1. Unit 8 Bonding and Nomenclature
Bonding Notes

2. Covalent and Ionic Bonds

3. TYPES OF CHEMICAL BONDS
Chemical bond: a force that holds groups of 2 or more atoms together to function as a unit (molecule).
Ionic bond: An atom (metal) transfers electron(s) to another atom (nonmetal), which creates an electrostatic attraction of a positively charged ion and a negatively charged ion.
Cation: a positively charged ion which is created when an atom loses one or more e-.
Anion: a negatively charged ion which is created when an atom gains one or more e-.

4. TYPES OF CHEMICAL BONDS
Covalent bond: The type of bonding where electrons are shared by two nuclei.
Nonpolar covalent bond: A covalent bond with equal sharing of electrons between the two nuclei.
Polar covalent bond: A covalent bond with unequal sharing of electrons between the two nuclei

5. Dogs Teaching Chemistry

6. Polar/ Nonpolar
Polar Covalent
Nonpolar Covalent

7. ELECTRONEGATIVITY
Electronegativity: the ability of an atom in a molecule to attract shared electrons to itself

8. Electronegativity cont.
Generally increases going from left to right, decreases going down the periodic table.
Metals relatively low electronegativity
Nonmetals  relatively high electronegativity
The higher the atom’s electronegativity value, the closer the shared electrons tend to be to that atom when it forms a bond.

9. ELECTRONEGATIVITY CHART

10. THE RELATIONSHIP BETWEEN ELECTRONEGATIVITY AND BOND TYPE

11. Using Electronegativity to Determine the Type of Bond
Nonpolar: a covalent bond with equal sharing of electrons between the two nuclei. <0.3
Polar: a covalent bond with unequal sharing of electrons between the two nuclei
Above 1.5 is ionic

12. Electronegativity Problems
H2 H-H
H= 2.1
Find the difference: = 0
Nonpolar covalent
HCl H-Cl
H= 2.1, Cl=3.0
Find the difference: = 0.9
Polar covalent

13. Electronegativity Problems con’t
NaCl Na – Cl
Na = 0.9 Cl = 3.0
Find the difference: = 2.1
Ionic bond
What type of bond would be found in between the following atoms?
a. Mg and O b. S and O c. O and O
= 2.3
= 1.0
= 0.0
Ionic
Polar
Non-polar

14. End of Day 1

15. What Holds Bonded Atoms Together
Atoms bond when their valence electrons interact
Remember that atoms with full s and p orbitals are more stable than atoms with only partially full s and p orbitals
Atoms tend to bond together so that they have stable s and p orbitals (stable electron configuration) like the noble gases

16. Valence Electrons
Valence electrons are the electrons that can participate in the formation of a chemical bond. These electrons can be found in the highest energy level of an atom.
In electron configuration, biggest number and usually s and/or p orbitals.

17. Examples of valence electrons
Lithium- Li
Calcium- Ca
Silicon- Si
1s22s1 or [He]2s1
Number of valence electrons?
[Ar]4s2
[Ne] 3s23p2 1 2 2

18. Lewis Structures

19. Lewis Dot Structures
Lewis Structure: a representation of a molecule that shows how the valence e- are arranged among the atoms in the molecule.
Provide a way to draw molecules and better visualize it.
How to draw Lewis Dot structures
First write the element symbol
Determine the number of valence electrons
Add valence electrons around element symbol (one dot represents one valence electron)
Add one electron to each side before adding two!!

20. Lewis Dot structure Practice
Draw Lewis Structures for each of the following atoms
Li C P Br Kr

21. Lewis Dot structures for Molecular Compounds
The most important requirement for the formation of a stable compound is that the atoms achieve a noble gas e- configuration.
Octet rule: 8 electrons are required to fill in the s and the p orbitals in the valence energy level. Exception: Duet rule (2 electrons) for hydrogen.
Bonding pairs: e- that are shared between 2 atoms
Single bond: one pair of e- are shared between 2 atoms
Double bond: two pairs of e- are shared between 2 atoms
Triple bond: three pairs of e- are shared between 2 atoms
Lone pairs: e- that are NOT involved in bonding.

22. Lewis Structures When drawing Lewis structures:
Remember to think about bonding sites (terminal atoms on the “outside”)
N – A = S
electrons needed–electrons available= electrons shared
octet or duet valence electrons bonding electrons

23. Practice drawing Lewis Structures
Draw Lewis electron dot diagrams for each molecule listed below.
1. Br HI NH PCl3
5. H2O CBr SeO N2  

24. End of Day 2

25. Ionic Bonds
Ionic Bonds form between ions with opposite charges (between a metal with a positive charge and a non-metal with a negative charge)
One atom will lose electron(s) and the other will gain electron(s) so that they both have stable electron configurations (8 valence electrons)

26. Ionic Bonds (con’t) EX. NaCl, sodium chloride
Sodium has 1 valence electron, it loses the electron to become a positive ion, Na+
Chlorine has 7 valence electrons, it gains one electron to become a negative ion, Cl-
The oppositely charged ions attract each other and form an ionic bond
Ionic compounds do not form molecules
Ions always exist in the same ratio to one another

27. Ionic Solids
Ionic solids have high melting points, will not conduct electricity as a solid (ions are kept in place), but they will when dissolved in water or when melted (ions are free to move)

28. Lewis Dot structures for Ionic compounds
Write element symbols
Transfer the electrons!
Add the charge!

29. Ionic compounds Lewis practice
Draw the Lewis Structures for the following ionic compounds
1. MgCl MgO K3P

30. Structural Formula
Structural Formula: The overall shape of the molecule is represented by elements symbols and a single covalent bond is represented by a straight line or dash.
Isomers: substances having the same chemical formula but different structures.
Go back to the Lewis Structures for Molecular Compounds and write out the structural formulars