1. Introduction to Chemicals and Safety
Section 5.2

2. Safety Skills
Read pages & and answer the following questions.
What does HHPS stand for?
What does WHMIS stand for?
Name 5 household products that are not reactive.
Why do you think we use hazardous chemicals when safer methods and materials are known?
Give an example of a hazardous substance that might be found in the
Kitchen
Bathroom
Garage
garden
Make a list of safety rules that should be followed in the lab. What sort of consequences should be put in place for students who do not follow the lab safety rules?

3. Assignment: Label the HHPS and WHMIS symbols. (handout)
Read and sign the Lab Safety Contract and have it signed by a parent/guardian.
Lab Safety Worksheet

4. HHPS (Hazardous Household Products Symbols)

5. WHMIS (Workplace Hazardous Materials Information System)

6. Unit 2
Chemical Processes

7. WHAT IS CHEMISTRY? Chemistry Matter
Is the study of matter, its properties, and its changes or transformations.
Matter
Is anything that has mass and takes up space.
Everything around us is made up of matter (water, sand, DNA, oxygen, etc)

8. States of Matter There are four states of matter. They are:
Solid (s)
Liquid (l)
Gas (g)
Plasma
does not have a definite shape or volume unless enclosed in a container
it may form structures such as filaments, beams and double layers
some common plasmas are stars and neon signs.
As you move from solid to a gas some of the atoms spread apart and the matter gains energy

9. Physical and Chemical Properties Song
All matter has physical and chemical properties.
Physical Properties of chemicals describe the way they look. Examples include: colour, shine, strength, and state. A change in the way something looks is called a physical change. These include:
Evaporation (liquid to a gas)
Freezing (liquid to a solid)
Sublimation (solid to a gas)
Melting (solid to a liquid)
Condensation (gas to a liquid)

10. A physical change is not permanent
A physical change is not permanent. It is easily reversed and does not change the chemical formula or properties of the matter.
Chemical Properties are the behaviors that happen when a substance changes into a new one. The change itself is called a chemical change.

11. Reactants are the starting materials in a reaction.
Products are the new materials made in the reaction.
Reactants Products
Physical & Chemical Changes Song H2 O2
H2O

12. Assignment:
Complete the handout “Physical or Chemical Change?”

13. Physical or Chemical Changes?
Type of Change?
1. The steel wool turned the block pot shiny silver.
2. Lightning flashed across the sky.
3. Acid caused limestone to fizz.
4. The back steps are rotting out.
5. The piece of lemon turned the tea cloudy.
6. Spilt gasoline dried, but left a bad odor in the room.
7. The gravy in the refrigerator jelled.
8. A blue-green lobster turned red as it cooked.
9. A cavity appeared in a tooth.
10. Flash cameras flashed after the wedding photo.

14. Chemicals and Chemical Changes
Section 5.1

15. Heterogeneous Mixtures
Classifying Matter
Matter
Pure Substances
Compounds
Elements
Mixtures
Homogeneous Mixtures
(Solutions)
Heterogeneous Mixtures

16. Definitions Pure Substances Elements Compounds
All the particles are the same. The substance has constant properties.
Elements
Are pure substances that are made up of only 1 type of atom. They can be identified with a chemical symbol.
Ex. Carbon = C
Compounds
Are pure substances that contain two or more different elements in fixed amounts. They can be written as a chemical formula. Ex. Water = H2O

17. Mixtures Video Clip Mixtures Homogeneous Mixtures
Are a combination of two or more substances that are not chemically combined. They are simply in the same place at the same time.
Homogeneous Mixtures
Mixtures are uniform. It’s parts are not distinguishable. Ex. Pop, salt water.
Heterogeneous Mixtures
Mixtures are not uniform. It’s parts are easily distinguishable. Ex. Pizza, salad, sand and water.

