1. Science 10 Introduction

2. Science 10 Outline Resource text: Review of Science 10 Chemistry
Science and Safety
Classifying matter
Atomic structure
Chemical names and formula
Chemical reactions

3. Evaluation: Course work Tests & Quizzes, Labs 40%
Assignments and labs 40%
Final Exam 20%
Marks may be deducted for work handed in late. Time factor will be addressed.

4. Lab Safety Common Sense! No fooling around Wear proper clothing
Use equipment correctly
No unauthorized experiments
Hand chemicals with respect - never taste
Clean up - wash/return/turn off

5. IUPAC – International Union of Pure and Applied Chemistry
TERMS
IUPAC – International Union of Pure and Applied Chemistry
MSDS – Materials Safety Data Sheet

6. WHMIS Provides: 1) information 2) safety & training 3) labeling
Workplace Hazardous Materials Information System
Provides:
1) information
2) safety & training
3) labeling

7. WHMIS SYMBOLS
Class A
Compressed Gas
ie) acetylene cylinder (welding)

8. Class B
Flammable & Combustible Material
ie) methane

9. Class C
Oxidizing Material
ie) Acid

10. Class D1 Materials causing immediate & serious toxic effects
ie) Acids & Bases

11. Class D2
Materials causing other toxic effects.
ie) Heavy metals (Pb, Hg)

12. viruses, biological weapons
Class D3
Biohazardous Infectious material
viruses, biological weapons

13. Class E
Corrosive Material
ie) Acids & Bases

14. Class F
Dangerously Reactive Material
ie) sodium

15. The Scientific
Process

16. Questioning
Problem statement: What affect does the (manipulated) have on the (responding)
Manipulated variable: changed
Responding variable: reacts
Controlled variable: stays the same

17. Proposing Ideas
Background info: Info you know from class and experience
Hypothesis: I hypothesize that (answer to problem) because (reason based on background)
Prediction: statement using values

18. Designing Experiments
Overview: Design in 2-3 sentences.
Experimental group vs control group - one variable changed
Pre-lab calculations
Materials & procedure:steps

19. Observing & Measuring Data
Observations: qualitative vs quantitative Be specific & detailed
Record Data: - in tables & charts. Remember a title.

20. Processing Data
Analysis: answer any questions, complete calculations or graph
Graphing: follow the rules outlined in the booklet. Remember a title & use pencil. The 3 types of graphs are pie , bar & linear(straight, sigmoid, exponential, or curved).

21. Interpreting Data Evaluation: 1) Hypoth/pred/purp supported?
2) Experimental design?
3) Errors? - human or equipment
4) Conclusions
Extension: Other related experiments.
Bibliography/References

22. Answers to Sig. Digs W/S 1. a. 3 2. a. 5.808 x 103 3.a. 60
b. 3 b. 6.3 x b. 6200
c. 2 c x c. 740
d. 2 d x d
e. 3 e. 7.0 x e
f. 2 f. 5.8 x f. 4.3

23. Answers to Sig. Digs W/S 4. a. 400.9 g b. 333.5 J c. 83 m d. 57 kg
e g
f x 104 s
g. 4 x 105 g
h. 8.1 x 1010 i

24. Answers to Sig. Digs W/S 5.
Slope = 2.11 mmol/L/s

25. Review of Science 10

26. Organization

27. Some Additional Definitions
Matter - object that has mass & occupies space
Pure Substance - substance that has no impurities
Mixture - made of 2 or more substances

28. Compound - two or more elements in fixed proportions
Element - chemical building block; an atom table.

29. Solution - a uniform mixture where the parts are not visible
(Also called a homogeneous mix)
ie) salt water
The two parts to a solution are:
Solvent-large dissolving part-often H2O
Solute - smaller dissolved part-salt

30. Mechanical mixture: non-uniform mixture where the parts remain visible and intact.
Also called heterogeneous mixture
ie) oil and water

31. Physical Properties of Matter Color Malleability Lustre Shape State
Heat capacity
Density
Ductility
Malleability
Lustre
State Bp Mp
Solubility

