1. Properties of Molecular Substances
Chapter 8
Covalent Bonding
VSEPR Theory
Molecular Shape
Polar or NonPolar
Properties of Molecular Substances

2. Why Do Atoms Bond?
The stability of an atom, ion or compound is related to its energy
lower energy states are more stable.
Metals and nonmetals gain stability by transferring electrons (gaining or losing) to form ions that have stable noble-gas electron configurations.
Ionic Bonding
Another way atoms can gain stability is by sharing valence electrons with other atoms, which also results in noble-gas electron configurations.
Covalent Bonding

3. The most stable arrangement of atoms exists at the point of maximum net attraction, where the atoms bond covalently and form a molecule.

4. The Covalent Bond
Atoms will share electrons in order to form a stable octet.
Covalent bond : the chemical bond that results from the sharing of valence electrons
also called a molecular bond

5. The Molecule formed when two or more atoms bond covalently
The smallest piece in a covalent compound
Formed when the proton of one atom is attracted to the electron cloud of another atom.

6. Models
Molecules

7. Single Covalent Bonds
In a single covalent bond a single pair of electrons is shared
This can be represented with a Lewis structure
A single line represents a single covalent bond
A single pair of electrons

8. Bonding pair: a pair of electrons shared by two atoms
Lone pair: an unshared pair of electrons on an atom

9. Formation of Water

10. Group 17 elements will form one covalent bond.

11. Group 16 elements will form two covalent bonds.

12. Group 15 elements will form three covalent bonds.

13. Group 14 elements will form four covalent bonds.

15. Strength of Covalent Bonds
The strength of covalent bonds is determined by the bond length:
distance between the bond nuclei
Bond length is determined by:
The size of the atoms involved—larger atoms have longer bond lengths
How many pairs of electrons are shared—the more pairs of electrons shared, the shorter the bond length is.

16. Bond Dissociation Energy
the amount of energy required to break a bond
Indicates the strength of a covalent bond
When a bond forms, energy is released;
When a bond breaks, energy must be added
Each covalent bond has a specific value for its bond dissociation energy.

17. Bond Energy and Bond Length
A direct relationship exists between bond energy and bond length
Shorter Bond
Stronger Bond
Higher Bond Dissociation Energy
Longer Bond
Weaker Bond
Lower Bond Dissociation Energy

18. Energy Changes
An endothermic reaction is one where a greater amount of energy is required to break a bond in reactants than is released when the new bonds form in the products.
An exothermic reaction is one where more energy is released than is required to break the bonds in the initial reactants.

19. Common Names
Many compounds were discovered and given common names long before the present naming system was developed (water, ammonia, hydrazine, nitric oxide).

20. Structural Formulas
A structural formula uses letter symbols and bonds to show relative positions of atoms.

21. Lewis Structures Used to predict the structural formula
Show arrangement of the atoms and un-bonded electrons

22. Five steps to draw Lewis structures:
Count the total number of valence electrons in all atoms involved.
Decide how the elements are arranged in the structure and draw it out.
Hydrogen is always an end atom.
Central atom is usually written first in compound
Central atom has least attraction for the electrons
Usually closer to left on periodic table
Subtract the # of electrons used in the bonds.

23. Satisfy the octets of the terminal atoms.
Place any remaining electrons around the central atom to satisfy its octet. If the central atom cannot be satisfied, make a multiple bond using a lone pair from the terminal atoms.
Check your work 

24. Examples

25. Count the total number of valence electrons in all atoms involved.
Drawing Lewis structures for polyatomic ions is very similar to drawing Lewis structures for covalent compounds EXCEPT in finding the number of electrons available for bonding
Count the total number of valence electrons in all atoms involved.
If the polyatomic ion is negatively charged, ADD the charge to the number of valence electrons.
If the ion is positively charged, SUBRACT the charge from the number of valence electrons.
Follow the rest of the steps to drawing Lewis structures.

26. Examples

27. Sub Octet
Some compounds form with fewer than 8 electrons present around an atom.
Boron
BF3

28. Molecular Shape How a molecule “looks” Determines properties
The shape of a molecule determines whether or not two molecules can get close enough to react
We describe shape using the VSEPR model

29. VSEPR
This model is based on the fact that electrons pairs will stay as far away from each other as possible
Valence
Shell
Electron
Pair
Repulsion

30. How to apply VSEPR Draw the Lewis Structure for a Molecule
Count the pairs of bonded electrons
Count the pairs of unbonded electrons
Match the information with the VSEPR chart to classify the shape of the molecule

33. Atoms will assume certain bond angles: the angle formed by any two terminal atoms and the central atom
Lone pairs take up more space than bonded pairs do.

35. Molecular Polarity
Molecules are either polar or nonpolar depending on the bonds in the molecule.
We must look at the shape (geometry) of a molecule to determine polarity.
Symmetric molecules are nonpolar.
Asymmetric Molecules are Polar

36. Polar or NonPolar Determine if a molecule is polar or nonpolar by
Looking at a model of the molecule
Looking at a Lewis Structure of the molecule

37. Solubility of Polar Molecules
Bond type and shape of the molecule determine solubility
Polar substances and ionic substances will dissolve in polar solvents
Nonpolar substances will only dissolve in nonpolar substances

38. Intermolecular Forces
Nonpolar Molecules: Van der Waals intermolecular Forces
Very Weak forces between molecules
Polar Molecules: have dipole-dipole intermolecular bonding.
Stronger intermolecular Forces
Polar Molecules with Hydrogen Bonding: hydrogen bonded to nitrogen, oxygen or fluorine, it will have hydrogen bonding between molecules. A very strong dipole-dipole interaction
Very strong intermolecular Forces
High boiling points, high melting points