1. CHEMICAL BONDING AND NOMENCLATURE

2. CHEMICAL BONDING
When atoms interact to form a chemical bond, only the outer energy level electrons, the valence electrons interact.
Electron Dot Diagrams are used to keep track of the valence electrons.

3. CHEMICAL BONDING
A Chemical Bond is the force that holds atoms together in a compound.
They are formed in definite ways according to certain rules.

4. PRACTICE
Draw the Electron Dot Diagrams for the following elements: K Se Be Ga Ge P Ar

5. CHEMICAL STABILITY
It is the goal of every atom to have a full outer energy level. Remember full is eight electrons (octet) except for hydrogen and helium which have two.
The atom gets a full energy level by losing or gaining electrons to form an octet.
This is the most stable formation.

6. CHEMICAL STABILITY
The number of valence electrons determines whether or not the atom will bond with another element in a chemical reaction.
Noble gases are chemically stable. Their outer energy levels are full.
Every element wants the same number of valence electrons as a noble gas and will gain electrons or lose electrons to get it.

7. IONS
IONS: Ions are charged particles that have more or fewer electrons than protons.
Cation: positive ion (lost electrons)
Anion: negative ion (gained electrons)

8. OXIDATION NUMBER
positive or negative number that indicates how many electrons an atom has gained, lost, or shared to become stable.
Is written as a superscript with a positive or negative sign.

9. OXIDATION NUMBER
Na+1 means that sodium has p+ = 11 no = 12 but e- = 10. It is a +1 charge because there is one more proton than there are electrons.
Cl-1 means that chlorine has p+ = 17 no = 18 but e- = 18. It is a -1 charge because there is one more electron than there are protons.

10. OXIDATION NUMBER
In order for sodium to be “happy”/stable, it wants to lose an electron. In order for chlorine to be “happy”/stable, it wants to gain an electron.

11. PRACTICE
Look at the electron dot diagrams you did above. Predict the ion formed for each below:
K ___ Se___ Be___
Ga___ Ge___ P___
Ar___

12. PRACTICE Write the correct ion for each below: K ___ Se___ Be___
Ga___ Ge___ P___
Ar___

13. PRACTICE Group 1 ve: _____ ion: ____ Group 2 ve: _____ ion: ____
Write the number of valence electrons in each of the groups listed. Then write the ion formed.
Group 1 ve: _____ ion: ____
Group 2 ve: _____ ion: ____
Group 13 ve: _____ ion: _____ Group 14 ve: _____ ion: _____ Group 15 ve: _____ ion: _____ Group 16 ve: _____ ion: _____
Group 17 ve: _____ ion: _____
Group 18 ve: _____ ion: _____

14. IONIC BONDING
Video

15. IONIC BONDS involves a TRANSFER of electrons from one atom to another.
is held together by electric/electrostatic force (+/-).
happens between positive ions and negative ions.
When the cation and the anion come together, there is zero net charge.
The basic rule is that opposites attract. A strong attraction between oppositely charged ions (formed by the transfer of electrons) hold the ions together in an ionic bond.

16. EXAMPLE

17. CHARACTERISTICS high melting and boiling points
tend to be soluble in water
Solids are poor conductors of heat and electricity.
When dissolved in a liquid or melted, they become good conductors.
Generally forms between elements on opposite sides of the periodic table.

18. COVALENT BONDS
Atoms with similar electron affinities tend to SHARE electrons.
There are not enough electrons for each individual atom to have a full octet.
Neither atom wants to lose electrons; both want to gain them.
Usually forms between nonmetals.

19. COVALENT BONDING
The positively charged nucleus of both atoms attracts the negative electrons to be shared.
Covalent bonds form a unit called a MOLECULE.
A molecule is the smallest particle of a covalent bond that has all the properties of that substance.

20. COVALENT BONDING

21. CHARACTERISTICS low melting points compared to ionic
many exist as gases or vaporize easily at room temperature
relatively soft, brittle
Poor conductors of electricity.

22. TYPES OF COVALENT BONDS
1. Single - When a single pair of electrons is shared
2. Multiple - When more than one pair of electrons are shared.

