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 Unit 2
Matter and Chemical Change
 Topic 1
Key Concepts
 WHMIS (Workplace Hazardous Materials Information System)
 Substances and their properties
 Elements, Compounds and Atomic Theory
 Topic 1
My Learning Outcomes
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When I have completed this section I will be able to :
Identify and evaluate dangers of caustic materials and potentially explosive reactions.
Investigate and describe properties of matter
describe and apply different ways of classifying matter based on its composition and
properties.
 Topic 1
 Safety in the Science Classroom
 In any science activity, the safety of you, your classmates,
and your teacher are of the utmost importance.
 There are two areas of special consideration:
 Understanding warning labels
 Following safety procedures
 Keep Safety in Mind
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Before you start any science activity:
Follow all safety instructions from the teacher or from the text book.
Identify possible hazards and report them immediately
Show respect and concern for your own safety and the safety of
others
 Read Toolbox 1: Safety in the Laboratory (Page 478)
 Organizing Matter
 Almost all matter exists as either a solid, liquid or gas. (A fourth state of
matter is the plasma state - created when a large amount of energy is
added to a gas)
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 Physical Properties of Matter
Matter
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Colour
Lustre
Melting Point
Boiling Point
Hardness
Malleability
Ductility
Crystal Shape
Solubility
Density
Conductivity
Chemical Properties of
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Reaction with acids/bases
ability to burn
reaction with water
behavior in air
reaction to heat/light
 Pure Substances and Mixtures
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 Observing Changes in Matter
 Changes in matter are classified as physical or chemical.
 Physical Change - matter changes state
 Chemical Change- two or more materials react to create new materials
with different properties from the original substances.
 Evidence of Chemical Change
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A colour change
Formation of an odour
Formation of a gas (bubbles)
Formation of a solid (precipitate)
Release or absorption of energy (heat, light, sound)
Topic 2.0
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An understanding of the nature of matter has developed through
observations over time.
 Evolving Theories of Matter
 Organizing the Elements
 The Periodic Table Today
 Topic 2 – Key Concepts
 Substances and their properties
 Elements, compounds and atomic theory
 Periodic table
 Topic 2 - Learning Outcomes
 When I have completed this section I will be able to:
 Distinguish between observation and theory and provide examples of how models are
used to explain observations
 Demonstrate understanding of the origins of the periodic table and relate physical and
chemical properties of the elements to their position on the periodic table
 Use the periodic table to
 Identify numbers of protons and electrons in each atom
 Describe the relationships between the structure of atoms in each group and the
properties of the elements in that group.
 2.1 Evolving Theories of Matter
 People try to make sense of the world around them by suggesting
explanations (theories) to explain their observations. Over time these
theories are modified as new evidence is discovered….
 Ex. Human beings once thought the earth was flat and as we began to
discover more about our planet the theories changed.
 Stone Age Chemists
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8000B.C.
In the Middle East
Metals not discovered yet
Learned to start and control fire
Controlling fire led to producing glass and ceramic tile.
 Early Interest in Metals
and Liquid Matter
 6000 B.C. – 1000 B.C.
 Chemists investigated valuable materials
 Many were metals like gold & copper
 Chemists wanted to understand how copper could be controlled.
 Original discovery that heat made copper very useful was accidental –
some fell in a fire.
 Later experiments (4500B.C.) led to creation of Bronze – made from
copper and tin.
 Iron Age
 1200 B.C.
 Hittites in Middle East discovered how to extract iron from
rocks and turn it into a useful material.
 Led to combining iron & carbon to produce an even harder
material – steel.
Emerging Ideas About the Composition of
Matter
 500 B.C. - Idea that all matter is made up of particles started with Greek philosophers
 400 B.C. –
 Democritus used word atomos to describe smallest particles that could not be broken
down. (means indivisible)
 Believed each type of material was made up of different atomos
 These particles gave each material its own unique set of properties.
 350 B.C. – Aristotle said everything was made of earth, air, fire, and water.
 He was so well known and respected that his ideas were accepted over Democritus
for 2000 years.
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 Chemistry Develops into a Science
 1660’s (A.D.) Robert Boyle - matter is made up of tiny particles.