32. Chemical Properties 1) precipitate 2) color change 3) odour formed
describe how a substance reacts
5 evidences of chemical change
1) precipitate
2) color change
3) odour formed
4) energy change
5) gas formed

33. Examples of Physical Change
1) dissolving
2) phase change
3) dividing up

34. Examples of chemical changes - reaction where
1) a new gas is released
2) new solid is formed
3) new odour is produced
4) energy change occurs
5) new color appears

35. The Periodic Table

37. Metals, metalloids and nonmetals

39. Information on the table
Name naturally occurring elements
Family or Group - 18 columns, each with similar properties
Period - 7 rows, each with same electron orbitals
Atomic number - # or protons
Atomic molar mass - ave mass
physical properties - bp, mp, etc

40. An Introduction to Chemistry

41. Atomic Structure
An atom is the smallest part of an element that retains the properties of that element. 
The modern atomic theory was developed based on the work of the many scientists.
Greek philosophers in the centuries before Christ’s birth believed that all matter was composed of tiny, indivisible particles that they called atoms.
Their ideas prevailed for almost 2000 years until the scientific revolution of the 19th century.

42. Dalton’s Atomic Theory
Dalton’s proposal was based on many experimental observations by him and other scientists when elements react and compounds are formed.
Dalton’s Atomic Theory states:
Atoms are tiny, indivisible particles of elements
All elements are composed of atoms
Atoms of the same element are identical, but the atoms of different elements are different

43. Atoms can combine in fixed ratios to form compounds.
Chemical reactions occur by a rearrangement of atoms NOT by changing atoms.
Dalton basically said that an atom is a solid sphere similar to a billiard ball:

44. J. J. Thompson (1897)
Thompson used a device called a cathode ray tube. This is essentially a vacuum tube containing a gas at low pressure.
He found evidence that a stream of negatively charged particles was produced from the atoms in the tube. Thompson called these negative particles electrons.
Shortly after electrons were discovered, scientists carried out experiments with hydrogen and discovered positively charged particles that they called protons.

45. Thompson’s atomic theory uses the following model:
Thompson said that an atom is a sphere with embedded electrons, similar to a raisin bun.
Positive sphere Negative electrons

46. Rutherford’s Atomic Theory
Rutherford experimented using gold foil and small positively particles called alpha particles.
He shot the alpha particles at a thin piece of gold foil. In his experiments, he found that the majority of the alpha particles passed straight through the gold foil, but some bounced back.
Based on this evidence, he concluded that there is a very dense region within the gold foil that was causing the deflection of the alpha particles.

47. Since like charges repel, Rutherford concluded that this small dense region must be positively charged.
He called this region the nucleus of the atom and suggested that it was surrounded by mostly empty space containing the electrons. The positively charged particles in the nucleus he called protons.
Later experiments (Chadwick, 1932) showed that the nucleus also contains dense, neutral particles that were called neutrons.

48. Bohr’s Atomic Theory
Bohr postulated a new revised atomic theory stating that electrons have specific energies and can only occupy certain areas around the nucleus
Bohr’s atomic theory is summarized as follows:
Electrons are found at different energy levels around the nucleus
Electrons closest to the nucleus have the least amount of energy
The higher the energy level of an electron, the further it is from the nucleus 

49. Each energy level can hold a maximum number of electrons:
Electrons can move from one energy level to another by gaining or losing energy, but cannot exist between energy levels
Each energy level can hold a maximum number of electrons:
The first energy level has a maximum of 2 electrons
The second energy level has a maximum of 8 electrons
The third energy level has a maximum of 8 electrons

50. An atom with the maximum number of electrons in its outermost energy level is stable and is therefore unreactive
Bohr’s atomic theory enables us to explain a great deal about the properties of elements and how they behave, particularly how and why they react in the way they do.
It also explains the groupings of elements in the periodic table and periodicity.

51. It is the arrangement of electrons at different energy levels that determines all these things, especially the electrons in the outermost energy level.
The Bohr model will be used primarily in this course to explain chemical properties of matter.

52. Some Bohr atom models:

53. Quantum Mechanical Theory of Atomic Structure
Although the Bohr model is useful in allowing us to explain matter and its properties, it is really an over simplification.
In reality the structure of an atom is much more complex and mathematical, particularly with regard to the position of electrons in their orbits around the nucleus.
This is referred to as Quantum mechanics.