18. A Chemical Reaction usually occurs when:
Color Change occurs
Gas is formed (bubbles)
Precipitate
A solid that forms in a liquid
Energy change
Exothermic – heat/light is given off (feels warm)
Endothermic – heat is absorbed (feels cold)
Very difficult to reverse

19. Assignment:
Read pages and complete questions 1, 2, 4, 5 on page 175.
Complete the handout “Elements, Compounds and Mixtures”

20. Atoms and Elements

21. The Atom
The atom is the smallest particle of an element that retains the same properties of that element (Democritus, 4th century BCE).
Atoms are the basis for
everything in the Universe.
Everything is made of matter
and all matter is made of atoms.

22. Dalton’s Atomic Theory
John Dalton (1700) stated the following ideas:
All elements are composed of tiny indivisible particles
called atoms.
Atoms of the same element are identical.
The atoms of any one element are different
from those of any other element.
Atoms of different elements can combine with
each other in simple whole number ratios to
form compounds.
Chemical reactions occur when atoms are separated,
joined, or rearranged. However, atoms of one element
are not changed into atoms of another by a chemical reaction.

23. Elements and Atomic Structure
The Bohr-Rutherford model of the atoms suggests that atoms are composed of three types of subatomic particles:
Protons: heavy, positively charged particles found in the core of the atom called a nucleus.
Neutrons: heavy, neutral particles found in the nucleus.
Electrons: light, negatively charged particles with almost no mass that circle the nucleus at different energy levels called orbits or shells.
Particle
Symbol
Charge
Approximate Mass
Proton p+ 1+ 1
Neutron n0
Electron e- 1-
1/1840

24. Video Clip on Atomic Structure

25. Atomic Structure
Atomic Number is the number of protons. This number is constant, it does not change.
Mass Number is the number of protons and neutrons. The number of neutrons can change, so the mass number is not constant.
Standard Atomic Notation
for an element shows it’s mass number
at the top, it’s chemical symbol
and it’s atomic number at the bottom.
Since atoms are neutral: # of electrons = # of protons.
Therefore:
mass number – atomic number = # of neutrons 12 C 6

26. Chemical Symbols
Each element has a one or two letter symbol (a few have three). The first letter is always capitalized, and the second it is always lowercase.
For many, it is the first letter or two of their name.
Some symbols came from the element’s name in Latin, so the letters don’t always match the name.
Lead = Plumbum = Pb

27. You are responsible for knowing the name and the symbol for the following elements (spelling counts):
Name H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Mn Sb Zn Rn Co I Br Ra Fe Xe Kr Pb Ni W Ag Hg Cu Pt Sn Au

28. Assignment: Subatomic Particles Worksheet (handout)
Quiz on the 40 elements (names/symbols)
The Elements Song

29. Subatomic Particles Worksheet (Blackline Master 3.3)
Element
Name
Atomic Number
Mass Number
Standard
Atomic
Notation
# of p
# of e
# of n
Aluminum
40 Ar 18 4 5 40
Carbon 19 2

30. Element
Name
Atomic Number
Mass Number
Standard
Atomic
Notation
# of p
# of e
# of n 3 12
Neon 7 8 31 19 14 16
Sodium

31. The Periodic Table Section 5.5
The Elements Song-Animated Periodic Table

32. History of the Periodic Table
The periodic table was first arranged by Mendeleev according to mass number.
Mosely (1913), a student of Mendeleev, decided to try to arrange the periodic table according to atomic number. This method worked better and it is how the periodic table is arranged today.

33. Each element behaves in a different way from all other elements because it has different properties.
However, some behave similarly so they are grouped together and arranged in rows and columns.

34. Horizontal rows (across) are called periods. There are 7 of them
Horizontal rows (across) are called periods. There are 7 of them. Elements are arranged by increasing atomic number.
A column (down) is known as a group. Elements in a family group share similar chemical properties.

35. Group A Elements Representative elements. For Group A elements ONLY,
Group 1A - Alkali metals react with water to form basic or alkaline solutions.
Group 2A - Alkaline Earth Metals react with oxygen to then react with water to form alkaline solutions.
Group 7A – Halogens combine with other elements to form salts.
Groups 8A - Noble Gases or Inert Gases are chemically unreactive.
For Group A elements ONLY,
if you know the group #, then
you know the number of electrons
in the outer shell or orbit.