23. Types of Covalent Bonds

24. SHARING ELECTRONS
Electrons are not always shared equally between atoms in a covalent bond.
The strength of the attraction of each atom is related to:
1. Size of the atom
2. Charge of the nucleus
3. Total number of electrons the atom contains.

25. SHARING ELECTRONS
An example of unequal sharing is HCl. Chlorine has a stronger attraction for electrons than hydrogen does.
As a result, the electrons will spend more time surrounding the chlorine nucleus than they will the hydrogen nucleus.
The chlorine atom is partially negative and the hydrogen atom is partially positive.

26. SHARING ELECTRONS
This is called a polar molecule – a molecule with a slightly positive end and a slightly negative end resulting in the electrons being shared unequally.
Molecules where the electrons are shared equally and do not have oppositely charged ends, like H2, N2, and O2, are nonpolar molecules.

27. SHARING ELECTRONS

28. NAMING COMPOUNDS AND FORMULAS - NOMENCLATURE
Three Goals:
Interpret information given by a chemical formula.
Write formulas for various compounds.
Name compounds given a formula.

29. NAMING COMPOUNDS AND FORMULAS - NOMENCLATURE
A chemical formula tells:
what type of atoms are present.
the number of each type of atom present.
ratio of one atom to another.
H2O Cl2 Al2O3

30. NAMING COMPOUNDS AND FORMULAS - NOMENCLATURE
Before naming a compound, you must determine which category it fits into:
Ionic OR
Molecular
(Covalent)

31. IONIC NOMENCLATURE contains cations and anions.
Includes metals and nonmetals.
Metals go first in name and formula.

32. IONIC NOMENCLATURE Which symbol goes first? S and Cu d. C and Mg
Bi and S e. Ta and Cl
c. N and Nb f. Al and As

33. NAMING IONIC COMPOUNDS
Write the name of the cation (metal)
Write the name of the anion and change the ending to “ide”.

34. PRACTICE
CaF2
CaCl2
MgBr2
BeO
BeS
B2S3
NaF
K2O

35. POLYATOMIC IONS
An ion made up of more than one atom NO3-1, SO4-2, C2H3O2-1;
The charges given to polyatomic ions apply to the whole group of atoms in the ion.
NEVER change the subscripts of a polyatomic ion.
If more than one ion is needed, a parenthesis is placed around the ion and the subscript is written outside; Example: Al(OH)3

36. POLYATOMIC IONS Acetate, C2H3O2-1 Ammonium, NH4+1 Bicarbonate, HCO3-1
Ones You Might Run Into:
Acetate, C2H3O2-1
Ammonium, NH4+1
Bicarbonate, HCO3-1
Carbonate, CO3-2
Hydroxide, OH-1
Nitrate, NO3-1
Phosphate, PO4-3
Sulfate, SO4-2

37. PRACTICE Ca(OH)2 Na2SO4 KC2H3O2 NH4NO3 Mg(HCO3)2 Al2(CO3)3 Al(NO3)3
AlPO4

38. WRITING SIMPLE IONIC FORMULAS
1. Write the charge for the element over the name. You determine the charge by looking at your periodic table or the list of ions in your notes.
2. Write the symbol for the cation, followed by the symbol for the anion.

39. WRITING SIMPLE IONIC FORMULAS
Monatomic Ions
An ion with only one atom: Mg+2, Na+1, O-2.
Most of the charges can be determined by the atom’s position on the periodic table.

40. WRITING SIMPLE IONIC FORMULAS
3. Crisscross the numbers only (not the signs), writing the oxidation number as a subscript for the other. Do not write the number 1. Balance the charges by placing the subscripts. Net charges must be zero.
4. Make sure the formula is in its simplest form. Mg2O2 is reduced to MgO, but MgCl2 cannot be reduced.

41. WRITING SIMPLE IONIC FORMULAS
5. Check your work by calculating the total positive and negative charges and confirming that the total charge on the compound is zero. MgO - Mg: +2 x 1 and O: -2 x 1 = +2 and -2 which is zero.