54. Quantum Theory

55. Structure of an atom
All atoms are composed of protons, neutrons and electrons.
The nucleus of an atom is composed of positively charged protons and neutrons, which are neutral. This contributes to the mass of the atom.
The electrons are found revolving in regions of space around the nucleus.

56. The nucleus is very small in comparison to the area occupied by the electrons.
The usual analogy is that of an ant, representing the nucleus, at the centre of a football field, which represents the total size of the atom.

57. Summary of Atomic Structure

58. Atomic number
This number is assigned based on number of protons in the nucleus.
fluorine has 9 protons in its nuclei, the atomic number is 9
oxygen has 8 protons in its nuclei, the atomic number is 8
Atoms are electrically neutral, which means the number of electrons is equal to the number of protons.
Therefore, the atomic number of an element is equal to both the number of protons and the number of electrons.

59. Mass Number
Is the total number of protons and neutrons in the nucleus.
Each proton and neutron has the same mass; therefore in a carbon atom with 6 protons and 6 neutrons, the mass number is the sum of the protons and neutrons, which are 12.
For Be the atomic number is 4, so there are 4 protons in the nucleus. But, if there are 5 neutrons the mass number is 9.

60. Since the mass of each sub-atomic particle is very small, they are assigned arbitrary units.
The mass of each proton and neutron is 1 atomic mass unit (amu).
The mass of an electron is so small as to be negligible and is not considered in the mass of an atom.

61. Isotopes:
Atoms that have the same number of protons but may contain a different number of neutrons.
Carbon has a number of naturally occurring isotopes, for example:
carbon carbon -14
mass number = 13 mass number = 14
# of protons is 6 # of protons is 6
# of neutrons is 7 # of neutrons is 8

62. We represent isotopes in the following way, this is carbon – 136
Here 13 is the mass number and 6 is the atomic number

63. Atomic mass
As given on the periodic table is the average mass of all naturally occurring isotopes in a pure sample of that element.
It is not a whole number.
It must not be used to determine the number of neutrons in an atom.

64. Valence electrons
Are the outermost electrons of an atom, occupying the outer energy level.
For the representative elements the number of valence electrons is the same as the last digit of the group number.
The number of energy levels that contain electrons is the same as the period (row) number for that element in the periodic table.

65. Elements and Isotopes

67. Chemical Names and Formula
The names of compounds and their formula follows the rules laid down by the International Union of Pure and Applied Chemistry (IUPAC)

68. Atoms and Ions
Ions are atoms or groups of atoms that gain or lose electrons to become:
Cations are positively charged ions formed by atoms losing electrons.
Formed from metallic elements, E.g:
Na+ (lost 1e-)
Mg2+ (lost 2e-)
Names of cations remain the same as that of the element with ion added,
e.g. sodium ion, magnesium ion

69. Anions are negatively charged ions formed by atoms gaining electrons.
Formed from non-metallic elements,e.g.
Cl- (gained 1e-)
O2- (gained 2e-)
Names of anions have their ending changed to –ide. E.g.
chloride ion, oxide ion

70. When ions are formed the number of protons remains the same while the number of electrons changes so that the electron arrangement is the same as the nearest noble gas.
This is more stable than the arrangement of the electrons in the atom.
Sodium atoms have 11 protons and so 11 electrons.
When an electron is lost, its ions have only 10 electrons with a net charge of 1+, i.e. Na+.
These electrons are arranged like those in neon, the nearest noble gas.

71. Atoms and Ions

72. Activity - Complete the table:
Symbol of element
Change in electrons
Formula of ion
Name of ion Ca
2 lost
 Ca2+
 calcium ion F
  1 gained F-
 fluoride ion
 Al
  3 lost
Al3+
 aluminium ion Se
2 gained
 Se2-
selenide ion 

73. Compounds
Compounds are pure substances consisting of more than one type of atom.
Formed when atoms of two or more elements combine chemically.
They do so in fixed proportions by mass, since they combine in fixed ratios.

