Transcript Energy and Matter in Chemical Change Science 10x

Energy and Matter in
Chemical Change
What is Chemistry?
The study of matter; its
properties, composition
and the changes it
undergoes.
Why Chemistry is Cool
It has the best jokes…
Only the coolest
people do
chemistry…
http://www.youtube.com/watch?v=bYtQ3_eh
094&feature=related
A Walk Down
WHMIS
Compressed Gas
Toxic
Biohazardous
Flammable
Poisonous
Oxidizing
Corrosive
Radioactive
HHPS
Danger
H azardous
H ousehold
P roduct
S ymbols
Warning
Caution
MSDS
M aterial
S afety
D ata
S heets
Science Lingo…
• Scientists use an experiment to search for cause
and effect relationships in nature. In other words,
they design an experiment so that changes to one
item cause something else to vary in a
predictable way.
• These changing quantities are called variables. A
variable is any factor, trait, or condition that can
exist in differing amounts or types. An
experiment usually has three kinds of variables:
independent, dependent, and controlled.
• The independent (manipulated) variable is the one
that is changed by the scientist. To insure a fair
test, a good experiment has only one independent
variable. As the scientist changes the independent
variable, he or she observes what happens.
• The scientist focuses his or her observations on the
dependent (responding) variable to see how it
responds to the change made to the independent
variable.
Let’s Look At An Example…
• For example, if you open a faucet (the
independent variable), the quantity of water
flowing (dependent variable) changes in
response--you observe that the water flow
increases. The number of dependent variables
in an experiment varies, but there is often
more than one.
• Controlled variables are quantities remain constant.
For example, if we want to measure how much
water flow increases when we open a faucet, it is
important to make sure that the water pressure
(the controlled variable) is held constant. That's
because both the water pressure and the opening
of a faucet have an impact on how much water
flows. If we change both of them at the same time,
we can't be sure how much of the change in water
flow is because of the faucet opening and how
much because of the water pressure. In other
words, it would not be a fair test. Most experiments
have more than one controlled variable. Some
people refer to controlled variables as "constant
variables."
An Example…
• A scientist is conducting an experiment to test if
taking vitamin A could extend a person’s lifeexpectancy.
Another Example…
• A scientist wants to see if temperature affects
the reaction rate of melting M&M’s. She sets
up three beakers with different temperatures
of water; hot, luke warm and ice cold.
Summary
• Independent (Manipulated) Variable: What
we are changing
• Dependent (Responding) Variable: The
Response (what happens)
• Controlled Variable: What remains the same
• A hypothesis is an educated guess about how
things work.
• Your hypothesis should be something that you
can actually test, what's called a testable
hypothesis.
• In other words, you need to be able to measure
both "what you do" and "what will happen."
Scientific Method
• What is the scientific method?
– It is a process that is used to find answers to questions about
the world around us
• Is there only one “scientific method”?
– No there are several versions of the scientific method. Some
versions have more steps, while others may have only a few.
However, they all begin with the identification of a problem
or a question to be answered based on observations of the
world around us and provide an organized method for
conducting and analyzing an experiment.
Hypothesis
• What is a hypothesis?
– It is an educated guess based on observations and
your knowledge of the topic.
• What is data?
Data
– It is information gathered during an
experiment
Identify
Form a Hypothesis
Create an Experiment
Perform an Experiment
Analyze Data
Is data accurate?
Modify the Experiment
Communicate the Results
Scientists
Aristotle
• 400 BC
• Fire, Earth, Water and Air
Democritus
• Believed that material was made up of small
individual particles called “atoms”
Alchemists
Lavoisier
• “Father of Modern Chemistry”
• Came up with the Law of Conservation of
Mass- “Mass is neither created nor destroyed”
The other person who did a lot of significant scientific work but did not
get credit because of her gender
Dalton
• In 1804 Dalton stated that atoms were
– Tiny individual particles
– All identical in a given element with identical
properties
– Two or more elements atoms will combine in fixed
ratios
– “Billiard Ball Model”
Dalton’s Matter Model
• All matter is made of atoms
• All atoms of an element are identical in
properties
• Atoms of different elements can combine in
specific ratios to form new substances
Law of Definite Proportions
• A compound always has the same ratio of
atoms, regardless of how it is made.
Example: water is always H2O, not HO, or HO2, or H2O2
30
J.J. Thompson
• Discovered that atoms had negative particles
(electron) contained inside the atom
• “Raisin Bun Model”
Nagoaka
• Placed all the + charges in the centre of the
atom and all the – charges in a ring around the
centre.
• “Saturn Ring” model
Rutherford
• Believed that electrons were outside the
dense positive part of the atom (nucleus)
• Used gold foil experiment
Chadwick
• Discovered the subatomic particles:
– Protons: p+
– Electrons: e– Neutrons: n°
Bohr
• Thought that electrons were arranged in
certain energy levels around a positive nucleus
• Electrons in orbitals
• “Bohr Model”
de Broglie and Schrodinger
• Electron cloud model
• Electrons are not considered to be point
charges but as a mist around the nucleus
• Theory has evolved into the quantum
mechanics model
http://www.youtube.com/watch?v=7SjFJImg2Z8
The Latest and Greatest- Quantum
Mechanics
• The latest experimental evidence has
disproved Bohr’s idea of fixed energy levels
• Now, they visualize electrons not as a particle,
but as a cloud of negative charge. Rather than
following little race tracks around the nucleus,
they occupy the whole space all at once at
different energy levels
• NOW GET YOUR HEAD AROUND THAT!
Modern Atomic Structure
• Current models are more complex and
involve many subatomic particles such as
quarks (protons and neutrons)
38
Your Task
• Check & Reflect p 25 # 1,5-8 & 11
• Section Review page 27 # 5, 10, 11, 12, 15, 16
• Expect a scientist quiz next class
Matter
• Matter: Anything that has mass and takes up
space.
States of Matter
• There are four states of
matter:
– Solid
– Liquid
– Gas
– Plasma
Particle Model of Matter
5 Main Points:
1) All substances are made of matter
2) All particles in a pure substance are the same.
Different pure substances are made of different
particles.
3) Particles are always in motion. The speed of the
particles increases when temperature increases.
4) Particles have space between them
5) Particles may have attractive forces between them
Organizing Matter
Pure Substances
• Made up of only one kind of matter and
has a unique set of properties.
• Can only be broken down by a chemical
decomposition
Elements
• pure substance that cannot be separated into
simpler substance by physical or chemical
means.
• Basic building blocks for all compounds
• Each element has it’s own symbol
Compounds
Pure substance composed of two or more different
elements joined by chemical bonds.
– Made of elements in a specific ratio
that is always the same
– Has a chemical formula
– Can only be separated by
chemical means, not physically
Mixtures
• A combination of two or more pure substances that
are not chemically combined.
• substances held together by physical forces, not
chemical
• No chemical change takes place
• Each item retains its properties
in the mixture
• They can be separated physically
Chem4kids.com
Heterogeneous Mixtures
• Heterogeneous mixtures are also known as
Mechanical Mixtures
• In a mechanical mixture, the different
substances that make up the mixture are
visible.
• Includes: Mechanical mixtures, suspensions,
emulsions and colloids
Homogeneous Mixtures
• A Homogeneous Mixture is also known as a
Solution
• In a solution, the different substances that
make it up are not separately visible.
