1. Adventures of Oxygen Clip
Chemical
Bonding 1
Adventures of Oxygen Clip

2. 2
GOALS
1. Compare & contrast ionic and covalent bonds in terms of electron position.
2. Predict formulas for stable binary ionic compounds based on balance of charges.
3. Determine the Types of ions formed by representative elements
4. Use IUPAC nomenclature for transition between chemical names and chemical formulas of
- binary ionic compounds
- binary covalent compounds 34 16
(IUPAC) International Union of Pure and Applied Chemistry

3. Why do Atoms Form Compounds?3
Stability.
What makes an atom stable?
Full outer energy level.
Eight.
They can either……
1) Gain electrons
2) Lose electrons
3) Share electrons

4. 4
A Chemical Bond holds atoms together in a compound.
Two basic types:
1. Ionic
2. Covalent

5. Ionic Bonding Transfer of electrons from one atom to another atom.5
Ionic Bonding
Transfer of electrons from one atom to another atom.
Occurs between metals & nonmetals.
Remember: Atoms need a full outer energy level to be stable. EIGHT!
Called compounds.

6. Ionic Bonding Opposites ATTRACT! Occurs between metals and nonmetals.6
Ionic Bonding
Occurs between metals and nonmetals.
Metals are electron donors.
SO, they become POSITIVE
Non-metals are electron accepters.
SO, they become NEGATIVE.
Opposites ATTRACT!

7. When Atoms gain or lose electrons, they are called Ions.3P 3P 3P
Anion 3P 3P 3P
Cation

8. Metals lose electrons to become stable.
Nonmetals gain electrons to become stable.

9. Ionization: requires energy Why do atoms lose and gain electrons?
Atoms can gain or lose electrons
Ionization: requires energy
Why do atoms lose and gain electrons?
To become more stable.
Stability=full outer energy level

11. 11
Animation Examples
Opposites Attract!

12. Properties of Ionic Compounds12
Properties of Ionic Compounds
Crystalline solids at room temperature.
Arranged in repeating three-dimensional patterns
Have high melting points
Can conduct electricity when melted or dissolved in water

13. 13
Ionic Bonding
CLIP

14. Occurs between nonmetals and nonmetals. 14
Covalent Bonding
Occurs between nonmetals and nonmetals.
The sharing of electrons between atoms.
Called Molecules.

15. 15

16. 16
Hydrogen and Fluorine
Hydrogen and Chlorine

17. 17
Single, Double, Triple
2 e e e-

18. 18
Clip

19. Polar molecules happen when one atom has a greater positive charge19
Unequal Sharing
Called Polar δ+ δ_
Polar molecules happen when one atom has a greater positive charge

20. Properties of Covalent Molecules20
Properties of Covalent Molecules
Many are gases or liquids at room temperature
Composed of two nonmetals.
Have low melting and boiling points… covalent bonds are weaker than ionic bonds

21. Covalent or Ionic? (write the formula, then write “C” or “I”)21
Covalent or Ionic? (write the formula, then write “C” or “I”)
CO2
NaCl
H2O
MgCl2
NO2
Li2S
NaF
BeCl2
BeO
HCl
KCl
H2O2 N2
Cl2
AgCl2
clip

22. Ionic and Covalent Bonding Review Clip

23. Writing
Chemical
Formulas
Goals revisited

24. 22 The chemical formula for water is H2O. Carbon Dioxide is CO2.
Writing chemical formulas is a shorthand way of indicating what a substance is made of. 
These formulas also let you know how many atoms of each type are found in a molecule. 
The chemical formula for water is H2O. 
Carbon Dioxide is CO2. 
Why does oxygen combine in different ratios, in different compounds? 
The chemical formula for table salt is NaCl. Calcium Chloride is CaCl2.
Why does chlorine combine in different ratios, in different compounds? 