74. Molecular Compounds
Tend to have low boiling and melting points, therefore are liquids or gases at room temperature.
Compounds of two or more non-metallic elements are composed of molecules.
Molecules are neutral groups of atoms that act together as a single unit due to the atoms being bound together by chemical bonds.

75. Ionic Compounds
Are composed of positive ions (cations) and negative ions (anions) arranged in a 3D pattern (lattice) to form crystals
In a crystal, each cation is surrounded by anions and visa versa
Since the negative and positive charges (ions) are balanced, ionic compounds are electrically neutral.
Tend to have high melting points and are therefore crystalline solids at room temperature
Formed from the combination of metal and non-metallic elements

76. Chemical Formula
Show the kind of atoms and the number of them present in a Representative Unit of the compound.
A complete chemical formula should also show the state of the material at SATP. (Standard Atmosphere Temperature and Pressure)

77. Molecular Formula Chemical formula of molecular compounds.
The number of each kind of atom is shown by a subscript after the symbol of the element, e.g. H2O(l)
There is no indication of the shape or structure of the molecule.

78. Ionic Formula Chemical formula of an ionic compound.
The formula shows the ions contained in the compound, placing the symbol of the metal first.
Ionic compounds do not exist as distinct units, unlike molecules they are composed of fixed ratios of ions to provide a neutral compound.

79. Formula of ionic compounds do not represent molecules but show the lowest whole number ratio of ions in the compound,
e.g. sodium chloride has a ratio of 1:1 sodium to chloride ions so the formula is NaCl(s)
This is known as a formula unit of an ionic compound.

80. Example:
Magnesium chloride is composed of Mg2+ ions and Cl- ions.
To make this compound electrically neutral, two chloride ions must be present to balance each 2+ magnesium ion.
The ratio of Mg2+ to Cl- ions is 1:2, therefore a formula unit of magnesium chloride is MgCl2(s)

81. Ionic Charges on the Elements
To be able to state what compounds are formed we need to know what types of ions are formed from different elements, i.e. their ionic charges.
For the representative elements the charge can be easily determined from the periodic table.

82. Metals:
In groups 1, 2 and 3 lose electrons to form positive ions (cations).
The charge on the ion is equal to the group number and equal to the number of valance electrons.

83. For example: magnesium:Mg 2- 2+

84. Non - Metals:
Here the opposite is true, these elements form negative ions (anions) by gaining electrons.
The charge can be determined by subtracting the last digit of the group number from 8.

85. For example: oxygen: O 2-

86. Elements in Group 14:  
These elements do not easily form ions, instead, they share electrons in covalent bonds
Elements in Group 18:
The noble gases have a full outer energy level and therefore are very stable and do not form ions.

87. Polyatomic Ions (complex ions)
Polyatomic ions are groups of atoms that together act as a single unit with an overall charge,
e.g. SO42- is a group of 1 sulfur and 4 oxygen atoms with an overall charge of 2-
The names and formula of polyatomic ions are provided on your periodic table. You should learn be able to be able to recognize these complex ions.

88. Binary Ionic Compounds
Are composed of two mono-atomic ions, one cation (metal) and one anion (non-metal).
The name of the cation is kept the same as the name of element and always comes first, while the ending of the name of the anion is changed to -ide (note neither are capitalized).
E.gs. sodium chloride; calcium oxide

89. The formula of binary ionic compounds must show that charges balance and must be in the lowest possible whole number ratio:
Sodium chloride is formed from the combination of Na+ and Cl- ions. Thus it becomes NaCl(s)
The lowest common multiple of the ion charges can be used to determine the correct ratio of ions in the formula:

90. Aluminium oxide is from Al3+ and O2-
The lowest common multiple is 6.
3 goes into 6 twice and so there must be two aluminium ions in the formula.
2 goes into 6 three times meaning that there will be 3 oxide ions.
Therefore, the formula is Al2O3(s)

91. The criss/cross method is the simplest way to write formulae of ionic compounds.
Here the numerical charge from each ion is crossed over and used as the subscript for the other ion:
E.g. iron (III) oxide is from Fe3+ and O2-
Criss/crossing the two charges gives Fe2O3(s)

92. Molecular Compounds
These compounds are combinations of non-metallic elements. They are composed of molecules and therefore ionic charges are NOT a factor.
When two non-metallic elements combine there are often different possible combinations producing different compounds.
These different compounds must be distinguished between since they have different chemical and physical properties.