• One substance is dissolved in another.
• Recall that a Solution is made up of a Solvent
and a Solute
Suspensions
• A suspension is a cloudy mixture where tiny
particles of one substance are held within
another.
• These particles can be separated out when the
mixture is poured through filter paper.
Colloids
• A colloid is also a cloudy mixture, but the
particles of the suspended substance are to
small that they cannot be easily separated out
from the substance.
Test Your Understanding
• A black solid with a constant melting point is
heated to a high temperature, producing a gas
and a shiny brown metal. The boiling point of
the gas is -183°C and the melting point of the
metal is 1085°C. Is the black solid an element,
compound or mixture? Explain
• Let’s break down the evidence:
– Constant melting point (not a mixture)
– But both a gas and a solid were formed. So it is
not an element
 Compound
Separation of Mixtures
• Separating Mixtures to identify the ingredients
can be accomplished through the following
procedures.
– Mechanical Mixtures: sifting or filtering
– Solutions: distillation
– Suspensions: filtration and centrifuging
– Colloidal Solutions: centrifuging
– Compounds: chromatography
Phase Changes
Physical vs. Chemical Change
Physical
Chemical
http://cwx.prenhall.com/petrucci/chapter1/medialib/tutor/f20/0103.html
Chemical Change
• A chemical change always results in the
formation of a different substance or
substances.
Physical Properties
2 Categories of Physical Properties:
1) Qualitative Properties ( sensed/ observed)
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•
•
•
•
Color
Texture
Smell
Taste
Malleability
Texture
Change in state
Ductility
Physical Properties
2) Quantitative Properties (measured)
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•
•
•
•
Melting point (mp)
Boiling Point (bp)
Density  d= m/v (mass/volume)
Solubility
Viscosity (flow rate)
Conductivity
Chemical Reactions
• Characteristics
– Energy change (temperature, light, sound, electricity)
– Odor Change (appearance or disappearance)
– Color change
– Formation of a gas
– Formation of a solid (precipitate)
http://www.youtube.com/watch?v=66kuhJkQCVM
Test Your Understanding
• A blue crystal is placed in water and after
stirring it disappears and the water becomes
blue. The liquid is then heated and the water
evaporates and small blue crystals appear. Did
a chemical reaction take place?
=
Answer
• NO. The crystal dissolved and then the water
evaporated. This was only a phase change not a
chemical change. No new substance was
created.
And Now You Know!
Test Your Understanding
• If a marshmallow is cooked over a flame and
becomes black on the outside, did a chemical
reaction occur? Defend your answer
=
Answer
• YES! A new color appears (black crust) and a
new substance was produced (carbon crust)
Change
Frying and Egg
Boiling Water
Leaves changing in the
Fall
Rusting Iron
Physical
Chemical
Evidence
And One More Time…
1) Color change
2) Gas is formed
3) New odor is formed
4) Precipitate (solid) is formed
5) Temperature rise or drop
Your Task
• Create a set of notes summarizing the
differences between chemical and physical
changes/reactions
• Check & Reflect p 25 # 1,5-8 & 11
• Section Review page 27 # 5, 10, 11, 12, 15, 16
The Periodic Table
http://www.youtube.com/watch?v=lxJe7e5thkI&feature=relmfu
http://www.youtube.com/watch?v=0YmypUfvJvY
Elements
• The number of protons in the nucleus and the
distribution of electrons around it determines the
type of element and the chemical and physical
properties of the element.
Eg. hydrogen is 1 p surrounded by 1e- bromine is 35 p and 45 n
surrounded by 35 e-
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Elements
• Elements consist of three main classes
– 1) metals
– 2) non-metals
– 3) metalloids
• These classifications are based primarily on
conductivity
73
Elements
• Most elements occur naturally as single atoms
(monoatomic) eg. Ne
• Other elements occur naturally as
combinations of two or more atoms (molecular
elements)
– diatomics eg. O2, N2, H2, Cl2,
– polyatomics eg. P4, S8
74
Relationship of Chemical Symbols to Chemical
Names
• Alchemists used artistic symbols to
represent the elements
• Dalton updated the symbols but they were
still cumbersome
75
Relationship of Chemical Symbols to Chemical
Names
• Baron Jons Jakob Berzelius (1779 -1844)
developed the letter system in
1814 which was basically the same
as the modern system except it
used superscripts instead of
subscripts.
76
International Union of Pure and
Applied Chemistry -IUPAC
• Founded in 1919 IUPAC maintains an
international system for naming chemicals
77
International Union of Pure and
Applied Chemistry -IUPAC
• IUPAC develops rules, guidelines and
standard conventions for the study of all
aspects of Chemistry
• IUPAC allows chemists to communicate
clearly and precisely
78
Modern Atomic Symbols
• The same in every language.
• First letter upper case, second letter
lower case (eg Co not CO)
(if any)
NOTE: Figure A2.3 p 30
Memorize element names and symbols of rows1- 4
[The elements most prevalent in living things are H, C, N,
O, S, & Ca.]
Worth of the Human Body link1
Worth of the Human Body link2
79
Periodic Table
• John Newlands (1837-1898)
• English scientist organized the 62
known elements on basis of
atomic mass four years before
Mendeleev (1864).
H1
F8
Cl 15
Co/Ni 22
Br 29
Pd 36
I 42
Pt/Ir 50
Li 2
Na 9
K 16
Cu 23
Rb 30
Ag 37
Cs 44
Tl 53
Gl 3
Mg 10
Ca 17
Zn 25
Sr 31
Cd 34
Ba/V 45
Pb 54
Bo 4
Al 11
Cr 18
Y 24
Ce/La 33
U 40
Ta 46
Th 56
C5
Si 12
Ti 19
In 26
Zr 32
Sn 39
W 47
Hg 52
N6
P 13
Mn 20
As 27
Di/Mo 34
Sb 41
Nb 48
Bi 55
O7
S 14
Fe 21
Se 28
Ro/Ru 35
Te 43
Au 49
Os 51
80
Periodic Table - reference p 31
• Dimitri Mendeleev (1869) &
Lothar Meyer (1870)
individually arranged know
elements in periodic
repetition of characteristics
• The success of this model was
that it was predictive
81
The Periodic Table, Your Tool in Chemistry
• The elements on the periodic table are
organized according to their atomic number
• There are about 115 elements known
– Only 90 are naturally occurring
– All elements are divided into one of three
categories
Metalloids
Trends
• The periodic table is arranged according to
three
• Main things:
a) Increasing atomic mass
b) Grouped in families because they have
similar chemical properties
c) Reactivity increases from L to R and
from Top to Bottom
• Each horizontal row is called a period
– Grouped together because of energy levels
– Numbered 1-7
• Each vertical column is a group or family
– All members of a family have similar chemical
properties
– Numbered 1-18
Group (Medeleev used the term families)
• Vertical columns, elements with similar chemical
and physical properties
(Alkali metals, alkali earths, halogens, & noble gases)
- Group 1 & 2 are classified with increasing
reactivity from top to bottom
- Group 17 decreases in reactivity from top to
bottom (fluorine is one of the most reactive substances known)
(Note: Noble gases generally unreactive)
86
Period
• Horizontal rows with elements arranged so
that their properties repeat periodically
across the table
• Each new row represents repeating trends
in reactivity
87
Groups or Families
• Notable groups:
- Alkali metals
- Alkaline-earth metals
- Noble gases
- Halogens
Alkali Metals
• Soft, shiny, silver in colour
• Very reactive with water
• Compounds tend to be white solids that are
soluble in water
• They have 1 valence electron (e-)
* Note that Hydrogen does not belong to the
alkali family
Alkaline Earth Metals
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•
•
•
Shiny, silver, but not as soft as alkali
Compounds white, but less soluble than alkali
Tend to react with oxygen
2 valence e-
Halogens
• Non-metals
• Poisonous and react with alkali metals to form
salts
• Cl and F (g) Br (l), I (s)
• 7 Valence e-
Nobel Gases
• Colorless gases
• Very unreactive
Other Groups
Group 3-12: Transition Metals
• They tend to be shiny, malleable and conduct
thermal energy and electric current well
Staircase line
• Separates metals & non-metals
• Metalloids border the staircase
94
Metalloids
• Solids at room temperature
• Have properties of metals and non-metals
• Eight metalloids are: B, Si, Ge, As, Sb, Te, Po
At border the staircase line of the periodic
table
• Important in semi-conductor industry
95
Lanthanides
Lanthanides Elements in the first row that
follow lanthanum are called lanthanides. These
are also called rare-earth and inner-transition
metals. These can be found naturally (though
rarely, considering the name!) on earth.