25. The simplest compounds are ones with only two elements23
The simplest compounds are ones with only two elements
These are called binary
KI, CO, H2O, NaCl

26. Oxidation numbers
+4 -4 +1
Tell you how many electrons an atom must gain, lose or share to become stable. -2 +2 +3 -3 -1 24

27. All compounds are neutral That means the overall charge is ZERO!25
Oxidation numbers
We can predict the ratio of atoms in ionic compounds based on their oxidation numbers +1 -1
1 valence electron K Cl
7 valence electron
All compounds are neutral
Tells you how many electrons an atom must gain, lose or share to become stable.
KCl
That means the overall charge is ZERO!

28. To make it ZERO, you need 1 Ca & 2 Br.+1 -1 +2 -1 Na Br Ca Br
To make it ZERO, you need 1 Ca & 2 Br.
NaBr
CaBr2
Subscripts show the number of atoms of that kind in the compound 26

29. Now You Try writing Binary Ionic formulas27
Now You Try writing Binary Ionic formulas
K + Br
Mg + Cl
Ca + I
K + O
K + I
Sr + Br
Na + O
Ga + Br

30. Some elements have more than one oxidation number (Chart p588)+3 -2 +2 -2 Co O Co O
Co2O3
CoO
We call these elements- Multivalent Elements 28

31. 29
Multivalent Practice
Fe+2 + O
Fe+3 + S
Cu+2 + F
Cr+3 + Br

32. Naming
Chemical
Formulas

33. Naming Binary Compounds and Molecules32
Steps:
If it is Binary-
Decide if it is an ionic or covalent bond.
Metal- nonmetal…..
Ionic
Nonmetal- nonmetal….
Covalent
Example:
NaCl

34. If ionic ……. Examples: CaO K2S NaCl K2O AlCl3 BaF2 33 Example: NaCl
With a friend:
NaCl
K2O
AlCl3
BaF2
Check to see if any elements are multivalent.
If all single valent, write the name of the positive ion first.
Write the root of the negative ion and add –ide.

35. If ionic ……. 34
Check to see if any elements are multivalent.
If multivalent ions, determine the oxidation number of the element.
Use Roman numerals in parentheses after the name of the element.
Write the root of the negative ion and add –ide.
NiCl
Mn2S

36. Ionic-multivalent
Examples:
FeO
Fe2O3
CuO
Cu2O

37. 36
If Covalent...
Use Greek prefix to indicate how many atoms of each element are in the molecule
Add -ide to the more electronegative element
Greek Prefixes
1- mono-
2- di-
3- tri-
4- tetra-
5- penta-
6- hexa-
7- hepta-
8- octa-
Example:
PCl3
Phosphorous trichloride NO
Nitrogen Monoxide

38. Naming Covalent Practice
36 ¾
Naming Covalent Practice
P4S5
SF6
N2O5
H2O
NF3
SiO2
P2Br4
SO3
Tetraphosphorus Pentasulfide
Sulfur Hexafluoride
Dinitrogen Pentaoxide
Dihydrogen Monoxide
Nitrogen Trifluoride
Silicon Dioxide
Diphosphorus Tetrabromide
Sulfur Trioxide
Greek Prefixes
1- mono-
2- di-
3- tri-
4- tetra-
5- penta-
6- hexa-
7- hepta-
8- octa-

39. Name the following: Mixed Practice37
Name the following: Mixed Practice
KBr
HCl
MgO
CaCl2
H2O
NO2
Potassium Bromide
Hydrogen Monochloride
Magnesium Oxide
Calcium Chloride
Dihydrogen Monoxide
Nitrogen Dioxide
Calcium Sulfide
Chromium (III) Oxide
Iron (II) Oxide
Lithium Bromide
CaS
Cr2O3
FeO
LiBr

40. End of Chemical Bonding Study Packet

41. 38
Writing & Balancing
Chemical Equations
Goals revisited

42. GOALS
1. Apply the Law of Conservation of Matter by balancing following types of chemical equations:
Synthesis
Decomposition
Single Replacement
Double Replacement
2. Demonstrate the Law of Conservation of Matter in a chemical reaction