 1770 – 1780’s – Antoine Laurent Lavoisier
 develops a system for naming chemicals
 “the father of modern chemistry”
Atomic Theory
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Billiard Ball Model
 1808 – John Dalton
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Suggests matter is made up of elements
First person to define element
Each element is composed of an atom
Atoms cannot be broken into smaller particles
 Raisin Bun Model
 1897 – J.J. Thomson
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First person to discover a subatomic particle
Named them electrons
Thomson discovered the electron in 1897
So, the atom was not the smallest unit of matter
Since there are negative electrons in the atom,
there must also be positive substances
 Thomson’s model of the atom: the positively charged substance fills the
atom and the electrons are embedded throughout it
 The atom is like a “raisin bun”
 Rutherford Model
 1898 – 1907 Earnest Rutherford (working in Canada)
 Atoms mainly empty space
 Core (center) was a tiny positively charged nucleus (1/10000th the size of
the atom)
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 Bohr’s Model
 1913 Niels Bohr (while working with Rutherford)
 Electrons move in specific circular orbits – electron shells (not randomly)
 Electrons could jump shells by gaining or losing energy
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 Bohr model refined by James Chadwick
 Discovered positive protons and neutral charged neutrons in the nucleus.
 Neutron and proton same mass but 1800 times more than an electron.
 Quantum Mechanics Model
 Describes electrons as existing in a charged cloud around the
nucleus.
 2.2 Organizing the Elements
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1814 Jons Jacob Berzelius
 Suggested using letters rather than picture to represent elements..
 The first letter (capitalized) of an element would be the symbol. For elements with
same first letter (hydrogen & helium) a small second letter would be added.
 H – hydrogen & He for helium.
 1869 – Dmitri Mendeleev
 Organized elements in a way that reflected patterns in the properties of elements.
 Was able to show that the properties vary periodically with increasing atomic mass
 Had gaps in his chart of elements but predicted that new elements would be
discovered that would have the properties and atomic mass needed to fit.
 Within 16 years was proven right- gallium discovered.
The Periodic Table Today
Understanding the Periodic Table
Horizontal rows (periods)
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Numbered 1-7
From left – right elements change from metals to non-metals
Most reactive metals on the left
Atomic # increases as we move to the right
Groups (Families)
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Vertical columns
Elements in the same family have similar properties
Numbered 1 – 18
Special Names
Alkali metals (group 1) – most reactive metals
Alkaline Earth metals (group 2) very reactive
Halogens (Group 17) most reactive non-metals
Noble gases (group 18) – very stable, almost never react
Element Information
Element Symbols
 Some symbols are based on latin names of elements instead of the modern name
 Ex. Gold – Au (aurum)
 Lead – Pb(plumbum)
Atomic Number
 Tells us # of protons
 Atoms of same element all have the same # of protons
 Ex. Oxygen – 8 protons
 Also tells us # of electrons
Atomic Mass
 Tells us total mass of all protons and neutrons in an atom. (electrons are too tiny
to effect mass of an atom)
 Not all atoms of same element have same atomic mass (different #’s of neutrons)
Mass Number
 Is the sum of protons and neutrons in an atom
 Mass # - atomic # = neutrons
Patterns of Information
Metals
 Left ¾ of periodic table
 Shiny, malleable, ductile, conduct electricity
Non-metals
 Right side of periodic table
 Solid or gas
 Dull, brittle elements
 Do not conduct electricity (carbon is exception)
Metalloids
 Diagonal row of elements between metals & non-metals
 Have both metallic and non-metallic properties
Topic 3
Compounds form according to a set of rules.
 Naming Compounds
 Ionic Compounds
 Molecular Compounds
Topic 3 Key Concepts
 Periodic table
 Elements, compounds and atomic theory
 Chemical nomenclature
Topic 3 - My Learning Outcomes
When I have completed this section I will be able to:
 Distinguish between ionic and molecular compounds and describe examples of each.
 Read and interpret chemical formulas for compounds of two elements, and give the
IUPAC name and common name for these compounds
 Identify examples of combining rations/number of atoms per molecule found in
common materials and use ionic charges to predict combining rations in ionic
compounds.
3.1 Naming Compounds
Combining Elements to Make Compounds
Looking around our homes we are surrounded by chemicals - from water to
bleach.
Each compound has a chemical name and a chemical formula.
 This formula identifies which elements and how much of each are in the
compound.
 Sodium Chloride (aka. table salt) chemical formula NaCl.
 Sodium Bicarbonate (aka baking soda) NaHCO3
Naming Chemical Compounds
 Until 18th century there was no standard system. Names for chemicals differed
from country to country.