Actinides
• Elements in the second row that follow
actinium are called actinides.
• These are all radioactive and some are not
found in nature. Some of the elements with
higher atomic numbers have only been made
in labs
Some Things to be Aware of…
• Group 1 elements always donate 1 electron
• Group2 elements always donate 2 electrons
• Group 17 elements accept 1 electron
• Group 16 elements accept 2 electrons
Types of Elements
Metals:
– Solid at room temperature (except mercury)
– Most are lustrous, malleable and ductile
– Most will react with non-metals
– They are good conductors of electricity
– Exist to the left of the “staircase”
Characteristics of Metals
• 1. Conduction: Metals are good at conducting electricity. Silver
(Ag) and copper (Cu) are some of the most efficient metals and
are often used in electronics.
2. Reactivity: Metals are very reactive, some more than others,
but most form compounds with other elements quite easily.
Sodium (Na) and potassium (K) are some of the most reactive
metals.
3. Chemical: Metals usually make positive ions when the
compounds are dissolved in solution. Also, their metallic oxides
make hydroxides (bases) (OH-), and not acids,
when in solution.
4. Alloys: Metals are easily combined.
Mixtures of many metallic elements are called
alloys. Examples of alloys are steel and bronze.
Alloys
• solid solutions which
usually have a metal as the
solvent and a metal or a
non-metal as the solute.
Eg. Steel (Fe, C),
solder (Sn, Pb/Ag),
bronze (Cu, Sn),
amalgam (Hg, Ag, Sn)
101
Types of Elements
Non-metals:
- Find them right of the “staircase”
- Can exist as solid, liquid, or gas at room
temp
-Appearance varies, not very shiny
-Poor conductors of electricity
- Brittle/not ductile
Element Song - YouTube
Your Task
• Periodic table worksheet
• Read page 28-39 make sure to add notes for
understanding on the following:
–
–
–
–
–
SATP
Metals as elements
Non-metals as elements, molecules that are elements
Metalloids as elements
The Modern Periodic Table and
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•
•
•
•
•
•
•
•
Families or groups
Periods
The staircase
The table key
Alkali metals
Alkaline earth metals
Halogens
Noble gases
salts
Atomic Theory
Atom Song – YouTube
Atoms
• An atom is made up of three types of
subatomic particles:
– Protons  p+
– Electrons  e –
– Neutrons  n°
These subatomic particles were
discovered by…
Atoms
• The protons and neutrons are contained in the
nucleus, while the electrons are outside the
nucleus
• Neutral element (no charge)
– Electrons and protons are the
same
Organization
Atomic Number
• Elements are defined by the number of
protons contained in the nucleus
• The number of protons found in the nucleus
of an atom is its atomic number
Example:
Hydrogen (H) - 1 p - 1H
Gold (Au) - 79 p - 79Au
109
Isotopes
• Isotopes: differing numbers of neutrons in an
atom but the same number of protons
– Results in a different atomic mass
Note:
• All three isotopes of hydrogen contain 1 proton in the nucleus
• All three isotopes of hydrogen have 1 electron outside the nucleus
• The three isotopes of hydrogen differ in the number of neutrons in the nucleus
http://www.youtube.com/watch?v=BYX312koKps
Isotopes
PROTON AND ELECTRONS STAY
THE SAME
BUT
THE NEUTRON # IS DIFFERENT
112
113
Example
• Let’s consider the naturally occurring isotopes of
chlorine:
• 75.53% of all atoms are 3517 Cl isotopes; the
atomic number is 17 and the atomic mass is 35
• 24.47% of all atoms are 3717Cl isotopes; the
atomic number is 17 and the atomic mass is 37
• Note the atomic mass changes because the
number of neutrons increases the mass
Example:
• An isotope with 45 proton and 48 neutrons
would be what?
Symbol
• Ms. Godley remember to talk about how to
symbolize isotopes
Naming Isotopes
• “atom name” – “mass number”
Example:
Oxygen-18
# of e =
# of p =
# of n =
Symbol =
Silver-200
# of e =
# of p =
# of n =
Symbol =
Copper-150
# of e =
# of p =
# of n =
Symbol
Isotopes
• The average mass of each atom is what is
recorded for each element in the Periodic
Table.
• All elements in the Periodic Table have
isotopes that results in atomic numbers that
are not
whole numbers
Energy Levels - aka Energy Shells
• Electrons in atoms can only be at certain
energy levels. [cf. stairs]
• Transitions between energy levels involves
either a gain ("up") or loss ("down") of energy.
• Energy levels closest to the nucleus are always
filled up first
120
Energy Levels
The maximum number of electrons in each level is:
Level
1
2
3
4
# of electrons
2
8
8
18
This level is filled 1st
Bohr
Diagrams
*Lower levels must be filled before higher levels get any
electrons.
121
Let’s Reintroduce Ourselves to Bohr
Diagrams
Drawing E Level Diagrams
• Draw E level diagrams for various atoms on
board and state how many valence electrons
they have
123
Energy Levels & Chemical Properties
• Why would sodium with 11 electrons have
similar chemical properties as potassium with
19 electrons?
• Draw E level diagrams for each!
124
A Little Chemistry Humor…
• Why does hamburger have lower energy than
steak?
Because it's in the ground state.
Valence e- & Octet Rule
• The electrons occupying the outer- most E level are
called valence electrons
• Members of the same group (family) have the same
number of valence e-
127
Valence e- & Octet Rule
• Electrons are added as you move from left to right
in a period
• Notice that the period number is the same as the
number of occupied energy levels
• Notice also that the group number is the same as
the number of electrons in the valence shell (Note: if >
10 remove the 1)
128
Valence e- & Octet Rule
• Noble gases are stable because they have
completely filled energy shells
• Atoms bond in such a way as to create a stable
octet (duet)
[illustrate using NaCl]
129
Your Task
• Draw energy diagrams and valence diagrams
for 6 elements of your choice. You must show
all of your work/steps
Ions
Let’s Get a Few Things
Straight
• Atoms = Proton # = electron #
neutral charge
Strive to be like Noble Gases
Ions & Ionization - reference p 34 - 36
• Atoms "try" to become stable by having their
outer energy levels filled with electrons (like noble
gases). To do this they must either:
a) gain or lose electrons (ionization), or
b) share electrons with other atoms.
