The document discusses matter and chemical change. It defines matter as anything that has mass and takes up space. Matter exists in different states - solid, liquid, and gas - which are determined by the motion and spacing of particles. Changes between states of matter involve adding or removing heat. Chemical and physical properties are also discussed, where chemical changes alter the chemical composition and physical changes do not. The document provides examples of physical and chemical properties and changes. Classifying matter as elements, compounds, or mixtures is also covered.

1. Unit B
Matter and Chemical Change

2. What is matter?
• Matter: The stuff that makes up the universe that is
not energy
• Matter:
1. Has mass
2. Takes up space

3. The Particle Model of Matter
1. All matter is made of tiny particles
2. Particles are in constant motion
3. There are attractive & repulsive forces between particles
4. These are spaces between the particles
5. Difficult matter is made of different particles

5. States of Matter
States of Matter: distinct forms in which matter can
exist.

6. States of Matter
Bose-Einstein Condensate
5

7. Solid Liquid Gas
Motion of
Particles
Space between
particles
Behaviour in a
container
LOW LOW HIGH
Not affected
by the
container
Takes shape
of the
container
Fills the
container

9. Changes Between States of Matter
Melting: Going from a solid to a liquid when heat is added
Boiling: Going from a liquid to a gas when heat is added
Condensation: Going from a gas to a liquid when heat is
removed
Freezing: Going from a liquid to a solid when heat is removed
Sublimation: Going from a solid to a gas when heat is added
Deposition: Going from a gas to a solid when heat is removed

11. Physical & Chemical
Properties

12. Physical Properties
Physical Properties: a characteristic of matter that can be
observed and measured without changing the chemical
identity of the sample.
Note: It’s important to remember that when a substance
undergoes a physical change, such as melting, its
appearance or state may be altered, but its composition
stays the same.

13. Physical Properties:
• Melting Point
• Boiling Point
• Hardness
• Malleability
• Substance that can be rolled into a sheet
• Ductility
• Substance that can be stretched into a wire
• Crystal Shape
• Solubility
• Substance that can be dissolved into another
• Density
• Conductivity
• Ability to conduct electricity or heat

15. Chemical Properties
Chemical Properties: a property or characteristic of a
substance that is observed during a reaction in which
the chemical composition or identity of the substance
is changed
Note: A chemical change always results in the
formation of a different substance or substances.

16. Chemical Properties:
•Reaction with acids
•Ability to burn
•Reaction with water
•Behavior in the air
•Reaction to heat

17. Classifying Matter

18. Classifying Matter
Pure Substance:
• is made of only one kind of matter
• has a unique set of properties that sets it apart from any other
kind of matter.
• (It is further classifies into an element or a compound)
Mixture:
• is a combination of pure substances not combined chemically

19. Pure Substances
Element (Example Na or Sodium/O or Oxygen):
• Matter than cannot be broken down into any simpler substance.
• The most purest matter because it contains only one kind of atom
• Represented by a symbol on the periodic table
Compounds (H20 or CO2) :
• A substance formed by 2 or more elements chemically fixed in
proportion (molecule).

20. Pure Substances
Compounds (H20 or CO2) :
• A substance formed by 2 or more elements chemically fixed in
proportion (molecule: atoms bonded together).

21. Mixtures
Mechanical Mixture (Heterogeneous):
• Some or all particles can be seen and be separated (Ex. Salad)
Solution (Homogeneous):
• Not separate visibly, as one is dissolved in another (Ex. coffee)
• A substance dissolved in water is called an aqueous solution

22. Mixtures
Suspension:
• Mixture where tiny particles of one substance are held within
another. Substances can be separated by centrifuging or
filtration. (Ex. Salad dressing)
Colloid:
• Cloudy mixture, where the
particles are suspended
and difficult to separate
(Ex. milk)

23. Tyndall Effect
Solution vs Colloid
Solution vs Colloid

25. 1. How is an element different than a compound? Give an example of each.
2. What is the difference between a pure substance and a mixture?
3. How is a suspension different from a colloid?
4. Classify the following substances as an element, compound, or mixture:
a) Pop is composed of water, sugar, and carbon dioxide.
b) Graphite in a pencil is composed of carbon.
c) Carbon dioxide is composed of carbon and oxygen.