 Ex. Hydrochloric acid and muriatic acid refer to the same thing.
 1787 a French Chemist (Guyton de Morveau) created a naming system
(nomenclature) for compounds
 The chemical name for each element in the compound is used in the compound name.
Ex. Zinc and oxygen form zinc oxide.
 Since 1920, the International Union of Pure and Applied Chemistry (IUPAC)
decides on the name for every chemical compound discovered.
Interpreting Chemical Names and Formulas & Indicating Physical
State.
 If you know the formula you can determine the chemical name; if you know the
chemical name then you can determine the formula.
 Ex. H2O – the two indicates there are two atoms of hydrogen to go with every atom of oxygen in
water.
 NaCl – one atom of sodium to every one atom of chlorine.
 Another common notation added to chemical compounds indicates the state at room
temp.
 After the chemical formula a subscript (s) for solid, (l) for liquid, (g) for gas or an (aq)
to show something dissolved in water. (aqueous solution)
3.2 Ionic Compounds
Ionic Compounds are pure substances formed from the attraction
of (+) & (-) charged particles called ions. (ex. Table salt formed
from (+) sodium & (-) chloride ions.
All ionic compounds
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Are solid at room temp.
Conduct electricity when dissolved in water
Formed from metals (+) and non-metals (-)
Have high melting points
Ion Charges
An ion is an atom or group of atoms that has become electrically charged
through the loss or gain of electrons.
To indicate ions a (+) or (-) sign is placed to the upper right of the element
symbol (Ex hydrogen ions are h+ and chlorine ions are Cl - )
Some ions also form when certain atoms of elements combine. (ex carbonate
CO3 2-) These polyatomic ions act as one.
Naming Ionic Compounds
Two rules:
 The chemical name of the metal or (+) ion goes first followed by the
name of the non-metal or (-) ion.
 The name of the non-metal (-) ion changes its ending to ide. (NaCl is
not sodium chlorine but sodium chloride.
o ** there is one exception to these rules – where negative ions are
polyatomic ions the name remains unchanged.
Some elements have more than one ion charge. To show which ion is being
used a Roman Numeral is added.
Ex. iron(III) oxide contains the Fe3+ ion. (Fe2O3)
Ex. iron(II) oxide contains the Fe2+ ion (FeO)
Using Ion Charges to Write Chemical Names
 Step 1 – print the metal elements symbol with its ion charge. Next to it
print the non-metal elements symbol and ion charge.
Ca2+ Cl 1 Step 2 – balance the ion charges. You must have an equal number of (+)
charges & (-) charges.
Ca2+ Cl 1- Cl 1 Step 3 – Write the formula showing the number of atoms of each
element to balance the formula. ** Do Not Include The Ion Charge.
CaCl2(s)
Molecular Compounds
Writing Formulas for Molecular Compounds
Formed when non-metals combine.
 May be solids, liquids or gases
 Tend to be insulators, or poor conductors of electricity
 Have relatively low melting & boiling points
Writing Formulas
 no ions are present and the ion charge is not used in the formula
 Hard to predict combinations
 Formulas show they type and # of atoms of each element.
 Ex. Ammonia NH3(g)
Naming of Molecular Compounds
Many molecular compounds known by their common name. Ex. Water & ammonia
Names of these do not tell what elements they are made of
All molecular compounds except those with hydrogen can be named using
the following rules:
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The first element uses the element name
The second element has the suffix (ends in) ide.
More than one atom in the formula a prefix is used which specifies # of atoms.
If the first element is only one atom – no prefix.
Examples of Naming Molecular Compounds
Topic 4 – Substances undergo a chemical change when they
react to produce different substances
4.1 Chemical Reactions
4.2Conservation of Mass in Chemical Reactions
4.3Factors Affecting the Rate of a Chemical Reaction
Topic 4 - Key Concepts
Endothermic and exothermic reactions
Reactants and products
Conservation of mass
Factors affecting reaction rates
Topic 4 – My Learning Outcomes
When I have completed this section I will be able to:
Identify and critically evaluate if a new substance has been formed
Observe and describe evidence of chemical change in reactions
Distinguish between materials that react readily and those that do not
Describe chemical reactions and represent these reactions by using word equations
and chemical formulas and by constructing models of reactants and products.
