require breaking covalent bonds throughout the solid

32. Chapter 8 Assessment Page 247
Homework
Chapter 8 Assessment Page 247
#’s 39-41, 43-46, 51, 53, 54, 57-59, 61, 65, 68, 83, 85, 86, 89, 96, 99, 100