Eg.
•
Na loses 1 e-  Na+ ion.
Cl gains 1 e-  Cl- ion.
[draw Li and F ion E diagrams]
133
Ions & Ionization
• When atoms gain or lose e-, they become charged,
forming ionic species or ions
• Non-metal atoms tend to gain electrons to have a
full outer energy level. Negative ions are called
anions. [anion = a negative ion]
Eg. Cl + 1e-  Cl-
134
Ions & Ionization
• Metals tend to lose electrons to have a full outer
energy level. Positive ions are called cations.
Eg. Na(s)  Na+ + 1e-
135
• Summary:
Ions
– Ions are formed by the gain or
loss of electrons
– Anions: gain electrons
(negative)
– Cations: lose electrons
(positive)
– Non-metals gain electron(s)
to become negative (anions)
– Metals lose electron(s) to
become positive (cations)
– Members of the noble gas
family do not ionize. Why????
Naming Ions
• Metals do not change names when ionized
– Example: Sodium is still called sodium in a
compound. We simply add the word ion if it is
alone
• Non-metals have the last three letters of their
names changed to “ide:
– Example: Chlorine becomes chloride
Multivalent Elements
• Some atoms can exist as more than one ion
type
• On the periodic table the most common form
of the ion is listed first
Example: Cu2+ and Cu1+
138
Recall that…
• Ionic Compounds form when we transfer
valence electrons
• Molecular Compounds form when we share
valence electrons
Lewis Dot Diagrams
• Only looks at Valence electrons
Rules:
1) The element symbol represents the nucleus and
inner-filled energy levels
2) A “dot”, “x”, or “O” represents valence electrons
3) There are 4 valence orbitals of the valence energy
level
4) Each valence electron is distributed like compass
points but must fill up 1 each before doubling up
Lewis vs. Bohr
Your Task
• Check & Reflect p 39
#1 – 12
• Draw 5 Lewis dot
models NOT in the
examples and must be a
representative element
(#1-20)
• Practice Questions in
notes
Compounds
• Categories of Compounds:
1) Ionic (metal + non-metal)
2) Molecular (non-metal + non-metal)
3) Intermetallic (metal + metal combinations)
Compounds
1) Ionic Compounds: metal/non-metal
combination
-Positive charge and negative charge combinations
-They are all solid at SATP (standard ambient
temperature and pressure)
-Dissolve in water to form electrolytes
Conduct electricity
Take
Electrons
Ionic Compounds
Identify the ionic compound(s)
• Na3N
• CS2
• SbBr3
• H2O
End in “OH”- hydroxide
Bases are Ionic Compounds
•
•
•
•
Feel slippery
Turns litmus paper blue
Have a bitter taste
Have a pH greater than 7.0
Also known as
Covalent Compounds
Compounds
Share
electrons
2) Molecular Compounds: non-metal and
non-metal combinations (no charges!)
- They can be solid, liquid or gas @ SATP
-Do not conduct electricity
Acids are Molecular Compounds
• Unlike other molecular substances in water,
acids DO conduct electricity
• pH less than 7
• Taste sour
H+  Hydrogen ion
Recall…
• Acids and bases will neutralize each other to
form salts and water
Compounds
3) Intermetallics: metal and metal combinations
• Made of two or more different metal atoms
Example: brass copper + zinc
bronze  copper + tin
Compounds
Ionic Compounds
- Solid at SATP
- aq (soluble in
water)
- Conduct
electricity
- pH greater than 7
- Red blue
- Bitter, slippery
Molecular
Compounds
-
aq
Conduct electricity
pH less than 7
Blue  red
Sour
Corrosive
Chemical Bonds
• Forces of attraction between atoms
Ionic: Taken (stronger)
Molecular: Shared
155
Types of Bonds
Ionic Bonds
• An ionic bond is formed by the attraction of
oppositely charged atoms or groups of atoms.
When an atom (or group of atoms) gains or
loses one or more electrons, it forms an ion.
Remember:
-Cation = positive
-Anion = negative
Types of Bonds
Covalent/ Molecular Bonds
Let’s Draw a
Lewis Dot to
Explain!!!!
• A covalent chemical bond results from the
sharing of electrons between two atoms
• A single covalent bond represent the sharing of
two valence electrons (usually from two different
atoms).
• The Lewis structure below represents the
covalent bond between two hydrogen atoms in a
H2 molecule.
Summary
• Ionic Compounds form when we transfer
valence electrons
• Molecular Compounds form when we share
valence electrons
Remember…
• Metals do not change names when ionized
• Non-metals have the last three letters of their name
changed to “ide”
• Composed of metallic cations ( +charged) and a non
metal anion ( - charge)
• Form because of the attraction between + and charges
Naming Compounds
• Recall that compounds are pure substances which
are combinations of two or more elements.
• To form compounds, at least two atoms must form
ions. One atom will lose an electron, while the
other will gain the electron
• Compounds can be classified by the type of
chemical bonds they have:
1. Ionic
2. Molecular
160
http://www.youtube.com/watch?v=9Zl5Y4earhM
Two types of Compounds:
1. Compounds which are combinations of
metals and non-metals are called ionic
compounds.
• Number of electrons lost will always equal the
number of electrons gained  electrically
neutral
• Subscripts are used to represent the number
of ions in a compound
2. Combinations of non-metals with other
non-metals are called molecular (covalent)
compounds.
161
Ionic Compounds
• Collisions between non-metal atoms and metal
atoms results in electron transfer, forming
electrically charged ions. The attractions of
oppositely charged ions are called ionic bonds.
Eg. KCl
162
Ionic Compounds
• These bonds form a 3-D crystal lattice
arrangement of ions and the resulting compound
is called an ionic compound.
163
Ionic Compounds
Ionic compounds are; solids at room temperature,
dissolve in water to some extent, and conduct
electricity when in solution.
164
Types of Ions:
• Monoatomic (simple) ions
- Formed from single atoms.
Eg. Ag+, Al3+, etc.
• Polyatomic (complex) ions
- Formed from groups of atoms &
behave as a single unit.
Eg. CH3COO-, PO43-, NO3-, SO42166
Ionic Compounds: Important Rule
• IN AN IONIC COMPOUND ....…
TOTAL POSITIVE CHARGE = TOTAL NEGATIVE CHARGE
(Example: net charge = 0)
Eg. Sodium chloride
Na+ + Cl- -> NaCl
Ca 2+ + F- -> CaF2
You need to
balance the
charges
http://www.youtube.com/watch?v=h79HW83aoEw&feature=relmfu
Naming Ionic Compounds
reference p 40 - 46
NOMENCLATURE AND FORMULA WRITING
• IUPAC rules for naming (same in any language)
1. Binary Ionic Compounds
2. Polyatomic Ionic Compounds
http://www.youtube.com/watch?v=szhKhCrjUDI
http://www.youtube.com/watch?v=NUVO_TKi6Vo&feature=relmfu
168
1. Ionic Compounds
Nomenclature
- Name of metal ion (cation) followed by name of
non-metal (anion) ion suffix (ide).