36. Group B Elements Transition elements.
Found in the middle of the periodic table.
The group number does not indicate the number of electrons in the outer shell.

37. Metals
Nonmetals
Elements found to the left of the “staircase” are metals. Common properties include:
Shiny
Conduct electricity
Melt when heated
Malleable (can be beaten into thin sheets)
Tend to lose electrons
Ductile (can be formed into wire)
Elements found to the right of the “staircase” are nonmetals. Common properties include:
Not shiny
Not malleable
Low melting point
Poor conductor of electricity
Brittle
Tend to gain electrons

38. Metalloids
Elements found touching either side of the staircase are called metalloids or semi-metals. Have properties of both metals and nonmetals. Common properties include:
- Solids
- Can be shiny or dull
- Ductile
- Malleable
- Conduct electricity better than nonmetals but not as good as metals.

39. Assignment:
Color coding the periodic table.

41. Elements 43, 61, and from 84 on are all radioactive.
Elements 43, 61, and from 93 on have been made by man (synthetic).
There are 11 gases:
Hydrogen, Helium, Nitrogen, Oxygen, Fluorine, Neon, Chlorine, Argon, Krypton, Xenon, Radon
There are 2 liquids:
Bromine, Mercury
Inner Transition Elements
Lanthanide Series – Begins at Cerium #58
Actinide Series – Begins at Throrium #90
Blank spots on the table are for elements that have not
been made yet.

42. Lanthanide Series
Actinide Series

43. Electron Orbits and Energy Levels
Electrons are arranged in orbits around the nucleus of the atom. The orbits are also known as energy levels.
Orbits or Energy Levels
Number of Electrons/Level 1 2 8 3 4 18 5 6 32

44. Drawing Bohr-Rutherford Diagrams
Electron orbits are filled from the inside out. (The orbits closest to the nucleus fill first.)
Atoms are stable (do not react with other atoms) when their outermost orbit is full.

45. Assignment: Complete the sheet on Bohr-Rutherford Diagrams.
Fill in the worksheet with Bohr diagrams. (We are only going to look at the first 20 elements)
When you are finished see if you can observe any patterns in the organization of the table.

46. BOHR-RUTHEFORD

47. Lewis Dot Structures
Lewis Dot Structure shows only the valence electron arrangement for each element.
Valence electrons are electrons that are in the outermost energy level only. These are the electrons that will determine the ionic charge of an atom.
To draw Lewis dot diagrams write the symbol for the element, then draw dots around the symbol to represent the number of valence electrons. Electrons are placed one on each side going around the symbol.

48. Lewis Dot:
Oxygen Fluorine

49. HOMEWORK Draw the Lewis Dot Structure for the first 20 elements
List all elements that are stable.
List all elements that have only 1 electron in their outer level.
List all elements that are only 1 electron short of a full outer level.
Look for patterns on your periodic table.

50. LEWIS DOT

51. Isotopes
Isotopes are atoms that have the same number of protons, but different numbers of neutrons. Therefore, they also have different mass numbers and different atomic masses. Isotopes are still chemically alike because they have identical numbers of protons and electrons.
Ex. Three Isotopes of Carbon are Carbon-12, Carbon-13, and Carbon-14. Complete the table for each.
Symbol
12C 6
13C
14C
# neutrons
#protons
# electrons

52. Ions
Normally, atoms have no charge because the number of protons (+) equals the number of electrons (-). When forming a compound, an atom can gain or lose 1 or more electrons.
Ion - an atom which has a positive or negative charge.
Remember - atoms want a full outer orbit, so they tend to lose or gain electrons to do this.

