1. Bell Ringer 9/18/13 Copy and complete the following:
Locate the Halogens and Noble Gases on the Periodic Table.
To what Group are they each located?

2. Chemical Bonding
Physical Science

3. Chemical Bonds
When two or more atoms attach to each other, they form a chemical bond
Compounds are any two elements chemically bonded
Water
Sugar
Salt
And almost all other substances!!!!
Electrons are responsible for the type, strength, and size of a chemical bond

4. Lewis Structures
Bohr-Rutherford diagrams are large and difficult to show relationships between multiple atoms
Lewis diagrams are used to show multiple atoms
Lewis diagrams show only the valence electrons

5. Lewis Structures Valence electrons form the charge of an atom
Electrons are always trying to get together in groups of 8 (forget shells for a minute)
Elements that have 8 valence electrons have FULL outer groups
We call these elements NOBLE or INERT gases, they are found in group 8

6. Lewis Structures
Label the Nobel (Inert) gases on your chart

7. Lewis Structures
Elements with 1 valence electron are called the Alkali metals (group 1) (Label)

8. Lewis Structures
Elements with 2 valence electrons are called the Alkaline Earth metals (group 2) (Label)

9. Lewis Structures
Elements with 7 valence electrons are called the Halogens (group 7) (Label)

10. Sy Lewis Structures Consist of Element Symbol
Electrons in each open spot Sy

11. Cl 7 Lewis Structures Lewis Structure Draw the element symbol
Determine the # valence electrons
Starting at the top, going clockwise, place one electron in each spot around the element symbol
Lewis
Structure Cl

12. Lewis Structure
Draw the Lewis Structure for Aluminum! Al

13. Li Quick Recap! Draw the Lewis Structure for Lithium!
Draw the Rutherford-Bohr Diagram for Lithium! - - Li 3P 3N -

14. A note about charges… Writing a charge
Valence electrons, Bohr-Rutherford, and Lewis diagrams are used to determine charge
Charges are a shortcut to determining bonding properties
RULES OF CHARGE
IF the # of valence electrons is GREATER than 4, the charge is negative (Mostly)
IF the # of valence electrons is less than 4, the charge is positive (Mostly)
Charges are in reference to a full shell of 8

15. A note about charges… For example, Aluminum has 3 valance electrons
The possible charges are +3 OR -5
It either has 3 OVER a full shell, or 5 LESS than a full shell
Because the number 3 is less than 4, we use the charge of +3

16. Charges
Any element with a charge is called an ION, the charge is an ionic charge
What are the ionic charges of the elements in the table?
Sodium? +1
Nitrogen? -3
Oxygen? -2
Argon?
Sodium?
Nitrogen?
Oxygen?
Argon?

17. Charges

18. Charges A few exceptions! Example: Boron
Metals are always a positive charge!!
Non metals are always negative!!
Metalloids can go either way (you are not responsible for choosing – I will tell you)
Example: Boron
According to rule of 4’s….its a +3 charge
But since it’s a nonmetal, we use -5!

19. Rules of Bonding All compounds must have neutral charges
(That means the positive charges (cations) and the negative charges (anions) must equal
Subscript numbers are used to show the number of ions
Coefficients are used to show the number of molecules

20. “1’s” are implied and not written
Rules of Bonding
2H2O
Subscript
1 atom of O
“1’s” are implied and not written
Coefficient
Subscript
2 atoms of H

21. Rules of Bonding
H2O O H H

22. Rules of Bonding
H2O O H H

23. Rules of Bonding
Try this one!
NaCl (table salt) Na Cl

24. Rules of Bonding
Last One!
Aluminum Bromide Br Al

25. Chemical Bonding Several Types including Covalent Bonds* Ionic Bonds*
Metallic (only between metals)

26. Covalent Bonds Electrons are shared between two or more atoms
Covalent bonds can exist between atoms of the same type…for example N-N (N2) or O- O (O2)
Covalent bonds can form single, double, or triple bonds
Covalent bonds are strong and usually result in stable molecules
Carbon always forms covalent bonds and forms the basic molecules for all life substances

27. Ionic Bonds
Usually formed by members of the Alkali group (ones with +1 electron)
Electrons are donated to another molecule
Between elements from opposite sides of the chart
Forms crystals (salts) & most dissolve in water

28. Forming Compounds Write ions with charges Cross charges
Write subscripts (omit “1’s”)
Use parenthesis if needed

29. O H H+1 O-2 H2 O1 H2O Forming Compounds
What is the molecular formula of water?
H O-2
H2 O1 H O
H2O

30. Reduce like a fraction to lowest denominator
Forming Compounds
What is the molecular formula of carbon dioxide?
C O-2
C2 O4
Yikes!
Reduce like a fraction to lowest denominator
****note****
C2O4
CO2

31. Al+3 S-2 Al2 S3 Al2S3 Forming Compounds
What is the molecular formula of a compound that has aluminum and sulphur?
Al S-2
Al2 S3
Any guesses on the name?
Al2S3

32. Naming
Binary Covalent

33. Ionic or Covalent??? Ionic - Starts with a metal
Swaps valence electrons
Covalent - 2 non metals
Shares valence electrons

34. Naming Binary Covalent Compounds
Before you can name binary covalent compounds, you MUST know the prefixes!

35. Prefixes Mono Di Tri Tetra Penta Hexa Hepta Octa Nona Deca 1 2 3 4 5 67 8 9 10

36. Rules for naming Binary Covalent Compounds
Name the prefix for the number of atoms of the first element
Then name the first element
Name the prefix for the number of atoms of the second element
Than name the root of the second element with the ending -ide