93. Prefixes are used to show how many atoms of each element are present in the molecule.
In the formula these prefixes are represented by a subscript number after the symbol for the element.

94. Generally the name goes:
The name of the first element remains the same (after any prefix that is necessary) followed by the second element’s name, the ending of which is changed to -ide.
Generally the name goes:
prefix-element name, prefix-element root -ide.
Often mono- is dropped as a prefix to the first element.

95. tetraphosphorous decaoxide
Writing formula of molecular compounds is simply done by using the prefix to determine the subscript number of that element. For example:
Reactants
Product formula
Product name
C(s) and S8(s)
CS2(l)
carbon disulphide
N2(g) and I2(s)
NI3(s)
nitrogen triiodide
N2(g) and O2(g)
N2O(g)
dinitrogen monoxide
P4(s) and O2(g)
P4O10(s)
tetraphosphorous decaoxide

96. Molecular Elements:
Most elements are found as single atoms, i.e. they are mono-atomic.
All metals are like this and some non-metals, however, some elements are found as diatomic molecules and even polyatomic molecules.
This arrangement allows them to be more stable.

97. Hydrogen, nitrogen, oxygen and the halogens are all diatomic
Hydrogen, nitrogen, oxygen and the halogens are all diatomic. Their formulas are written H2(g), O2(g), Cl2(g) etc.
Sulfur (S8(s)) and phosphorus (P4(s)) are polyatomic.

98. Acids and Bases Acids have special names.
Acids are defined as compounds that release hydrogen ions when dissolved in water.
Acids and Bases will be reviewed and discussed in Science 30.

99. Chemical Reactions
Chemical reactions always involve reactants changing to products.
This involves the rearrangement of the atoms of the reactants into new substances as the products.
Particles of the reactants must collide for reactions to occur.
Reactants must have a certain minimum energy for reactions to occur.

100. Chemical reactions do not cause atoms to be created or destroyed.
Changes in energy are always associated with chemical reactions due to the breaking and/or formation of chemical bonds.
Other evidence for chemical reactions includes production of a solid, gas or odor or a change in colour.

101. Features of Chemical Reactions

102. Chemical Equations:
Chemical equations are used to represent what happens in a chemical reaction.
They always indicate:
Reactants  Products
They always show that matter is conserved i.e. that no matter is created or destroyed.

103. Word Equations:
Show the names of the chemicals involved in the reaction and the products:
E.g.
iron (III) + oxygen  iron (III) oxide

104. Balance Chemical Equations:
Show the quantities of the reactants and products involved in terms of the number of atoms of each substance.
Use a symbolic representation of the substances as their chemical formulas.
Have the same number of each type of atom in the reactants and the products, i.e. they are balanced in accordance with the law of conservation of matter.
Show the states of matter of the reactants and products.
E.g. 4Fe(s) O2(g)  2Fe2O3(s)

105. Rules for Balancing Chemical Equations
Determine the correct chemical formula for all reactants and products.
Write the formulas for the reactants on the left and those for the products on the right and separate them with an arrow. Place a plus sign between two or more reactants or products.
Indicate the state of each of the reactants and products.
Count the number of atoms of each element on both the reactant and product sides of the equation. Polyatomic ions that remain unchanged on both sides of the equation are counted as a single unit.

106. Rules for Balancing Chemical Equations
Balance the number of atoms of each element on both sides of the equation using coefficients.
Place the coefficient in front of the chemical formula; This multiplies all of the atoms in the formula by the value of the coefficient (when there is no coefficient it is taken to be 1)
Coefficient must be whole numbers
Start with elements other than hydrogen and oxygen
When balancing hydrocarbon combustion, start with carbon always, followed by hydrogen and finally oxygen
Check that the number of each atom or polyatomic ion is the same on both sides of the equation
Make sure that the coefficients are the lowest possible ratio
Practice, practice, practice!!!