53. There are two kinds of ions: Cations
An atom or group of atoms with a positive charge.
Compared to a neutral atom, cations have less electrons.
Metals tend to lose electrons and form cations.
Anions
An atom or group of atoms with a negative charge.
Compared to a neutral atom, anions have more electrons.
Nonmetals tend to gain electrons and form anions.
Names of anions end in “ide”

54. Using the Periodic Table to Find Ionic Charges
Metals: groups IA (1A), IIA (2A), and IIIA (3A)
Lose electrons when they form ions. Their ionic charge is positive (cations). Their charge equals the group number.
Ex. Group 1A = 1+ Li+ Group 2A = 2+ Mg2+ Group 3A = 3+ Al3+
Elements in groups 4A and 8A usually don’t form ions

55. Nonmetals: groups 5A, 6A, and 7A
Gain electrons when they form ions (anions). The charge is found by subtracting the 8 from the group number.
Ex. Group 5A = = 3- N 3- Group 6A = = 2- O 2- Group 7A = = 1- F-

56. 3. Transition Metals (Group B)
These charges can’t be read from the periodic table. These metals can have more than one charge, so a roman numeral must be used to indicate which ion is used.
Ex. Iron (II) – Fe2+ and Iron (III) – Fe3+
Lead (II) – Pb2+ and Lead (IV) – Pb4+

57. Ionic Compounds
Section 5.8

58. Ionic Compounds Binary Ionic Compounds
Only contain two elements
Two ions held together by electrical attraction.
Made of a cation (metal) and a anion (nonmetal)

59. Writing Formulas for Ionic Compounds
Write the symbol for your metal and nonmetal
Write the charges for each
Cross-over the charges from top to bottom
Remove the charges
Simplify the numbers and remove the 1’s

60. Sodium Chloride
Calcium Oxide
Calcium Phosphide
Mercury (II) Nitride

61. Naming Ionic Compounds
Name the cation first!
Figure out if you need a Roman Numeral next to the cation (Group B Metal)
Name the anion (ending in the suffix “ide”)

62. Formula
Name
LiF
Lithium Flouride
NaCl
CaCl2
Al2O3
CuBr2

63. Polyatomic Ions
Section 5.9

64. Polyatomic Compounds Ternary Ionic Compounds
Contain three or more elements
Joined by an electrical attraction
Pure substances that join:
A metal with a polyatomic ion
A polyatomic ion with a nonmetal
Two polyatomic ions
Polyatomic ions
Groups of atoms that stay together and carry an overall charge

65. Common Polyatomic Ions and their Charges
Formula
Charge
Hydroxide OH 1-
Nitrate
NO3
Chlorate
ClO3
Bicarbonate
HCO3
Carbonate
CO3 2-
Sulfate
SO4
Phosphate
PO4 3-
Ammonium
NH4 1+
Fluorate
FO3
Iodate
IO3
Bromate
BrO3

66. Writing Polyatomic Formulas
Following the same rules as for Ionic Compounds we….
Write the symbol for your metal and polyatomic ion
Write the charges for each
Cross-over the charges from top to bottom
(if needed)
Remove the charges
Simplify the numbers and remove the 1’s
When you are adding a subscript to a polyatomic ion you must put brackets around it.

67. Lithium nitrate
Sodium sulfate
Calcium carbonate
Ammonium chlorate
Aluminum hydroxide

68. Naming Polyatomic Compounds
Name the cation and the polyatomic ion
Cation is always named first!
Figure out if you need a Roman Numeral in the name
Which Roman Numeral
Check your work

69. Formula
Name
K2CO3
Na2SO4
AgNO3
Ni2(CO3)3

70. Comparing Ionic Compounds
Formulas
2 elements
(metal + nonmetal)
Names
Name always has an
–ide ending
Polyatomic Compounds
Formulas
3 or more elements
(Metal + polyatomic ion)
If there are brackets it is automatically a polyatomic compound
Names
Name always has an –ate or
–ite ending
EXCEPTIONS
OH – hydroxide
CN - cyanide

71. Molecular Compounds
Section 5.10

72. Molecular Compounds Molecular Compounds Covalent Bonds
Made of two nonmetals and have no charge.
Formed using covalent bonds.
Covalent Bonds
Shared pairs of electrons between two nonmetal elements that hold the atom together.

