Transcript Matter and Chemical Change

1.1 Safety in the Science Classroom
WHMIS – Workplace Hazardous Materials Information System
We will go in to the lab to work on and PRINT out a
worksheet on WHMIS and lab safety.
1.2 Organizing Matter
Matter exists in three states:
 Solid
 Liquid
 Gas
Six changes of state are:
 Solid to liquid = Melting
 Liquid to solid = Freezing or Solidification
 Solid to gas = Sublimation
 Gas to solid = Deposition
 Liquid to gas = Evaporation
 Gas to liquid = Condensation
Matter has two types of properties:
 Physical
Describe the physical characteristics of a substance such
as color, lustre, density, hardness, crystal shape.
 Chemical
Describe how the substance reacts with other
substances. An example would be vinegar and baking
soda reacting.
Classification of Matter
Matter
Pure Substances
Mixtures
Elements
Compounds
Mechanical
Mixtures
Solutions
Suspensions
Colloids
TAKE A BREAK
Pure Substances
Elements
 A substance that cannot be broken down further. For
example: Gold
 These are found on the periodic table of elements.
Compounds
 A chemical combination of two or more elements in a fixed
proportion.
For example: NaCl (Sodium Chloride)
 Compounds have very strong chemical bonds. As the
elements combine, they lose their individual
characteristics.
 Compounds cannot be separated by physical means
such as evaporation.
Mixtures
Mechanical Mixtures
 You can see more than one part in the mixture
For example: Cookie dough
 These are also called heterogeneous mixtures.
 These can be separated by mechanical means such as using
hands or filters.
Solution
 A mixture where you cannot see the different parts
For example: Vinegar
 These are also called homogeneous mixtures.
 There is no settling in a solution.
 There are no chemical bonds in solutions. When the substances
join, they keep their individual characteristics.
 Solutions happen because the particles slip between each other
so they are evenly distributed.
 They can be separated by physical means such as boiling.
Suspensions
 Suspensions are heterogeneous mixtures that look
homogeneous.
 Over time, the particles settle out in a suspension.
 The particles can be separated by filters or leaving it
stand over time.
Example: Silty water
Colloids
 These are heterogeneous mixtures in which the
particles don’t settle out.
 Colloid particles can often be seen in a beam of light.
Example: Milk is a colloid
Homogeneous and Heterogeneous Video
1.3 Observing Changes in
Matter
Physical Change
 A physical change is a change in appearance or state of
a substance. Composition of the substance always
stays the same.
Example: Melting, crystallization
Chemical Change
 This occurs when two or more substances react and
create new substances.
Example: Vinegar and baking soda reacting
Evidence of chemical change:
1.
Change in color – Dying clothes
2. Change in odour – Odour given off when striking a
match
3. Forming a solid or gas – Mixing vinegar and baking
soda you get carbon dioxide gas
4. Release or absorption of heat energy – When
gasoline burns in an engine it produces heat
However:
It is not always easy to tell if physical or chemical change
has occured. Sometimes a chemical analysis needs to
be done to determine the type of change.
Controlling Changes in Matter to Meet Human Needs
 Use of freeze-drying to make foods easier to prepare
and transport
 Chemicals from corn can make environmentally
friendly pop bottles, remove nail polish and make
ethanol fuel for cars
2.1 Evolving Theories of Matter
8000 BC - Stone Age
 Matter consisted of solid materials which could be
made into tools
Early Chemists
Between 6000 BC and 1000 BC
 Early chemists produced valuable metals such as gold
and copper
4500 BC - Bronze Age
 Heating copper led to the creation of bronze for use in
tools
1200 BC – Iron Age
 Iron was combined with carbon to make steel for
stronger tools
 The Hittites in the Middle East first discovered how to
extract iron from rocks and melt it into pots and
swords
500 - 350 BC
Democratis
 Greek philosopher that believed that matter could be
broken down into indivisible bits called “atomos”
Aristotle
 Another Greek philosopher that believed that
everything was made of earth, air, fire and water.
