1. Chemical Bonds

2. The Elements 90 naturally occurring elements
Most are not found as pure elements
The majority of elements are found combined with other elements to form compounds.
Gold, silver and platinum are rare examples of metals found in elemental form (precious metals)

3. Classifying
Since the 90 elements can form thousands of different compounds, classification systems have been developed.
The classification of compounds is based on their properties to help our understanding of compounds.
Example: melting point, boiling point, hardness, etc.

4. Electron Configuration
Atomic number = number of electrons
Electrons vary in the amount of energy they possess, and they occur at certain energy levels or electron shells.
Electron configuration determines how an atom behaves when it encounters other atoms

5. Electron Dot Structure
Symbols of atoms with dots to represent the valence-shell electrons H He:            Li Be  B   C   N   O  : F  :Ne :                    Na Mg  Al  Si  P S :Cl  :Ar :        

6. Chemical Bonds
Octet Rule- rule that states that atoms tend to gain, lose, or share electrons so that each atom has full outermost energy level which is typically 8 electrons. There are two ways to try to do this 1. Ionic bonds – 2. Covalent bonds –

7. Octet Rule = atoms tend to gain, lose or share electrons so as to have 8 electrons
C would like to
N would like to
O would like to
Gain 4 electrons
Gain 3 electrons
Gain 2 electrons

8. Properties of Ionic and Covalent Componds

9. Chemical Bonds
forces that attract atoms to each other to form compounds
involves the interactions of valence electrons between atoms
usually the bond forms a compound that is more stable than the atoms individually.

10. Ionic Bond
bond formed between two ions by the transfer of electrons

11. bond formed by the sharing of electrons
Covalent Bond
bond formed by the sharing of electrons

12. Formation of Ions from Metals
Ionic compounds result when metals react with nonmetals
Metals lose electrons to match the number of valence electrons of their nearest noble gas
Positive ions (cations) form when the number of electrons are less than the number of protons
Group 1 metals  ion 1+
Group 2 metals  ion 2+
Group 13 metals  ion 3+

13. Formation of Sodium Ion
Sodium atom Sodium ion
Na  – e  Na +
( = Ne)
11 p p+
11 e e-

14. Ions from Nonmetal Ions
In ionic compounds, nonmetals in 15, 16, and 17 gain electrons from metals (anions)
Nonmetal add electrons to achieve the octet arrangement
Nonmetal ionic charge:
3-, 2-, or 1-

15. Fluoride Ion     1 - : F  + e -> : F :     2-7 2-8 (= Ne)
unpaired electron octet
   
: F  e -> : F :
   
(= Ne)
9 p p+
9 e e-
ionic charge

16. Ionic Bond
Between atoms of metals and nonmetals (usually) with very different electronegativity
Electronegativity difference e.d. > 1.7
Bond formed by transfer of electrons
Examples; NaCl, CaCl2, K2O

17. Some characteristics of an Ionic Bond
Crystalline at room temperatures
Have higher melting points and boiling points compared to covalent compounds
Conduct electrical current in molten or solution state but not in the solid state
Polar bonds

18. Covalent Bond
Between nonmetallic elements of similar electronegativity.
Formed by sharing electron pairs
Examples; O2, CO2, C2H6, H2O, SiC

19. Some characteristics of a Covalent Bond
Covalent bonds have definite and predicable shapes.
Very strong
Low melting and boiling points

20. Covalent Bonds can have multiple bonds, so you should be familiar with the following…
Single Covalent Bond- chemical bond resulting from sharing of an electron pair between two atoms.
Triple Covalent Bond-chemical bond resulting from sharing of three electron pairs between two atoms.
Double Covalent Bond- chemical bond resulting from sharing of two electron pairs between two atoms.

21. Non-polar covalent Bonds
when electrons are shared equally

22. when electrons are shared but shared unequally
Polar Covalent Bonds
when electrons are shared but shared unequally

23. water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

24. Review of Bonds

25. Electronegativity (EN)
a measure of an atoms ability to attract electrons in a chemical bond.
a property of an atom involved in a bond

26. Using EN to predict the bond type
When two atoms form a bond the difference in electronegativity (ΔEN) can help to determine the bond type.
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27. Examples HCl 3.16- 2.2= 0.96 Polar Covalent CrO 3.44- 1.66= 1.78 Ionic
Br2
= Covalent