- Use lower case letters.
- Name gives no indication of the number of each
ion present.
(Can be determined by balancing overall charge)
Eg. calcium chloride - (CaCl2)
169
1. Ionic Compounds
• If the metal can exist as ions with differing
charges (eg. Cu), the ion charge is indicated
with a Roman numeral in parentheses after
the metal ions name.
Eg. copper (I) chloride = CuCl
copper (II) chloride = CuCl2
170
1. Ionic Compounds
Formula Writing
- Symbol of metal ion followed by symbol of
non-metal ion.
- Ion ratio is indicated by numerical subscripts (no
subscript means one).
Eg. NaCl, CuCl2, Fe2O3
1:1
1:2
2:3
171
1. Ionic Compounds
Ratios are determined by balancing overall
positive and negative charge
Eg.
potassium oxide = K2O
aluminium sulfide = Al2S3
aluminium oxide
2 Al3+
+ 3O2-  Al2O3
172
Summary of Writing and Naming Ionic
Compounds
1) Write the metal (cation) first using the element
name
2) Write the nonmetal (anion) 2nd using the anion
name and changing the last part to “ide”
3) Use the charges on each ion to determine the
lowest whole # ratio that produces an overall
charge of zero
Multivalent Ionic Compounds
• Ionic compounds where the metal can have
multiple charges
– Example:
Multivalent Compounds
• The charge that is listed first is the most
common charge
– Example:
Cu2+ Cu+
Fe3+ Fe2+
• When there is more than 1 charge listed we
MUST specify which charge was used with
Roman Numerals
Multivalent Compounds
Rules:
1. Name the metal using the element name and
follow with the roman numeral that corresponds
to the charge
2. The negative charged ion will take the anion name
 NO ROMAN Numerals
Your Task
• Practice Problems 1 & 2 p 43-44
Polyatomic Ionic Compounds(Table of
Polyatomics)
• Poly “many”
• Atomic “atoms”
• Groupings of atoms that make their own
special ion
• Unlike a molecule which is neutral, these
complex substances are like ions and carry a
charge
• CANNOT exist alone, they are neutralized by
combining with another ion
Polyatomic Ionic Compounds
Nomenclature
- Name of positive ion followed by name of
negative ion.
Eg. NH4Cl, NH4OH
Parentheses may be required to keep
numerical subscripts apart.
Eg. Cu(NO3)2
181
Polyatomic Ions aka Complex Ions
• Poly  “many”
• Atomic  “atoms”
• Groupings of atoms that make their own
special ion
• Unlike a molecule which is neutral, these
complex substances are like ions and carry an
electric charge
• CAN NOT exist alone, they are neutralized by
combining with another ion
Polyatomic Ions
• Naming Polyatomic Compounds
– The positive ion is named first followed by the
negative ion
– Parentheses may be required to keep numerical
subscripts apart
Polyatomics
–The positive ion is named first followed by
the negative ion
–The one with the lesser number of oxygen
atoms has an “ite” ending while the one
with the greater number of oxygen atoms
has an “ate” ending
• Nitrate vs. Nitrite
Polyatomic Ions
The most common polyatomic ions are:
•
•
•
•
•
•
•
•
•
•
•
•
Ammonium
Carbonate
Chlorate
Dihydrogen phosphate
Hydrogen carbonate
Hydroxide
Nitrate
Nitrite
Permanganate
Phosphate
Sulfate
Sulfite
NH4+
CO32ClO3H2PO4HCO3OHNO3NO2MnO4PO4-3
SO4-2
SO3-2
And the best part about these is that you will need to
memorize them all
2. Polyatomic Ionic Compounds
Formula Writing
- Positive ion first, negative ion last, using
numerical subscripts to indicate ion ratio
(found by balancing overall charge).
Eg. iron (III) oxalate,
ammonium chlorate,
ruthenium (III) borate
magnesium nitrate
You get to copy these down in
your notes at the bottom of
the page!!!
186
Solubility
• Ionic compounds usually dissolve in water(aq) BUT
some form a precipitate (solid matter)
• Can dissolve (soluble) = (aq) “aqueous”
• Cannot dissolve (non-soluble)= (s) – leave it solid
• We can use the solubility table to identify (aq) and
(s) ionic compounds
Solubility of Ions in Water
• Group 1 metals and NH4+, H+, ClO3-, NO3- and ClO3- are all soluble in
water
• CH3COO- when in compounds is soluble in water except when added
to Ag+
• Br-, Cl- and I- are soluble in water except when added to Ag+, Pb2+, and
Cu+
• SO4-2 is soluble in water unless mixed with Ca2+, Sr2+, Ba2+, Ra2+, Pb2+,
and Ag+
• OH- is soluble in water when combined with Group 1 elements, NH4+,
Sr2+ or Ba2+
• PO4-3, SO3-2 and CO3-2 are soluble in water only when added to Group
1 elements and NH4+
A Few Examples…
• Lead (II) iodide=
• Sodium bromide=
• Potassium sulfate=
• Magnesium hydroxide=
*If you DO NOT find the ions in the table
assume it is a solid
Ionic Compounds Review
• Five properties that ionic compounds have in
common
–Have high melting points
–Are solid at room temperature
–Form crystal lattices
–Will dissolve in water to a certain extent
–When dissolved in water, the solution will
conduct an electric current
Your Task
Practice Problems 3 & 4 p 46
State if the following are soluble in water
•
•
•
•
•
•
•
Sodium chloride
Ag(NO3)
Magnesium sulfate
Calcium sulfate
Be(OH)2
Na3(PO4)
Lithium chlorate
Name or write the formula for the following compounds. As
well as state if they are soluble or insoluble
•
•
•
•
•
•
Pb(SO4)2
Al(CN)3
Chromium (lll) hydroxide
Cu3P
Tin(ll) nitrite
Fe(NO3)3
Molecular Compounds
(aka covalent compounds)
If no metal atoms are available to give electrons to
non-metal atoms, the non-metal atoms share
electrons to fill their outer energy levels.
The bond which forms when electrons are shared is
called a covalent bond.
195
Molecular Compounds
(aka covalent compounds)
Covalent (molecular) compounds exist as separate
molecules, unlike the lattice of ionic compounds
Eg. F2, H2O
Covalent bonds are usually not as strong as ionic
bonds (more easily broken).
196
Molecular Compounds
(aka covalent compounds)
Molecular compounds can be solids, liquids or gases
at room temperature.
Molecular compounds tend to be poor conductors of
electricity
197
Electronegativity
The strength with which an atom holds onto its
outer (valence) electrons affects covalent
bonding.
On the periodic table, electronegativity increases
as you go up and to the right.
Fluorine is the most
electronegative
element.
198
Molecular Elements
Recall that some elements exist as molecular
elements when in pure form
ie. diatomic (HOBrFINCl At) and
polyatomic (P4 & S8)
199
Diatomic Molecular Elements
• ALWAYS have 2 atoms
• ALL ELEMENTS IN GROUP 17 ARE DIATOMICS
• H2(g), N2(g), O2(g), F2(g), Cl2(g), Br2(l), I2(s), At2(s)
Polyatomics (not the type in their own special table)
• Phosphorous and Sulfur
Molecular Compounds
- Two different kinds of non-metal atoms joined by
covalent bonds.