26. Physical & Chemical
Changes

27. Physical Changes
Physical Change: When a substance undergoes a
physical change, its state may be altered, but its
chemical composition is the same

29. Physical Changes
To identify a physical change:
• You can separate the end products (reactants) to
original substances. Ex. Salt water, sand & rocks
• You are able to re-freeze or melt the product again.
Ex. Ice cream

30. Chemical Changes
Chemical Change: When 2 or more substances react
and form 1 or more new substances (The new
substances formed are completely different from the
original reactants)

32. Chemical Changes
Four actions that identify a chemical change:
1. Change in color
2. Change of odour (if present)…safety!!
3. Formation of a solid (precipitate) or a gas
4. Release or absorption of energy in the form of light & heat

34. Timeline

35. 8000 BC
• Used stones & bone for tools
• Started to control Fire (cooking, baking)

36. 6000-1000 BC
• Found metals
• Gold became very valuable
• Chemists were only interested in materials of value

37. 4500 BC
• Copper was discovered and led to bronze
• Humans could now make hard strong tools (swords, spears &
other tools)

38. 1200 BC
• Hittites discovered how to extract iron from rocks
• This started the iron age (iron + copper = steel)
• Started to extract oils & juices from plants

39. 400 BC
• Greek philosopher (Democritus) used the word atomos to
describe the smallest particles that could not be broken down
any further

40. 350 BC
• Aristotle another more popular geek philosopher had the idea
that all matter came from Earth, Water, Fire and Air.
• Because Aristotle was more respected people accepted his
theory over Democritus' idea.

41. The Next 2000 Years
• Most work with matter was done by Alchemists, people who
were part chemist and part magician.
• Most of their job was working for the King to turn metals into
gold.

42. Late 1500
• People really started to investigate the world around them
• Started to form the scientific method.

43. 1660
• Robert Boyle experimented with the
behaviors of gases.
• He believed everything was made of tiny
particles of different size and shape.
• Boyle thought the purpose of chemistry
was to determine the types of particles in
each substance.

44. 1780
• Antoine Laurent Lavoisier "the father of
modern chemistry"
• Studied chemical interactions and
developed a system of naming them.
(everyone could use the same names)

45. 1808
• John Dalton suggested matter was made of
elements.
• Pure substances contained no other substances
(all one type of atom)
• Developed the billiard ball model (atoms are
solid spheres)
C
Cl
H
Fe
N
O
Na

46. 1897
• J. J. Thomson discovered electrons
• Developed a raisin bun model of matter.
• Negative electrons balance out positive sphere so atom has no
electrical charge.

47. 1904
• Hantaro Nagaoka modeled matter after the solar system
• Large positive center with negative electrons orbiting around

48. 1908
• Ernest Rutherford (working at McGill in
Montreal) found evidence of central
nucleus.
• In his experiment he shot positively
charge particles through thin foil and
found mostly empty space and
the nucleus was the core that would stop
the particles.

49. 1922
• Niels Bohr working with Rutherford found electrons orbit is
specific shells (circular orbits)
• Electrons jump between shells by gaining or losing energy

50. More 1922
• James Chadwick found protons & neutrons
(particles with no charge) inside the
nucleus of an atom.

51. Today
• People still use Bohr model of the atom
• Quantum mechanics found electrons exist
in a charged cloud around the nucleus.

52. History of Periodic
Table of Elements

53. Atom
• Proton: Part of an atom’s nucleus,
positive electric charge. Roughly
equal mass to neutrons.
• Neutron: Part of an atom’s nucleus,
no electric charge (Neutral).
Roughly equal mass to protons.
• Electron: Part of an atom, negative
electric charge

54. Size of an Atom
If an atomic nucleus was the size of a
pea, and that pea was inside a football
stadium... then a mosquito flying
around in the football stadium would
represent an electron and the football
stadium would represent the size of
the atom.