107. Interpreting balanced chemical equations
So far we have considered that we are working with individual atoms, ions or molecules of the reactants and products.
However, these entities are much too small to see and so observable changes in chemical reactions must involve very large numbers of particles.

108. A convenient way to communicate these enormous numbers has been developed by scientist.
The term mole is used to represent these large numbers in a convenient way. Just like one dozen always represents 12 eggs or donuts, one mole represents a certain number of particles.
This number is called Avogadro’s Number and it is a rather unusual number, 6.02 x 1023.

109. One mole of any substance is always 6
One mole of any substance is always x 1023 particles of that substance.
Essentially, a mole represents a certain number of something just like a dozen does.
Now we can interpret chemical equations a little differently.
For the following reaction we can talk about individual particles, dozens of particles or Avogadro’s number of particles:

110. We will deal with the mole concept in much more detail later.
4Fe(s) O2(g)  2Fe2O3(s)
4 atoms 3 molecules 2 formula units
4 dozen atoms 3 dozen molecules 2 dozen form u’ts
24.08 x 1023 atoms x 1023 molecules x 1023form u’ts
4 mol of atoms 3 mol of molecules 2 mol of form units
The mole concept becomes even more useful when we find that we can relate moles to the mass or volume of reacting substances.
We will deal with the mole concept in much more detail later.

111. Types of Chemical Reaction
To be able to predict the outcome of a chemical reaction and write a complete equation, you must be able to recognize the type of reaction that is occurring. There are five basic reaction types:
Formation
Decomposition
Single replacement
Double replacement (neutralization reactions are a special kind of double replacement)
Combustion

112. Types of Chemical Reaction

113. Formation reaction
This is a reaction of two or more elements forming an ionic or molecular compound.
E.g. magnesium reacts with oxygen to form magnesium oxide.
2 Mg(s) O2(g)  2 MgO(s)
Summary:element A + element B  compound AB

114. Decomposition reaction
This is a breakdown of a compound into simpler compounds or into the component elements.
E.g.1 calcium carbonate decomposes into calcium oxide and carbon dioxide.
CaCO3(s)  CaO(s) CO2(g)
E.g.2 water decomposes into its constituent elements by using electricity, electrolysis
  2 H2O(l)  2 H2(g) O2(g)
Summary: compound AB  element A + element B

115. Single replacement reaction
This is a reaction of an element with a compound to produce a new element and a new compound.
Basically the two elements switch around.
Occurs in aqueous solutions, usually involving replacement of metal ions.
E.g. copper solid reacts with a silver nitrate solution.
Cu(s) AgNO3(aq)  2Ag(s) + Cu(NO3)2(aq)
Summary: element A + compound BC  element B + compoundAC

116. This type of reaction can also happen with halogens.
E.g.
Cl2(g) + 2NaBr(aq)  Br2(l) NaCl(aq)

117. Double replacement Reaction
This is a reaction between two ionic compounds in solution to produce two new compounds.
Occurs in aqueous solution.
Generally, the new compounds are produced by switching the partners of the negative ions that are found in the reactants.
E.g. calcium chloride solution is mixed with a sodium carbonate solution .
CaCl2(aq) + Na2CO3(aq)  CaCO3(s) NaCl(aq)
compound AB + compound CD  compound AD + compound CB

118. A specific double replacement reaction called a neutralization reaction occurs when an acid reacts with a base to produce water and an ionic salt.
E.g. Hydrochloric acid reacts with potassium hydroxide.
HCl(aq) + KOH(aq)  KCl(aq) + H2O(l)
acid base  salt + water
Double replacement reactions always produce either a solid or a molecular compound.

119. Combustion reactions This reaction type involves oxygen as a reactant.
It is an exothermic reaction.
The products being produced in a combustion reaction must be determined by the composition of the substance being burned.

120. CHEMICAL REACTIONS
If the substance that is undergoing combustion contains:
carbon, then CO2(g) is produced.
hydrogen, then H2O(g) is produced.
carbon and hydrogen, then CO2(g) & H2O(g) is produced. This is a hydrocarbon Combustion.
sulfur, then SO2(g) is produced.
a metal, then the oxide of the metal with the most common ion charge is produced.

121. Endothermic and Exothermic Reactions

122. End of Review