1500 AD
Alchemists
 alchemists were the first modern chemists although
they weren’t scientists
 Alchemists believed it should be possible to change
metals into gold
 they conducted the first experiments
Modern Chemistry
1600’s - Robert Boyle
 Experimented with gases and what happened to them
under pressure
 Through experiments he became convinced that
matter was made of tiny particles
1780’s - Antoine Lavoisier
 Known as the “father of modern chemistry,” he
developed a system for naming chemicals
1808 - John Dalton
 First person to define an element as a pure substance
 Came up with the “billiard ball” model of the atom,
which was the first theory of atomic structure
1897 – J.J. Thomson
 Discovered electrons and came up with the “raisin
bun” or “plum pudding” model of the atom: This
showed the atom as a positive sphere with electrons
(negative charges) embedded in it
1904 – Hantaro Nagaoka
 proposed the “solar system” model of the atom: This
showed the atom as a positive sphere with electrons
orbiting it
1922 – Niels Bohr
 Develops the “electron shell” model where electrons
move in specific shells or orbits around the nucleus
1930’s – James Chadwick
 Discovered the proton (positive charge) and the
neutron (no charge) in the nucleus
Today
 The quantum mechanics model of the atom is used
today: This says that an atom is a cloud of electrons
surrounding a nucleus
Electrons, Neutrons, Protons, Atoms and Molecules
 Electrons - Negatively charged subatomic particles found
in the cloud region of atoms
 Protons - Positively charged subatomic particles found in
the nucleus
 Neutrons - Neutral subatomic particles found in the
nucleus
 Atoms – The basic building blocks of all matter.
 Molecules – A group of at least two atoms held together by
chemical bonds. (NaCl = this is a molecule of sodium)
ATOMS EXPLAINED
2.2 Organizing The Elements
Looking for Patterns
 Early Greeks first used Celestial bodies for symbols
 In the 1800’s John Dalton came up with new symbols
for the elements
 In 1814 Jons Jacob Berzelius created the system still
used today.
 He used letters for the elements instead of symbols
Ex. O for Oxygen and He for Helium
An Order for the Elements
 It was realized that elements could be listed in order of
increasing atomic mass
 Atomic mass is the mass of one atom of an element
 By comparing all elements to Carbon, the relative
atomic masses of elements can be figured out
 John Newlands recognized patterns in properties when
elements were listed in order of atomic mass. He said
properties of elements seemed to repeat every 8
intervals on the list
Finding a Pattern
 Dmitri Mendeleev then came up with a pattern for the
elements based on their properties
 He listed the 63 known elements at the time according
to their properties and left gaps for elements he was
sure would be discovered in the future
2.3 The Periodic Table Today
Periodic Tables Today
 Periodic tables today can have a variety of information
on them but they all have the same basic set-up:
Atomic
Number
Ionic
Charge
Element
Symbol
Atomic
Mass
Element
Name
Element Names
 The elements have names that come from: Greek,
Latin, geographic location, planets, minerals, famous
scientists
 Modern symbols are 1-2 letters (first is always a
capital, second is always lowercase)
 The symbol doesn’t always match the English name for
an element
Example: Pb is lead
TAKE A BREAK
Periods or Orbits
 The periodic table is organized into 7 horizontal rows
called periods or orbits
 Atomic mass increases when moving left to right
across a period
 The properties of elements change as you move across
a period but the same pattern of change is repeated in
each period
Groups or Families
 The table is also organized into 18 vertical columns
called groups or families
Example: Alkali metals, alkaline-earth metals,
halogens etc
 The elements in a group tend to have similar
properties
Example: Noble gases (group 18) are typically unreactive
Types of Elements
Non-metal
Nonmetals
Metals
TAKE A BREAK
Types of Elements
 The table has a “staircase” line that separates elements with
certain properties .
 Metals are on the left side of the staircase…they are shiny,
malleable, ductile and conduct electricity.
 Alkali metals are the most reactive metals.
 Non-metals are on the right side of the staircase…they are
dull, brittle and nonconductive. They can be a solid or gas.
 Halogens (group 17) are the most reactive non-metals.
 Noble gases are in group 18 and are the most
stable and un-reactive elements.
 Metalloids are directly above and below the
staircase…they show properties of both metals and
non-metals.
Atomic Number vs. Atomic Mass
 Atomic number is the number of protons (p+) in an
element
Example - Cl (chlorine) has 17 p+
 Atoms are neutral (not charged) which means the
number of protons is equal to the number of electrons
(e)
 Atomic mass = mass of protons (p+) + mass of
neutrons (n0)
 Mass number = # of p+ + # of n0
 To figure out the mass number of an element, round
the atomic mass (on periodic table) to the nearest
whole number
Example - Sodium (Na) has an atomic mass of 22.99
g/mol, so its mass number is 23
How to calculate the number of neutrons
 To calculate the number of neutrons (n0) in an
element:
# n0 = mass number – atomic number (#p+)
Example: Ne (Neon)
Atomic Number = 10
Atomic Mass = 20.18
Mass Number = 20
Neutrons = 20-10 = 10
3.1 Naming Compounds
Chemical Formulas
 Chemical formulas tell us two things:
What elements are in the compound
2. How many atoms of each element are in the
compound – determined by the script number next
to each symbol.
1.