42. PRACTICE lithium fluoride lithium chloride Barium bromide barium oxide
Aluminum iodide
aluminum oxide

43. MULTIPLE OXIDATION NUMBERS
Some cations (positive ions) can have more than one charge.
Most are transition elements. Lead and tin are exceptions.
It is important to distinguish which ion is in the compound.

44. MULTIPLE OXIDATION NUMBERS
Cu copper (I)
Cu copper (II)
Fe iron (II)
Fe iron (III)
Sn tin(II)
Sn tin (IV)
Pb lead (II)
Pb lead(IV)

45. WRITING THE FORMULAS FOR THESE TYPES
Write the name of the cation (metal)
Write the name of the anion ending in “ide”.
Just remember that the charge for the cation is in the name.
Examples: copper (II) sulfate, CuSO4 iron (II) oxide, FeO

46. PRACTICE Copper(I) fluoride Copper(II) oxide Chromium (II) bromide
Manganese (II) iodide
Lead(II) chloride
Cobalt (III) sulfide

47. NAMING MORE COMPLEX IONIC COMPOUNDS
1. Write the name of the cation (metal). 2. Write the name of the anion. If the cation (metal) has more than one charge, designate the correct charge with the roman numeral using the formula as a guide to determine the charge. Example: Fe(OH)2 iron (II) hydroxide Fe(OH)3 iron (III) hydroxide

48. PRACTICE
1. FeO
2. Fe2O3
3. CuCl2
4. CuCl
5. PbF4
6. SnCl2

49. PRACTICE 1. Potassium hydrogen carbonate 2. Iron(II) hydroxide
3. Iron(III) nitrate
4. Beryllium carbonate
5. Sodium acetate
6. Potassium sulfate

50. HYDRATES – write this on second to last page
A crystalline compound in which its ions are attached to one or more water molecules.
Hydrates are solids with water molecules trapped in them.
Example: CuSO4  5H2O
copper (II) sulfate pentahydrate

51. PRACTICE
1. Na2PO4  2H2O
2. SrCl2  6H2O
3. Ba(OH)2  8H2O

52. NPR REPORT
Dihydrogen monoxide

54. FAQ
Dihydrogen Monoxide (DHMO) is a colorless and odorless chemical compound, also referred to by some as Dihydrogen Oxide, Hydrogen Hydroxide, Hydronium Hydroxide, or simply Hydric acid. Its basis is the highly reactive hydroxyl radical, a species shown to mutate DNA, denature proteins, disrupt cell membranes, and chemically alter critical neurotransmitters. The atomic components of DHMO are found in a number of caustic, explosive and poisonous compounds such as Sulfuric Acid, Nitroglycerine and Ethyl Alcohol.

55. MOLECULAR COMPOUNDS Another word for covalent molecules
usually binary;
called a molecule
usually contain two nonmetals; solid, liquid or gas
Sometimes contains a metalloid and a nonmetal.

56. NAMING MOLECULAR COMPOUNDS
Rules:
Use prefixes to identify the number of atoms present.
If only one of the first atom (nonmetal), then NO prefix is used.
Last element ends in “ide”.
* Can combine in different ratios so MUST use prefixes.

57. NAMING MOLECULAR COMPOUNDS
Prefixes:
1 - mono hexa
2 - di hepta
3 - tri octa
4 – tetra nona
5 - penta deca

58. NAMING MOLECULAR COMPOUNDS
Examples:
CO2 carbon dioxide
PCl3 phosphorus trichloride
N2O dinitrogen monoxide
P2O3 diphosphorus trioxide

59. PRACTICE
1. NCl3
2. CS2
3. Si2Br6
4. S4N4
5. SF6
6. N2O5

60. WRITING MOLECULAR COMPOUNDS
1. Write the symbol of the first element adding a subscript of the number of atoms present. If only one of the first atom (nonmetal), then NO number is used.
2. Write the symbol of the second element adding a subscript of the number of atoms present.

61. PRACTICE boron trifluoride silicon tetrachloride iodine heptafluoride
tetraphosphorus decoxide
chlorine trifluoride
sulfur dichloride

62. PUTTING IT ALTOGETHER
You need to make sure that you can distinguish between ionic AND covalent compounds.