Nomenclature - use lower case
- (numerical prefix) + first name; (numerical prefix) +
second name with (ide) suffix.
(The prefix "mono" is usually omitted)
Eg. N2O5, H2O, H2S
201
Prefixes
• Mono 1
• Di 2
• Tri  3
• Tetra  4
• Penta  5
• Hexa 6
• Hepta  7
• Octa  8
• Nona 9
• Deca
Binary Molecular Compounds
Formula Writing
- Write symbols in same order as in name
- numerical prefix = numerical subscript
(omitted when "one")
eg. silicon tetrahydride
oxygen dibromide
carbon tetrafluoride
204
Molecular Compounds
Memorize Common Molecular Names and Formulas;
ammonia - NH3
glucose - C6H12O6
sucrose - C12H22O11
hydrogen peroxide - H2O2
hydrogen sulfide - H2S
methane - CH4
Propane- C3H8(g)
Methanol- CH3OH(l)
Acetic Acid- CH3COOH(aq)
Ethanol
- C2H5OH(l)
http://www.youtube.com/watch?v=2_W177RKt0U&feature=relmfu
205
Anomalous Properties of Water Reference
p 60
Water is a polar molecule and exhibits some
‘unusual’ properties:
•
•
•
•
adhesion & cohesion (surface tension)
high c
high mp & bp
density
206
Naming Molecules that Contain
Hydrogen
•
•
•
•
•
•
•
•
•
•
•
H2O(l) – water
H2O2(l) – hydrogen peroxide
NH3(g)- ammonia
C12H22O11(s)- sucrose
C6H12O6(s) – glucose
H2S(g)- hydrogen sulfide
CH4(g) – methane
C3H8(g) – propane
CH3OH(l)- methanol
CH3COOH(aq) – acetic acid vinegar
C2H5OH(l) - ethanol
You need to
memorize
these !
Your Task
Practise Problem 5 p 49
Check and Reflect p 50 # 1 -12
Supplemental p 77 # 28, 36 Check Answers p 495
Do Check and Reflect
p 61 # 6 & 8
Acids and Bases
Arrhenius
• Arrhenius proposed that any
substance that:
– Produces hydrogen ions when dissolved
in water = acid
– Produces hydroxide ions when dissolved in water
= base
• We now refer to these as Arrhenius Acids and
Bases
Acids
•
•
•
•
•
•
•
•
(aq), (s), (l), or (g) at room temp
Soluble in water (to some degree)
Solutions conduct electricity
React with active metals to produce H2 (g).
Taste sour
Turns blue litmus paper red
pH of <7
Identified by the presence of a H+
http://www.youtube.com/watch?v=70fot8t9zts&feature=relmfu
Acids
• Strong Acids:
– Hydrochloric acid(HCl)
– Nitric acid (HNO3 )
– Sulfuric acid (H2SO4 )
– Hydrobromic acid (HBr)
– Hydroiodic acid (HI)
– Perchloric acid (HClO4 )
It is in your best interest
to become very familiar
with these acids. This will
save you significant
amounts of time,
especially on exams.
• Strong acids will completely ionize 
conductive
• Weak acids will not ionize  low conductivity
Bases
•
•
•
•
•
•
•
•
Usually solids
Usually soluble in water.
Do not react with active metals.
Solutions conduct electricity.
Feel slippery, taste bitter.
pH >7
Red litmus paper turns blue
Identified by the presence of a OH213
Summary
Property
Acid
Base
Taste
Sour
Bitter
Feel
Wet
Slippery
Reaction with Mg
Produce Hydrogen gas
No reaction
Indicator
Turns blue litmus red
Turns red litmus blue
Conductive
Yes
Yes
pH
<7
>7
pH scale
• Describes whether a substance is acidic,
neutral or basic (alkaline)
• Refers to the potency or strength of
hydrogen ions present in the substance
215
pH Scale
• Each number change represents a ten-fold
change in the hydrogen ion concentration.
– For example, a pH of 1 is 10 times stronger than a
pH of 2
Indicators
• Chemicals used to determine whether a solution
is acidic or basic
• Change color at different pH
Example: litmus paper
- bases -> turn red
litmus blue
- acids -> turn blue
litmus paper red
217
Indicators
Reference p 63
• Some other
indicators 
• The Universal indicator
is a mixture of several
liquid indicators that
turn different colors
when exposed to
hydrogen ions
218
Buffers
• Any substance that keeps the pH of a solution
nearly neutral even is a small amount of acid or
base is added
• Example: sodium hydrogencarbonate (NaHCO3) – is
in the bloodstream. It neutralizes both acids and
bases that enter the blood.
Acid Nomenclature
Acid Nomenclature - name as an ionic compound
and then switch name according to acid naming
table (next slide) or....
"aqueous" before the ionic name.
No one really uses this nomenclature anymore. So you do not
need to worry about it for the purposes of this course
An acid formula should always be followed by the
subscript (aq), meaning dissolved in water.
220
Acid Nomenclature
• Hydrogen ion acts like a
positive charge
– Will be paired up with a
negative charge
Classical Naming System for Acids
1. If a substance ends with “ide” the
acid becomes “hydro
ic acid”
- Hydrogen ide becomes
hydro
Example: See notes
ic acid
Classical Naming
2. If the substance ends with “ate” the acid
becomes “
ic acid”
- Hydrogen ate becomes
ic acid
Examples: See notes
Classical Naming
3. If the substance ends with “ite” the
acid becomes “
ous acid”
-Hydrogen
ite becomes
ous acid
Examples: See notes
Acid Naming Table
• hydrogen _ide becomes hydro_ic acid
• hydrogen _ate becomes _ic acid
• hydrogen _ite becomes _ous acid
http://www.youtube.com/watch?v=70fot8t9zts
226
Or….
• H+
ide  hydro
Hydrochloric acid 
ic acid
• H+
Sulfurous acid
ite 
ous acid
• H+
Nitric acid
ate
ic acid
Exceptions
• Hydrogen ions can hide at the end of the formula
• Organic acids  Formulas begin with carbon and
end with H ions (hard to tell apart from bases)
– Ethanoic acid (CH3COOH)  acetic acid
•
•
•
•
•
•
•
•
•
•
Acid Naming
H2S
HClO3
H2SO4
HNO3
H2SO3
HNO2
HClO
Phosphoric acid
Hydroiodic acid
Hypochlorous acid
Let’s name
and write
the
formulas for
these
wonderful
acids!!!
Nomenclature of Bases
• For this course we will consider bases to
contain the hydroxide ion (OH)
• Name as an ionic compound
[Note: -COOH is an acid]
230
Naming Bases
• The name of a base always ends in “hydroxide”
• Some common Bases
– Sodium hydroxide
– Ammonium hydroxide
– Calcium hydroxide
– Magnesium hydroxide
Neutralization Reaction
Reference p 68
When an acid and a base are mixed they
neutralize each other by producing water
and a salt. (more later but be aware that salt doesn’t always
mean NaCl)
ACID
H+
+
BASE
OH-
=
HOH
Buffer - has the ability to neutralize either an
acid or base solution
232
Task as a Class
Check & Reflect p 69 # 3,4 & 7
Acid and Bases Naming Worksheets (included in
your notes)
Chemical Change
• Is a process that
involves recombining
chemical bonds
between atoms and
energy flow.