73. Diatomic Elements I2, H2, N2, Br2, O2, Cl2, F2
If these elements are alone they pair up with another of their kind.
They are written with the subscript 2 (H2)
Share electrons in a covalent bond instead of existing alone by themselves.
These elements are known as the Magic 7 (they make the shape of a 7 on the periodic table…except for H)
I Have No Bright Or Clever Friends.
I2, H2, N2, Br2, O2, Cl2, F2

74. Naming Molecular Compounds
The first element in the name is named using the full name
The second element is named using the ending –ide.
Prefixes are used for each element to indicate the number of atoms present.
The prefix –mono is only used to name the second element
CO = Carbon monoxide NOT
Monocarbon monoxide.

75. Prefixes in Molecular Compounds
Number
Example
Mono- 1
Carbon monoxide CO
Di- 2
Carbon dioxide
CO2
Tri- 3
Sulfur trioxide
SO3
Tetra- 4
Carbon tetrafluoride
CF4
Penta- 5
Phosphorus pentabromide
PBr5
Hexa- 6
Sulfur hexafluoride
SF6
Hepta- 7
Dichlorine heptoxide
Cl2O7
Octa- 8
Trichlorine octoxide
Cl3O8
Nona- 9
Nonaphosphorus dichloride
P9Cl2
Deca- 10
Tetraphosphorus decoxide
P4O10

76. Many molecular compounds have both common names and scientific names.
Water
Dihydrogen monoxide
H2O
Ammonia
Nitrogen trihydride
NH3
Ozone
Trioxide O3

77. Writing Formulas for Molecular Compounds
Write the symbol for each element
Write the subscript from the prefix (no criss cross)
Do not write 1’s.
Do not reduce any numbers.

78. Try the Following: Dinitrogen tetrahydride Heptafluorine hexoxide
Se8Cl3
S9I

79. Characteristics of Ionic and Molecular Compounds
Ionic Compounds
High melting and boiling points
Usually solid at room temperature
Will dissolve in water
Conduct electricity
Atoms held together by strong bonds
Made of metals and nonmetals
Examples:
NaCl, AgI, Fe2O3
Molecular Compounds
Low melting and boiling points
Many are gases at room temperature
Will not dissolve in water, but can dissolve in non-aqueous solutions (don’t contain H2O)
Cannot conduct electricity
Weak bonds between molecules
Made of nonmetals
Examples:
H2O, CO2, CI2

80. Naming Compounds Review
What is it made of?
2 elements
(metal + nonmetal)
Name the metal
Name the nonmetal (“ide” ending)
Do you need Roman Numerals?
3 or more elements
Name the polyatomic ion (“ate or “ite” ending)
(nonmetals)
Put prefixes in front of the names to tell how many atoms there are.
*No Mono in the first name
The second name ends in “ide”

81. Writing Formulas Review
What is it made of?
2 elements
(metal + nonmetal)
“ide” ending
1. Write the symbols
2. Write the Charges
3. Cross the Charges
4. Remove the Charges
5. Reduce and Remove 1’s
3 or more elements
“ate or ite” ending
*Brackets if necessary
(nonmetals)
1. Write the symbol
2. Write the subscript for the prefix.
3. Repeat for the next element.

82. Word Equations
Section 6.1

83. Word Equations
A word equation is one way of representing a chemical reaction. It tells you what reacts and what is produced. Word equations are written like this:
Reactants -> Products
When hot steel wool (iron) is put into a bottle of oxygen, there is a spectacular reaction and iron (III) oxide is produced. The word equation would be:
Iron + Oxygen -> Iron (III) oxide

84. Try These…
Write the word equations for each of the following examples:
A coil of copper wire is placed in a glass of colorless silver nitrate solution. Silver ions form on the coil, and the solution turns into copper (II) nitrate (which is blue).
Hydrogen and zinc chloride are produced when zinc is added to hydrochloric acid.