37. Note… No charges are used in Binary Covalent Compounds
If the 1st prefix is mono….DROP IT!
When the prefix ends in an o or a, and the name of the element begins with a vowel, the o or a is often dropped

38. Examples What is the name of N2O4? N2  di nitrogen O4  tetra oxide
Since oxide begins with a vowel, we will
drop the a in tetra
Dinitrogen tetroxide

39. More examples Name SO2 S  mono sulfur
But mono is with the 1st element, so it will be dropped  sulfur
O2  dioxide
Sulfur dioxide

40. More examples Write the formula for dichlorine monoxide
Dichlorine  Cl2
Monoxide  O
Cl2O

41. More examples Write the formula for disulfur dichloride Disulfur  S2
Dichloride  Cl2
S2Cl2

42. Ionic

43. Writing Chemical Formulas
Before we can name compounds, we must know the rules for writing formulas
The modern model of how atoms react to form compounds is based on the fact that the stability of a noble gas results from the arrangement of its valence electrons.
This model of chemical stability is called the octet rule.

44. Charges or Oxidation Numbers
Elements will gain or loose electrons to have 8 valence electrons.
To review, how many valence electrons does Na have? 1
What will happen to that valence electron
You will loose that electron
If you loose that electron, what charge will you have? +1

45. Charges or Oxidation Numbers
Group 1A  +1
Group 2A  +2
Group 3A  +3
Group 4A  +4 / - 4
Group 5A  -3
Group 6A  -2
Group 7A  -1
Group 8A  STABLE
The charges of monatomic ions, or ions containing only one atom, can often be determined by referring to the periodic table

46. Ions
An ion is an atom or group of combined atoms that has a charge because of the loss or gain of electrons.
A compound that is composed of ions is called an ionic compound.
Only the arrangement of electrons has changed. Nothing about the atom’s nucleus has changed.
Ionic compounds are usually start with a metal or a polyatomic ion (like ammonium)
In ionic compounds, you will SWAP valence electrons

47. Ions
A cation, or positive ion, is formed when an atom loses one or more electrons.
An anion, or negative ion, is formed when an atom gains one or more electrons.
A monatomic ion is one element with a charge
A polyatomic ion is more that one element with a charge

48. Formation of Ionic Compounds
Remember that objects with opposite charges attract each other.
The strong attractive force between ions of opposite charge is called an ionic bond.
Don’t forget that even though the ions have charges, the overall charge of the compound will be …
ZERO!

49. Examples of Formula Writing
Write the formula for the compound formed between sodium and chloride
Na  Na +1
Cl  Cl -1
Na+1 Cl-1 (criss-cross charges & reduce)
Na1Cl (What’s wrong here?)
Do Not write the 1’s
NaCl

50. More examples Write the formula between Mg and Br Mg  +2 Br  -1
Mg+2 Br -1 (Bring down charges)
MgBr2

51. More examples
Write the formula for the compound formed between Ca and S
Ca  Ca +2
S  S -2
Ca S-2 (Bring down charges)
Ca2S (What’s wrong?)
Simplify!
CaS

52. More examples Copper (II) and chlorine CuCl2 Silver and nitrogen Ag3N
Magnesium and sulfur
Calcium and selenium
Potassium and oxygen
Lithium and phosphate
sodium and chlorine
CuCl2
Ag3N
MgS
CaSe
K2O
Li3PO4
NaCl

53. Don’t Forget!
You have to remember the elements that form multiple charges (the ones with the roman numerals)
That roman numeral will tell you the charge!
For example: Copper (II)  Cu +2

54. Multiple Charges to Remember
Copper
Iron
Lead
Tin
Gold
+1 and +2
+2 and +3
+2 and +4
+1 and +3

55. Naming ionic compounds
In naming ionic compounds, name the cation first, then the anion.
Monatomic cations use the element name.
Monatomic anions use the root of the element name plus the suffix -ide.
(This means 1 element with a negative charge will end in –ide).

56. Naming ionic compounds
If an element can have more than one oxidation number, use a Roman numeral in parentheses after the element name, for example, iron(II) to indicate the Fe 2+ ion.
For polyatomic ions, use the name of the ion.

57. Simply put..
All you have to do is name the 1st thing then name the 2nd thing

58. Examples NaCl MgS K3P CaCl2 CuBr AlCl3 CuS Fe3N2 Sodium chloride
Magnesium sulfide
Potassium phosphide
Calcium chloride
Copper (I) bromide
Aluminum chloride
Copper (II) sulfide
Iron (II) nitride

59. More examples Lead (IV) Oxide Potassium fluoride Iron (II) chloride
Calcium sulfide
Lithium nitrode
Sodium selenide
Tin (II) chloride
PbO2 KF
FeCl2
CaS
Li3N
Na2Se
SnCl2

60. Mixed examples (remember to figure out what type of compound it is 1st
KCl
CO2
Na2S
LiBr
CuI2
Fe2O3
Al2O3
Potassium chloride
Carbon dioxide
Sodium sulfide
Lithium bromide
Copper (II) iodide
Iron (III) oxide
Aluminum oxide

61. More Mixed Examples Carbon tetrachloride Phosphorous pentachloride
Aluminum oxide
Copper (II) nitrate
Chlorous acid
Hydrophosphoric acid
Iron (III) hydroxide
CCl4
PCl5
Al2O3
Cu(NO3)2
HClO2
H3P
Fe(OH)3