43. 39
Chemical Reactions
A chemical reaction is a change in which one or more substances are converted into new substances.
Rearrangement of bonds in compounds and molecules.
Chemical Equations make it possible to see clearly what is happening during a chemical reaction

44. Chemical equations are a shorthand way to show chemical reactions.40
Chemical equations are a shorthand way to show chemical reactions.
Reactants
Products
H2 + O2
H2O

45. The mass of the products always equals the mass of the reactants41
Conservation of Mass
The mass of the products always equals the mass of the reactants

46. 2 Hydrogen atoms & one Oxygen atom42
H2 + O2
H2O
Does this meet the Conservation of Mass Law?
2 Hydrogen atoms
2 Oxygen atoms
2 Hydrogen atoms & one Oxygen atom
Must Balance the Equation to show Conservation of Mass.

47. 43
Can add coefficients to Balance equations. 2
H2 + O2 2
H2O 4 2 2 2 4 2 1
Balanced!!
Steps:
1. Count Atoms on both sides
2. If not Balanced, add coefficients to balance.
3. Recount atoms after adding each coefficient.
4. Keep adding coefficients until balanced.

48. Types of
Chemical
Reactions 44

49. 45
Chemical Reactions
You start with one or more compounds and turn it into different compounds.
Vapors of hydrogen chloride in a beaker and ammonia in a test tube meet to form a cloud of a new substance, ammonium chloride.

50. Single Replacement (Single Displacement) 46
Synthesis
Decomposition
Single Replacement (Single Displacement)
Double Replacement (Double Displacement)

51. Synthesis
“to make”
A + B AB
2Cu + O CuO
2H2 + O H2O

52. Decomposition 2H2O 2H2 + O2 NaOH Na + OH AB A + B “to breakdown”
Animation
Decomposition
“to breakdown”
AB A + B
2H2O H O2
NaOH Na + OH

53. When one element replaces another element in a compound
Single Replacement
When one element replaces another element in a compound
A + BC AC + B
Cu+AgNO3 Cu(NO3)2+ 2Ag
The more reactive metal will always replace the less reactive metal. (p749)

54. Single Replacement
Clip

55. Double Replacement AB + CD AD + CB Ba(NO3)2+KSO4 2KNO3 + BaSO4
Positive Ion of One compound replaces the positive ion of another compound and a Precipitate is formed.
AB + CD AD + CB
Ba(NO3)2+KSO KNO3 + BaSO4
Review Clip

56. Clip
Animation

57. Review Clip

59. Combustion

60. Chemical Reactions
and
Energy

61. Chemical Reactions and Energy51
Chemical Reactions and Energy
All chemical reactions release or absorb energy.
Heat, light, sound
Chemical reactions are the making and breaking or bonds.

62. 1. Exergonic
Chemical reactions that releases energy are called exergonic.
Glow sticks
If heat is released, it is called exothermic.

63. 2. Endergonic
Chemical reactions that require energy are called endergonic.
Ex: Cold Packs
If heat is absorbed, it is called endothermic

64. Catalysts and Inhibitors
Some reactions proceed slowly.
They can be sped up by a catalysts.
Catalysts are not used up in the reaction.
EX: enzymes (biological catalysts)
Some reactions proceed too fast.
They can be slowed down by inhibitors.
EX: Preservatives in food

65. GOALS Revisited…..
1. Compare & contrast ionic and covalent bonds in terms of electron position.
2. Predict formulas for stable binary ionic compounds based on balance of charges.
3. Use IUPAC nomenclature for transition between chemical names and chemical formulas of
• binary ionic compounds
• binary covalent compounds
4. Apply the Law of Conservation of Matter by balancing the following types of chemical equations:
• Synthesis
• Decomposition
• Single Replacement
• Double Replacement