85. Assignment
Do questions 1, 2, 3 & 4 on page 219

86. Conserving Mass
Section 6.3

87. Conserving Mass
The Law of Conservation of Mass states that in a chemical reaction, the total mass of the reactants is always equal to the total mass of the products.
Example 1:
Methane oxygen -> water carbon dioxide
The mass of the methane and the oxygen combined should equal the mass of the water and the carbon dioxide.
If methane = 20g, oxygen = 30 g, and water = 40g what is the mass of the carbon dioxide?

89. Example 2:
Solution A has a mass of 50g. Solution B has a mass of 30g. When they are mixed together a chemical reaction occurs and a gas and a liquid are produced. If the mass of the final liquid is 60g, then how much gas was produced?

90. Assignment
Do questions 1-5 on page 223.

91. Counting Atoms Subscripts Coefficients Brackets C N2 3C 2N2H4
Al2(SO4)3
3Ca3(PO4)2

92. Complete the examples below:
CaF2
NH4NO3
Type of Atom
# of Atoms
Type of Atom
# of Atoms
2CaF2
3Be(OH)2
Type of Atom
# of Atoms
Type of Atom
# of Atoms

93. Balancing Chemical Equations
Section 6.5

94. Skeleton Equations
A skeleton equation represents all chemicals by their formulas. It is NOT balanced. Skeleton equations and chemical equations are universal. They can be understood in every language all over the world!
Word Equation:
Methane oxygen -> water carbon dioxide
Skeleton Equation:
CH4 + O2 -> H2O + CO2

95. Skeleton Equations
A skeleton equation is NOT balanced. These equations are universal and can be understood all over the world!
Word Equation:
Methane oxygen -> water carbon dioxide
Skeleton Equation:
CH4 + O2 -> H2O + CO2

96. Balanced Chemical Equations
In a balanced chemical equation the reactants and the products are equal in the number of atoms. (This follows the Law of Conservation of Mass)

97. Rules: 1. Begin balancing with the most complicated-looking group.
A polyatomic ion, or the group that has the most elements and numbers in it.
2. Lock this in with a coefficient of 1 and balance any uneven
elements.
3. Save any single elements in the reactant and products for last.
Especially if it is hydrogen or oxygen.
4. If you get stuck, double the most complicated-looking group
and try again.
5. Make sure that all coefficients are in the lowest-possible ratio.
Know when to quit! None of the reactions you will encounter will be that difficult. If the coefficients are getting wild, double-check what you’ve done since you may have a simple mistake.

98. Rules: 1. Begin balancing with the most complicated-looking group.
A polyatomic ion, or the group that has the most elements and numbers in it.
2. Lock this in with a coefficient of 1 and balance any uneven
elements.
3. Save any single elements in the reactant and products for last.
Especially if it is hydrogen or oxygen.
4. If you get stuck, double the most complicated-looking group
and try again.
5. Make sure that all coefficients are in the lowest-possible ratio.

99. Balancing Examples H2 + N2 -> NH3 Type of Atom Reactants Products H

100. Balancing Examples NaCl + BeF2 -> NaF + BeCl2 Type of Atom
Reactants
Products Na Be Cl F

101. Balancing Examples K2O + H2O -> KOH K2O + HOH -> KOH
Type of Atom
Reactants
Products K O H

102. Balancing Examples FeCl3 + Be3(PO4)2 -> BeCl2 + FePO4 Type of Atom
Reactants
Products Fe Be Cl
PO4

103. Try These: Mg + HNO3 -> H2 + Mg(NO3)2 Zn + HCl -> ZnCl2 + H2
Al Fe3N > AlN Fe
Fe O >   Fe2O3
KClO3 ->   KCl   O2
O C6H12O6 -> H2O   CO2
S F >   SF6

104. Assignment:
Practice sheet balancing equations.

105. Types of Reactions
Section 6.6

106. There are five main categories of chemical reactions:
Combustion
Synthesis
Decomposition
Single Displacement
Double Displacement

107. Combustion
Combustion is the very rapid reaction of a substance with oxygen to produce compounds called oxides (CO2 and H2O). We often call this process burning.
The fuel can be a variety of things but it is often a hydrocarbon such as gasoline, natural gas, kerosene, or candles.