Chemical Reactions
Writing and Identifying Chemical Reactions
Chemical changes result from chemical reactions
Starting materials – reactants
New materials created – products
Ex. Campfire – burning wood undergoes a combustion reaction. Wood and oxygen are the reactants.
The products are carbon dioxide and water
wood + oxygen  carbon dioxide + water + heat
Evidence of a chemical change (5 clues)
A colour change
A formation of an odour
The formation of a gas
The formation of a precipitate (solid)
The release or absorption o energy (heat/light/sound)
Common Types of Chemical Reactions
Exothermic & Endothermic Reactions
Chemical changes involving oxygen
Combustion reactions (burning)
Corrosion (rusting)
Cellular respiration (body cells producing energy)
Conservation of Mass in Chemical Reactions
 Products are formed when reactants undergo a chemical change
 Often look very different from reactants
Conservation of Mass states that:
 Matter is not created or destroyed in a chemical reaction
 The mass of the products is always the same as the total mass of the reactants
Ex. Combining 24.3g of Mg and 32.1 g of S creates 56.4g of MgS (magnesium sulfide)
Closed System – no product escapes
Open System – products escape and affect total product mass.
Factors Affecting the Rate of a Chemical
Reaction
There are four factors that can affect the rate of a reaction.
 Presence of a catalyst
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Help reactions proceed faster – ex. Enzymes
 Concentration
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Greater the concentration, the faster the reaction
 Temperature
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More heat added to the reactants, the faster the reaction
 Surface Area
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The greater the surface area of the reactants, the faster the reaction
Chapter 2 - Study Guide
Topic 1
Classroom Safety
Whmis
Stands for
Whmis symbols
Organizing Matter
Physical Properties
State, colour, lustre, malleability, solubility, density, hardness, boiling/melting points,
conductivity
Chemical Properties
Reaction with acids & waterbehavior in air, reaction to heat, ability to burn,
Pure Substances
Elements & compounds
Mixtures
Mechanical mixtures, solutions, suspensions & colloids
Physical Changes
Change state (Phase Change comp. doesn’t change)
Evaporation, melting, sublimation, condensation, freezing,.
Chemical Changes
2 or more materials react to create new materials
5 clues that a chemical reaction has occurred
Metals, Non-metals, Metalloids
Alkali Metals, Alkaline Earth Metals
Halogens , Noble Gases
Chapter 2 Study Guide
Topic 2
2.1 Theories of Matter
Timeline of discovery
Democritus and Aristotle
What were their ideas about matter
Alchemists
Lavoisier – naming chemicals
Atomic Theory
John Dalton – Billiard ball model
JJ Thomson – Raisin Bun Model
Nagaoka – Electrons in orbit
Ernest Rutherford – famous experiment, his theory
Neils Bohr – Bohr Model
Electorn Shells
James Chadwick
Protons & neutrons
Quantum Theory
2.2 Organizing the Elements
Mendeleev and the Periodic Table
Atomic Mass
Definition
Qhy is it important
2.3 The Periodic Table
Periods – horizontal
Groups (families) – Vertical
Element Names – capital then lower case
Atomic # vs Mass #
Mass # = protons + neutrons
Chapter 2 Study Guide
Topic 3
Forming & Naming Compounds
Chemical Names and Chemical Formulas
Ionic Compounds
Properties
Solid at room temp.
High melting points
Good conductor
Formed from metal (+) and non-metal (-) ions
Rules for Naming Ionic Compounds
Metal comes first and keeps name –Non-metal second – end of name changes to ide.
Balance ion charge to determine ratio of atoms.
Molecular Compounds
Properties
Solid, Liquid or gas at room temp
Relatively low melting/boiling points
Insulator (poor conductor)
Formed from two non-metals
Rules for Naming Molecular Compounds
First element keeps element name
Second element changes ending to ide
If there is more than one atom a prefix is attached
Exception to rule 3 when the first atom is only one we do not use the mono prefix.
Chapter 2 Study Guide
Topic 4
Chemical Reactions
5 Clues that a new substance has formed
A colour change
Formation of an odour
Formation of a gas (bubbles)
Formation of a solid (precipitate)
Release or absorption of energy (heat, light, sound)
Types of Reactions
Endothermic & Exothermic
Combustion
Corrosion
Cellular respiration
Law of Conservation of Mass
Reaction Rates Affected By
Catalysts
Concentration
Temperature
Surface Area