• New substances are
formed.
234
Collision Reaction Theory
• For a chemical reaction to take place:
– Particles of the reactants
before a rearrangement can
occur
– A minimum energy is
particles
– A certain orientation is
particles for successful
rearrangement
must collide
required of the
required of the
235
Reactions
Reactants - the substances which go into a
chemical change.
Products - the substances which come out of
a chemical change.
Reactants
Products
C6H12O6(s) + 6 O2(g) --> 6 CO2(g) + 6 H2O(g)
236
Rules for Chemical Equations
• Use proper chemical symbols
• Reactants on left, arrow pointing to
products on right
• Show states of matter (s, l, g, or aq)
[Note: p 80]
• Are balanced (matter is conserved)
• include energy when possible
More to come
237
Energy Changes
Reference p 81 - 82
• In the formation of ionic
compounds a considerable
amount of energy (heat) may be
absorbed (endothermic) or
released (exothermic).
• Spontaneous reaction - occurs
on its own - when reactants are
in contact.
238
Exothermic Reaction Reference p 81
LINK & LINK TOO
• The energy not used to fuel the reaction (excess
energy) is released to the surroundings. (energy
written as product)
Eg. combustion of methane
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g) + E
• energy stored in newly formed bonds is less
than the energy available from the breaking of
bonds in the reactants.
Endothermic Reaction Reference p 81 link link too
• Reactions which require energy from outside
the system to drive the reaction. (energy written as
reactant)
Eg. Cold pack
NH4Cl(s) + H2O(l) + E  NH4+(aq) + Cl2 (aq) + H2O(l)
• Energy required to form the new bonds is
greater than the energy stored in the bonds of
the reactants.
241
Biochemical Reactions Reference p 82
• Essential to all life on Earth
Respiration
C6H12O6(s) + 6 O2(g) --> 6 CO2(g) + 6 H2O(g) + E
Photosynthesis
6 CO2(g) + 6 H2O(g) + E --> C6H12O6(s) + 6 O2(g)
[also chemosynthesis]
242
Characteristics of Chemical Reactions
Reference p 84
• Colour or density
change
• Gas produced (bubbles,
new odour)
• Energy change (heat,
light, electricity)
• A solid (precipitate)
forms or dissolves
243
Conservation of Mass Reference p 84
• Antoine Lavoisier
showed that the total
number of each kind
of atom remains the
same in a reaction.
That lady who did a lot of work
but because of her gender
doesn’t get any credit
Father of Modern Chemistry
Your Task
• Check and Reflect
p 85 #1-7,10,11
Recall…
• Reactants- the substances which go into a
chemical change
• Products- the substances which come out of a
chemical change
Recall…
• Word equations: use the “ chemical formula”
• Formula Equations: use the “ chemical
formulas” of reactants and products to
represent a reaction
Recall
• Skeleton Equations: a formula equation that
shows the identities of the substances
involved in the reaction and which elements
are present.
Recall…
• Balanced Equations: integers called
“Coefficients” are used to show equal number
of each element present on both sides
Conservation of Mass
• The amount of the reactants equals the
amount of the products
Law of Conservation of Mass &
Balanced Equations
• Chemical equations must
be balanced this is
important when writing
chemical equations
251
Chemical Equation Basics 101
• Coefficient- the number in
front of a symbol; represents the number of atoms
or molecules
- 3Na
Coefficient
• Subscript- the number behind the symbol;
represents the amount of each atom
- N2O5
Subscript
Chemical Equations Basics 102
• Recall that the following elements NEVER exist
as single atoms. They are represented as
shown:
N2
H2
O2
F2
Cl2
Br2
I2
At2
S8
P4
* Sometimes H2O is written as HOH
Chemical Equations
Reference p 86 - 88
• Adding coefficients to a skeleton equation,
balances the equation to show the same
number of atoms, of each type, on either side
of the equation.
• Coefficients in equations may be read as # of
molecules,
or # of moles.
254
Hints for Balancing Equations (NIT)
1. Make sure all reactants and products are written
with correct formulas.
2. Balance each atom, one at a time, using numerical
coefficients (leave O and H to the end).
2a. Balance complex ions as groups
3. Check if the coefficients can be reduced by a
common divisor.
4. Sometimes writing water as HOH can make the
solution easier to see.
Remember
• When counting elements don’t forget to look at the
subscript and coefficient
– P2O5 (has 2 phosphorous and 5 oxygen atoms)
– 2P2O5 (has 4 phosphorous and 10 oxygen atoms)
• Because there are 2 molecules (indicated by the
coefficient) and 2 atoms in each molecule (indicated
by the subscript)
• NEVER CHANGE A SUBSCRIPT TO BALANCE AN
EQUATION!!!!!! USE THE COEFFIECIENT INSTEAD!
Let’s Do An Example Together
• The decomposition of dinitrogen pentaoxide
gas
Step 1: Represent the reactant and products
with formulas
N2O5(s)  N2(g) + O2(g)
Notice that I included the states
• Step 2: When there is an odd number of
atoms of an element on one side and an even
number on the other side, find a common
multiple.
• For oxygen the common multiple of 5 and 2 is 10
This requires you to understand
what a multiple is
• Step 3: Multiply each formula by a coefficient
to represent the same number of atoms of
oxygen on both side
2N2O5(s)  N2(g) +5 O2(g)
• Step 4: Add coefficients to the remaining
formulas to balance the remaining atoms
2N2O5(s)  2N2(g) +5 O2(g)
• Step 5: Check to make sure all atoms are
balanced
Left side
Right Side
4 N and 10 O
4 N and 10 O
Balance the following example
equations:
A. ___ Al(s) + ___ Br2(l) --> ___ AlBr3(s)
B. _ C6H12O6(s)+ _ O2(g) --> _ CO2(g)+ _ H2O(g)
C. Solid iron reacts with aqueous copper (II) chloride to
form aqueous iron (II) chloride and solid copper
261
Further Assistance on Balancing
Equations
• http://www.youtube.com/watch?v=dQrV8Rdu
ttU
• http://www.youtube.com/watch?v=le5zr1kLE
4U
• http://www.youtube.com/watch?v=3UeD32Q
sKYM&feature=relmfu
Your Task
Example Problem p 89
Practice Problem p 89 # 1 Check & Reflect p 90
# 3,6 – 9
[W/S Balancing Chemical Equations – Supplemental]
Five Reaction Types Reference p 91 - 106
1. Formation
(synthesis)
2. Decomposition
3. Hydrocarbon
Combustion
4. Single Replacement
5. Double
Replacement
264
1. Formation (synthesis)
Reference p 91-93
• Composition reactions
• Two simple elements combine to form one complex
compounds.