108. Complete Combustion The formula for a complete combustion is:
Fuel oxygen oxides
CX HX O > CO H2O energy
Reaction when you light a lighter:
Butane oxygen -> carbon water energy
dioxide
2C4H O > 8CO H2O energy

109. Complete Combustion The formula for combustion of a hydrocarbon is:
CX HX O > CO H2O energy
When you light a lighter the following reaction occurs:
Butane oxygen -> carbon water energy
dioxide
2C4H O > 8CO H2O energy

110. Incomplete Combustion
The formula for Incomplete combustion occurs when there is not enough oxygen:
CXHX O2 -> CO CO C H2O + energy
Carbon monoxide is an odourless, colourless gas that is extremely poisonous. It blocks the bodies ability to use oxygen, suffocating it and causing death.

111. Incomplete Combustion
Incomplete combustion occurs when there is not enough oxygen. The formula is:
CXHX O2 -> CO CO C H2O + energy
Carbon monoxide is an odourless, colourless gas that is extremely poisonous. It blocks the bodies ability to use oxygen, suffocating it and causing death.

112. Synthesis and Decomposition
Section 6.7

113. Synthesis
Synthesis Reactions involve the combination of smaller atoms or compounds into larger compounds. They are also known as combination reactions. They have the following general formula:
A + B ➜ AB
If both reactants are element, then the reaction MUST be synthesis.
Eg. A B AB
2H2 + O2 ➜ 2H2O

114. Synthesis reactions sometimes involve joining two compounds into a larger one.
Eg.
Hydrogen ammonia ➜ ammonia
chloride chloride
HCl NH3 ➜ NH4Cl

115. Decomposition
Decomposition Reactions involve the splitting of a large compound into smaller molecules or elements. They have the following general formula:
AB ➜ A + B
If there is only 1 reactant, then the reaction MUST be decomposition.
Eg. AB A B
2H2O ➜ 2H2 + O2

116. Single and Double Replacement
Section 6.9

117. Single Replacement/Displacement
Single Displacement - The reactants MUST be an element and a compound. The general formula is:
A + BC ➜ B + AC
Single displacement can involve metals:
Eg. Na + KCl ➜ K + NaCl
Single displacement can involve nonmetals:
Eg. F2 + LiCl ➜ LiF + Cl2
Remember: If the single element is a metal it will replace the metal.
If the single element is a nonmetal it will replace the nonmetal.

118. Single Replacement/Displacement
Single Displacement involves one element replacing another in a chemical reaction. The reactants MUST be an element and a compound. The general formula is:
A + BC ➜ B + AC
Single displacement can involve metals:
Eg. Na + KCl ➜ K + NaCl
Single displacement can involve nonmetals:
Eg. F2 + LiCl ➜ LiF + Cl2
Remember: If the single element is a metal it will replace the metal.
If the single element is a nonmetal it will replace the nonmetal.

119. Double Replacement/Displacement
Double Displacement involves two elements replacing one another. The reactants must be two ionic compounds. The positive ions stay in the same position (A and C) and the negative ions change partners (B and D). The general formula is:
AB + CD ➜ AD + CB
Eg.
NaOH + FeCl3 ➜ Fe +3 OH Na +1 Cl -1
When you write the two new compounds formed, you must cross charges to get the correct new formulas.
NaOH + FeCl3 ➜ Fe(OH)3 + NaCl

120. Double Replacement/Displacement
Double Displacement - The reactants are two compounds. Two elements replace one another to create new products.
AB + CD ➜ AD + CB
Eg.
NaOH + FeCl3 ➜ Fe(OH)3 + NaCl
When you write the two new compounds formed, you must cross charges to get the correct new formulas.