55. Early Organizing of the Elements
• Early chemists used symbols of the sun and planets to
represent seven metallic elements
• Early 1800’s more than 30 elements had been identified and as
chemistry developed even more were identified
• Chemists tried to group elements by properties but everyone
was using a different system and it was very confusing. They
needed a new system

56. Early Organizing of the Elements

57. Early Organizing of the Elements
• John Dalton developed a new set
of symbols in early 1800s

58. Early Organizing of the Elements
• Swedish chemist Jon Jacob Berzeliusin, 1814, suggested
using letters rather than pictures to represent the elements.
The first letter capitalized of an element would become the
symbol. A second smaller letter would be added for elements
that had the same first letter.
• H - Hydrogen, He - Helium

59. Early Organizing of the Elements
How did they put these elements into an order?
• Elements were organized in order of increasing
atomic mass. Atomic mass is the mass of one
atom of an element and is measured in an
atomic mass unit (amu). Scientists were able to
determine the average atomic mass unit of an
atom by comparing it to a carbon atom (12.0).

60. Early Organizing of the Elements
How did they put these elements into an order?
• 1864, English Chemist John Newlands recognized a pattern when
elements were listed by increasing atomic mass – they seemed to
repeat properties at regular intervals. He called this pattern “the law
of Octaves” as the pattern was similar to the octave scale on a piano.
Many scientists thought the idea was silly and refused to accept it.

61. Early Organizing of the Elements
Who found the pattern for the
elements?
• Dmitri Mendeleev collected
63 elements known to exist
in his time and used the
properties to organize them
into a pattern. He sorted
them using cards.

63. Based on Mendeleev’s Periodic Law

64. Spiral Variation (Theodor Benfrey’s 1960)

65. Discrete Levels (Timothy Stowe)

66. Triangular

67. Reading the Periodic
Table of Elements

68. Label your blank table as we go!

69. Groups/Families & Periods
Groups/Families: Vertical Columns – organized because of
their similar chemical behaviour, same number of electrons in
their outermost shell.
Periods: Horizontal Rows – same number of electron shells

70. Groups/Families & Periods
Groups/Families
Periods
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
1
2
3
4
5
6
7

71. Metals, Metalloids, Non-Metals
Metals: Shiny, Malleable, Ductile and conduct electricity
Non-Metals: Are dull, brittle and only carbon conducts
electricity
Metalloids: have properties of both metals and non-metals

72. Color the One Blank Periodic Table

73. Metals, Metalloids, Non-Metals

74. Metals, Metalloids, Non-Metals

75. Periodic Table Categories (In more Detail)

76. Group 1 & 2 – Alkali/Alkaline-Earth Metals
Alkali Metals (Group 1): Most reactive metal (only need to lose
1 electron in their outer shell)
Alkaline-Earth Metals (Group 2): Also very reactive but not as
reactive as group 1 (Need to lose 2 electrons)

77. Color the 2nd Blank Periodic Table and
Label Categories

78. Group 1 – Alkali Metals

79. Group 2 – Alkaline-Earth Metals

80. Group 17 & 18 – Halogens/Noble Gases
Halogens (Group 17): Most reactive, typically react with
metals (only need to gain 1 electron in their outer shell)
Noble Gases (Group 18): Odourless and colorless gasses,
essentially non reactive. (Complete outer shells)

81. Group 17 – Halogens

82. Group 18 – Noble Gases

83. Transition Metals

84. Other Metals

85. Rare Earths

86. Radioactive Rare Earths

87. Periodic Table Categories (In more Detail)

88. What is in the Squares?
Element Name
Element Symbol
Atomic Number
Atomic Mass

89. What is in the Squares?
Atomic Number: Number of protons in the atom
Atomic Mass: Mass of
Protons + Neutrons + Electrons

90. Ion Charge
Ion Charge: Most common charge
of the atom in it’s ionic state

92. Periodic Table Scavenger Hunt!

93. Electron Shell Levels

94. Bohr Model Diagram

95. Bohr Model Diagram
What is this element?
Potassium

96. Compounds

97. Compounds
Compound: A chemical combination of two or more elements
in a specific ratio
Examples (Write One):
• Table Salt : NaCl (sodium chloride)
• Baking Soda : NaHCO3 (sodium bicarbonate)
• Water : H2O (dihydrogen oxide)