• element + element --> compound
– A + B  AB
– Reactant + reactant = product
2 K(s) + I2(s) --> 2 KI(s) + thermal E
Example Problems p 92 & 93
Let’s Do This Together…
• Example Problems p 92 & 93
2. Decomposition
Reference p 94
• Complex compounds break down to form
elements
• Compound --> element + element
You have to add energy to
• ab  a + b
Eg.
the system in order for
water to separate
2 H2O(l) + electrical E --> 2 H2(g) + O2(g)
http://www.youtube.com/watch?v=FyUvZ3ldtUU&feature=fvsr
3. Single Replacement
Reference p 96 - 97
• Element + aqueous ionic compound --> new
Element + new ionic compound
• a + bc  b +ac
Cu(s) + 2 AgNO3(aq) -> Cu(NO3)2(aq) + 2 Ag(s)
A + BC
AC
+ B
http://www.youtube.com/watch?v=WwH8I_K3yYM&feature=fvsr
4. Double Replacement
Reference p 100 - 101
• Ionic compound + ionic compound new ionic
compound + new ionic compound
• ab + cd  ad + bc
Example
Pb(NO3)2(aq) + 2 NaI(aq) -> 2 NaNO3(aq) + PbI2(s)
Example Problem p 100
Practice Problem p 100 # 9
http://www.youtube.com/watch?v=7TtOUf91VSU&feature=fvsr
5. Hydrocarbon Combustion
(oxidation reactions)
Reference p 95
• A compound reacts with oxygen to form the most
common oxide
• If O2 is not limited -> produces H2O(g), CO2, and
thermal/light energy
Eg. burning propane
C3H8(g) + 5 O2(g) -> 3 CO2(g) + 4 H2O(g) + E
http://www.youtube.com/watch?v=CY9ldQjdgT8
Oxygen Reactions
Oxygen is always present in these reactions!
1. Combustion Reactions
2. Corrosion Reactions: metal and oxygen
always the reactants
3. Cellular Respiration: Glucose always one of
the reactants
Predicting Products
Reference p 102 - 105
• Using knowledge gained regarding
reaction types it is now possible to
predict products of reactions.
Your Task
Example Problems p 100 - 105
Practice Problem p 103 #10 - 12
Check & Reflect p 106 # 1 – 8
[W/S Translating Formulae – Supplemental]
What is a Mole?
• We use “convenient numbers” in everyday
situations
• A “convenient number” is a simplified way to
describe the number of things
• It is often easier to group objects when
counting them
So What is a Mole?
• A mole is a counting unit. It tells us how many
objects there are.
• A pair and a dozen are also counting units
Pair
means
2 objects
Dozen
means
12 objects
Mole
means
6.02x1023
objects
So what is a Mole?
• A term used to represent the number of
objects
• The amount of any substance containing
6.02x1023 particles
– To define a mole, chemists measure carbon
– And we also call 6.02x1023 Avogadro's Number
Where does this Number come from?
• Scientists took 12g of pure carbon-12
(composed of 6 protons+ 6 neutrons; mass
number =12) and counted how many carbon
atoms were in the sample
• Guess how many there were?
Avogadro’s Number -The Mole Reference
p 107
• # of atoms in 12.0000 g of 12C is called
a mole (mol) = 6.02 x 1023
Avogadro’s #
• this number of atoms is observable and
measurable.
• coefficients in equations may be read
as # of molecules, or # of moles.
http://www.youtube.com/watch?v=xiVweBpjXJo&feature=related
http://www.youtube.com/watch?v=xqw2BWdKl1Q&feature=relmfu
http://www.youtube.com/watch?v=O7qjYRYxkso&feature=relmfu
283
Mole Ratio
• The ratio of the coefficients gives us the mole
ratio.
2 Al(s) + 3 Br2(l)  2 AlBr3(s)
• In the balanced chemical equation above..
- if 3 mol of Br2 reacts, 2 mol of AlBr3 forms
- if 6 mol of Br2 reacts, 4 mol of AlBr3 forms... etc.
284
Molar Mass
Reference p 108 - 109
• We can now calculate the mass of one mole
of a compound (molar mass ).
• Use the atomic molar mass (M), in g/mol
multiplied by the number of moles in the
chemical formula then add the masses of all
elements together
Example Problem p 108
Eg. Find the molar mass of H2SO4(aq)
hydrogen atomic mass = 1.01 g/mol
# of moles in formula = 2
2 mol x 1.01 g/mol = 2.02 g
sulfur atomic mass = 32.06 g/mol
# of moles in formula = 1
1 mol x 32.06 g/mol = 32.06 g
oxygen atomic mass = 16.00 g/mol
# of moles in formula = 4
4 mol x 16.00 g/mol = 64.00 g
...molar mass of H2SO4 =
2.02 g + 32.06 g + 64.00 g = 98.08 g/mol
286
Your Task
• PP p 108 # 13 – 16
Mass <--> Mole Conversions
Chemists created the concept of molar mass
to convert between mass and chemical
amount
These conversions can be performed using a
formula (M=m/n) or the factor-label (unit
analysis) method.
288
n= m/M
Where:
n= number of moles
m= mass
M= molar mass
However, you are much too
advance to use this formula
and are capable of using unit
analysis
Molar mass is often used as a conversion factor
because it contains the units grams and moles. Thus,
it can be used to convert grams to moles and moles
to grams.
Example
• How many moles are there in 20 g of carbon?
• Find the number of moles of H2SO4 in a 100.0g
sample.
Mole to Mass Conversion
• If we had a 4.60 mol amount of H2SO4, its
mass would be........
m = (4.60 mol)(98.08 g/mol)
m = 451 g
291
Mass to Mole Conversion
• If we had a 0.58 g of H2SO4, its mole
amount would be........
n = (0.58 g)/(98.08 g/mol)
n = 56.8864 mol
n = 5.9 x 10 -3 mol
Example Problems p 109
Practice Problems p 108 - 109 # 13 - 20
Check & Reflect p 112 # 1 -12
292
Your Task
Example Problems p 109
Practice Problems p 108 - 109 # 13 - 20
Check & Reflect p 112 # 1 -12
Nuclear Reactions (NIT)
Chemical reactions involve the valence electrons
and the bonds which they form (formation of new
molecules).
Nuclear reactions changes in
the atom's nucleus, the
formation of a different atom
and thus a new element.
294
Nuclear Reactions (NIT)
In nuclear reactions mass
changes into electromagnetic
radiation (E = mc2).
Nuclear reactions involve up to
106 times more energy than
chemical reactions. The high
energy radiation given off is
usually X-rays or gamma rays.
295
Nuclear Reactions (NIT)
Radioactivity is the emission of particles and
radiation from an atom's nucleus.
Some isotopes of an atom's nucleus have an
unstable ratio of protons to neutrons which makes
them radioactive (radioisotopes).
These atoms undergo radioactive decay to become
more stable isotopes.
Eg. Carbon- 14
14C

14N
+ radiation
296
Unit REVIEW
• Use Section Reviews
• p 27
• p 76 - 77
• p 113
and Unit Summary/Review
• p 116 - 121
to help you prepare for the Unit Exam
(selected answers on pg 498)
But not really because that
would be mean and
inappropriate
297
1.
2.
3.
4.
5.
6.
7.
8.
9.
Quiz
Acetic Acid
Hydrochloric acid
Nitrous acid
H2CO3
H2SO4
Magnesium hydroxide
Chloric Acid
Calcium hydroxide
What is the molar mass of an unknown oil if
57.29g of the oil contains 0.15mol?
10.How many grams are in 5.2 moles of potassium
sulfide?