Hydrogen peroxide
symbol for hydrogen
Subscript of “2” after
“H” means there are 2
atoms of hydrogen.
symbol for oxygen
Subscript of “2” after
“O” means there are 2
atoms of oxygen.
Water
Symbol for hydrogen
Subscript of “2” after
“H” means there are 2
atoms of hydrogen.
Symbol for oxygen
No subscript after “O”
means there is 1 atom
of oxygen.
Physical States
 States of matter can be written in brackets beside the
compound
Example: CH4(g)
(s) means it is in the solid state
(l) means it is in the liquid state
(g) means it is in the gas/vapour state
(aq) means it is dissolved in water
Check and Reflect
On page 143 please answer the following questions:
Questions 1,2,3 and 4 
* You also need to finish the worksheet provided (you
can work in a group if you wish)
3.2 Ionic Compounds
Ions
 An ion is an atom or group of atoms that has become
electrically charged
 Ionic compounds are pure substances that form as a
result of the attraction of positive and negative ions
Ionic Compounds
They:
 Are solids at room temperature
 Have defined crystal structure
 Conduct electricity when liquid or dissolved in water
 Have high melting points
Ions are formed when atoms either lose or gain
electrons…atoms lose or gain electrons to become
more stable. Noble gases are the most stable elements.
Metals:
 Lose electrons and form positive ions called cations
 Are named by their metal name followed by the word
“ion”
Example: Sodium ion, lithium ion
Non-metals
 Gain electrons and form negative ions called anions
 Are named by changing the ending of the non-metal
to “ide” then followed by the word “ion”
Example: Oxide ion, chloride ion
Ion charge can be found on the periodic table
Ion charge
Writing the symbol
 When writing the symbol for an ion, you write the
symbol for the element then write the charge as a
superscript (number first unless it is 1, then sign)
Example:
 sodium ion is Na+
 calcium ion is Ca2+
 oxide ion is O2 nitride ion is N3-
Ionic Bonds
 Ionic bonds form when positive and negative ions
(charges) attract each other
+ charge attracts the  charge
[Na]+
ionic bond
[Cl]
3.3 Molecular Compounds
 A molecular compound forms when two or more
non-metals join together
 They are joined by the sharing of electrons which is
called a covalent bond
Molecular compounds:
 Can be solids, liquids or gases at room temperature
 Do NOT conduct electricity when liquid or dissolved
in water
 Have low melting points
4.1 Chemical Reactions
Writing Chemical Equations
A chemical reaction occurs when
 a substance or substances react to produce a new
substance
A chemical equation is used to show the chemical
reaction
reactants
products
A + B  C + D
reaction direction
Reactants are
 the substance(s) that you start with and are used up
Products are
 the substance(s) that you make or produce
You can write a chemical formula in two ways:
as a word equation – Hydrogen + Oxygen Water
vapour
2. using chemical formulas and states of matter –
H2 (g) + O2 (g)
H2O (g)
1.
TAKE A BREAK
Exothermic Chemical Reactions
 a chemical reaction that releases energy is called
exothermic
 energy released can be heat, light, sound, electricity
Example – Burning wood
Endothermic Chemical Reactions
 a chemical reaction that absorbs or requires energy is
called endothermic
 many require heat but could also require electricity or
UV radiation
Example – Cooking food or recharging a battery
TAKE A BREAK
Reactions with Oxygen
1.

Combustion
occurs when oxygen reacts with a substance to
produce a new substance and energy
2. Corrosion

occurs when oxygen in the atmosphere gradually
reacts with a metal producing rust
3. Cellular Respiration

occurs in only in cells and is the reaction of organic
material (food) with oxygen
4.2 Conservation of Mass in
Chemical Reactions
The law of conservation of mass:
 states that the total mass of the reactants is equal to
the total mass of the products in every chemical
reaction
 when a chemical reaction takes place, matter is not
created or destroyed, it is only rearranged
A closed system:
 is a system where matter cannot escape
Example: a reaction in a sealed flask
An open system:
 is a system where matter can escape
Example: a reaction in an open beaker
4.3 Factors Affecting the Rate
of Chemical Reactions
Reaction rate:
 the speed at which a chemical reaction takes place
There are four factors that affect the rate of a reaction:
1. Presence of a Catalyst:
 substances that help a reaction proceed faster
without being used up themselves
 can be found in non-living and living systems
Example – enzymes in your body
2. Concentration of the Reactants
 a higher concentration of reactants means there are
more particles around to bump into each other and
react
3. Temperature of the System
 higher temperature speeds up the motion of the
particles which makes them bump into each other
and react
4. Surface area of the Reactants
 a higher surface area of the reactants means more
particles are exposed and able to react with other
reactant particles