98. Compounds
Chemical Formula: Identifies the elements in the compound
and the number of each element.
Examples (Write One):
• Table Salt : NaCl (sodium chloride)
• Baking Soda : NaHCO3 (sodium bicarbonate)
• Water : H2O (dihydrogen oxide)

99. Compounds
Nomenclature: Naming system created by French chemist
Guyton de Morveau in 1787.
If you know the formula you can determine the name. If you
know the name you can determine the formula.

100. NaCl Salt
(sodium ​chloride)
1 atom Na
1 atom Cl
H2O Water
(dihydrogen ​oxide
)
2 atoms H
1 atom O
C6H12O6 Glucose 6 atoms C
12 atoms H
6 atoms O
H2SO4 Battery Acid
​(sulfuric acid)
2 atoms H
1 atom S
4 atoms O
CH3COOH Vinegar
(acetic ​acid)
2 atoms C
4 atoms H
2 atoms O

101. Compounds
The Subscript (the small number behind some elements) tells you how
many of the element are in that compound.
After a chemical formula we can also have subscript letters in parentheses.
This lets us know the state of the compound.
(s) is a solid
(l) is a liquid
(g) is a gas
(aq) is an aqueous solution (a substance dissolved in water)

103. Compounds
There are two different types of compounds depending on the types of
elements that make the compound.
They Are:
Ionic
Compounds
Molecular
Compounds

104. Ionic Compounds
Ionic Compounds: Contain at least one metal and one or more non-metal
Examples:
NaCl : Sodium chloride CuSO4 : Copper (II) Sulphate

105. Ionic Compounds
When a metal and non-metal combine they form IONIC bonds
In order to form a compound the metal loses electrons and becomes
positively charged.
The non-metal gains electrons and becomes negatively charged.
These charged elements are known as Ions.

106. Ionic Compounds

107. Ionic Compounds
Magnesium chloride is an ionic compound.
Its chemical formula is MgCl2
That is 1 Mg and 2 Cl
Mg
-
-
-
-
- -
-
-
--
- -
The two outer electrons ​of the
metal are given to the chlorine
atoms, ​filling their
outer ​electron shell

108. Ionic Compounds
Properties of Ionic Compounds:
• High melting point
• Good conductivity
• Distinct crystal shape
• Solid at room temperature
Polyatomic ions - some atoms bond together and have an overall positive or
negative charge. these groups of atoms act as one.

109. Molecular Compounds
Molecular Compounds: Molecular compounds contain only non-metals
Examples:
N2O4: Dinitrogen tetraoxide P4O10: tetraphosphorus decaoxide

110. Molecular Compounds
Molecular compounds form
COVALENT bonds
In order to form a compound
the non-metals must share
electrons in order to fill the
outer shell of each element.

111. Molecular Compounds
N
-
-
-
-
-
- -
Nitrogen trichloride is a
molecular ​compound. Its chemical
formula is NCl3
Each chlorine ​shares one electron
with the nitrogen. ​The
nitrogen ​shares 3 ​electrons, one
for ​each chlorine.

112. Molecular Compounds
Molecular Compounds:
• Solid, liquid or gas at room temperature
• Lower melting and boiling points because the covalent bond is weaker
than ionic bond

113. Practice
Identify the following chemicals as either ionic ​compounds or molecular
compounds
Sodium
Chloride
Ionic Magnesium
​chloride
Ionic
H20 Molecular Tin(IV) oxide Ionic
SO2 Molecular Dinitrogen
pentaphosphide
Molecular
Nickel (II)
bromide
Ionic SiO2 Molecular
Carbon
monoxide
Molecular Mn2O7 Ionic