121. synthesis

122. Reaction Types Review:
Questions to Ask:
One Reactant?
Two Elements?
A Hydrocarbon and oxygen?
An Element and a compound?
Two Compounds
Types of Reaction
Decomposition
Synthesis
Combustion
Single Displacement
Double Displacement
Yes No
Yes No
Yes No
Yes No
Yes

123. Hydrocarbons
Section 5.12

124. Organic compounds are molecular substances that contain carbon and hydrogen atoms as building blocks. Methane is the simplest organic compound (CH4) and there are tens of thousands of others.
Carbon is the element that is most commonly found in molecules. That is because of it’s great combining capacity. Carbon has four lone valence electrons so it would like to bond or share with 4 other atoms. This is the maximum number of bonds a single atom can make!

125. (Methane - CH4 , butane - C4H10 , propane - C5H12)
Most foods, and most of our bodies, are made up of organic compounds. The two main sources of organic compounds are:
The ones that plants produce during photosynthesis including sugars, carbohydrates, and fats.
Fossil fuels found underground that were made millions of years ago from decaying plant and animal material.
Hydrocarbons are compounds with only hydrogen and carbon in them. They are also known as fossil fuels.
(Methane - CH4 , butane - C4H10 , propane - C5H12)

126. Synthetic Substances
Section 5.13

127. Chemical Wizardry: Synthetic Substances
Natural products are obtained from natural sources like animals, plants, and minerals. Leather, wool, and cotton are examples.
Synthetic products are made by man. Lycra, kevlar, and vulcanized rubber are just a few examples.
Polymers are long thin molecules formed from smaller molecules linked together. They have very different properties than the molecules they are formed from. There are natural and artificial polymers. Plastics are made from polymers and so are starches found in potatoes.

128. Acids and Bases
Section 8.2 & 8.3

129. Acids and bases are molecular compounds that break into ions in water.
An acid: - tastes sour
- reacts with some metals to produce hydrogen gas
- is a good conductor of electricity
- reacts with a base to produce salt and water
Common acids include:
Common Name
Found in
vinegar (acetic acid)
salad dressing
citric acid
oranges and lemons
ascorbic acid
vitamin C
lactic acid
sour milk
carbonic acid
carbonated drinks, pop
acetylsalicylic acid
aspirin
sulfuric acid
car batteries
hydrochloric acid
your stomach for digestion

130. - is a good conductor of electricity
A base: - tastes bitter
- feels slippery
- is a good conductor of electricity
- reacts with an acid to produce salt and water
Common bases include:
Common Name
Found in
sodium hydroxide
drain cleaner
potassium hydroxide
soap, cosmetics
aluminum hydroxide
antacids (Tums)
ammonium hydroxide
window cleaner
sodium bicarbonate
baking soda
potassium sulfite
food preservatives

131. Bases react with proteins to break them into smaller molecules
Bases react with proteins to break them into smaller molecules. A hair clogged drain may be cleared by treating it with a drain cleaner that contains sodium hydroxide.
Indicators are substances that turn different colours to indicate the presence of acids or bases. Litmus paper is one example of an indicator. Phenolphthalein is another popular indicator.
Chemists use the pH scale to represent how acidic or basic a solution is. It was developed in 1909 to allow scientists to compare the hydrogen ion concentration in various solutions. (pH stands for percent Hydrogen) The scale goes from 1 to 14. Most acids and bases can be ranked on this scale.

132. pH 1. 0 Highly Acidic pH 1. 5 stomach acid pH 2. 0 lemons pH 3
pH 1.0 Highly Acidic pH 1.5 stomach acid pH 2.0 lemons pH 3.0 apples pH 5.6 normal rain pH 6.5 milk pH 7.0 Neutral pH 8.2 baking soda pH 10.0 milk of magnesia pH 12.4 bleach pH 13.5 drain cleaner pH 14.0 Highly Basic

133. http://www. brainpopjr