114. Naming Ionic
Compounds

115. Naming Ionic Compounds
Step 1: Determine if the substance is a molecular or an ionic compound
Example: NaCl (Metal & Non-metal)
Step 2: Write the full name of the metal first
Na = Sodium
Step 3: Add the name of the non-metal second
Cl = chlorine
Step 4: Drop the ending of the non-metal, the “ine”, and change it to “ide”
chlorine  chloride
NaCl = Sodium chloride

116. Naming Ionic Compounds
** Step 5. If the metal is not from group 1 or 2 (i.e.3-17) we may need to
determine and state the ionic charge of the metal.
For example SnO2.
There are 2 Oxygen atoms and 1 tin atom.
Each oxygen atom has a 2- charge (see common ion charges). That is 4-
charges in total.
There is only 1 tin atom and its charge must equal and
balance the negatives. The tin must have a 4+ charge, one of its common
ion charges (top right of the box).

117. Naming Ionic Compounds
For example SnO2.
Sn (charge +4) ----- O (charge -2)x2
Convert the ionic charge of the metal into a Roman
numerals number. (4 is IV)
Insert the Roman numeral between the metal and non
metal using brackets
Ex. Tin (IV) oxide

118. Simple Examples
Together:
Rb2O
Rubidium oxide

119. Simple Examples
Try Yourself:
CaO
Calcium Oxide
Sr3N2
Strontium Nitride
BaBr2
Barium Bromide

120. Complex Examples
Together:
Cu2O
Copper (I) oxide

121. Complex Examples
Try Yourself:
CrN
Chromium (III) nitride
Nb3P5
Niobium phosphide
Mn2O7
Manganese (VII) oxide

122. Naming Molecular
Compounds

123. Naming Molecular Compounds
Step 1: Determine if the substance is a molecular or an ionic compound
Example: N2P5 (Non-metals)
Molecular compounds use Greek prefixes to tell us the number of each
atom. Copy the prefixes in your notes:

124. Naming Molecular Compounds
Example: N2P5 (Non-metals)
Step 2: Determine how many atoms of the first listed substance ​there are. (there
are 2 N’s)
Step 3: Attach the prefix for that number to the beginning of ​the elements name.
**If there is only one do not add a ​prefix to the first element.
Ex. 2=Di Dinitrogen
Step 4: Repeat step 2 and 3 for the second element.
** If there is only one atom use the prefix mono.
Ex. 5 = penta Dinitrogen pentaphosphorous
Step 5: Change the ending of the second element to and -ide ​ending.
Ex. Dinitrogen pentaphosphide

125. Simple Examples
Together:
NO2
Nitrogen Dioxide

126. Simple Examples
Try Yourself:
CO
Carbon Monoxide
S2Cl4
Disulfur Tetrachloride
B3I5
Triboron Pentaiodide

127. Creating Chemical
Formulas from Names

128. Creating Chemical Formulas
Example: Dicarbon tetrafluoride / Sodium Chloride
Step 1: Find the symbols for each
C F / Na Cl
Step 2: If the substance is molecular use the prefixes to tell you the
number of each atom.
Ex: C2F4
If the substance is ionic we must find the ion charges for each
atom.
Ex. Na +1 Cl -1
Step 3: Balance the charges so that there are the same number of
positives and negatives.
Ex. Only one of each is needed to balance Na and Cl
Sodium Chloride = NaCl

129. Examples
Try Yourself:
Calcium iodide
CaI2
diphosphorous trinitride
P2N3

130. Naming with
Polyatomic Ions

131. Naming with Polyatomic Ions
Sometimes when atoms group ​together, many atoms can work
as ​one.
Ex: SO4 -2
When this group of atoms act as a ​charged particle it is called
a ​polyatomic ion.
These polyatomic ions are only found in ​IONIC COMPOUNDS

132. Naming with Polyatomic Ions
Polyatomic ions have special rules ​for writing chemical formulas:
Step 1: Positive ions are written first ​negatives second (similar to
positive metals and ​negative non-metals)
Step 2: If more than one polyatomic ion is ​required to balance the
charges the entire ion is put in brackets.
Ex: Iron (III) Carbonate
Fe +3 CO3 -2 = Fe2(CO3)3
Note: the ending for polyatomic does not have an ​-ide ending. Each
are different