Science Focus 10
Unit 1 Energy and Matter in Chemical Change pp 2-137
In this unit we will explore:



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Atoms, elements and compounds
Ionic, molecular, acidic and basic compounds
Naming chemical compounds
The mole
Water
Chemical equations and reactions
Chapter 1: Atoms, Elements and
Compounds pp 2-39
1.1 Working with Chemicals p 6
Aboriginal peoples have been using chemical substances
for thousands of years to preserve food, treat illness,
build tools and decorate clothing. Many of these
traditional processes are still used today.
Many chemicals, even household products, have dangerous
properties and must be handled properly. An MSDS, or
Material Safety Data Sheet, lists important information
including physical properties (ie melting and boiling
points), chemical dangers, and how to store and dispose
of the chemical. See figure 1.3 p7.
The WHMIS or Workplace Hazardous Materials
Information Systems informs users of the dangers of
working with specific chemicals by informative labels on
the chemical container, MSDS paperwork and
worker training programs. See figure 1.4 p8
for examples of WHMIS symbols.
Do BLM 1-1 (MSDS).
Classifying Matter
Matter is anything that has mass and occupies space.
 Mixtures can be mechanical (heterogeneous) where
the separate parts are visible or solutions
(homogeneous) where the different parts are not
visible.
 Pure substances include elements (such as Na or Cl)
or compounds (such as NaCl). Compounds can be
chemically separated into simpler substances,
elements cannot.
 See figure 1.5 p10.
Do Practice Problems #1-4 p10
Do Check Your Understanding #1-3 p11
1.2
Developing Atomic Theories p12
The idea that matter is made up of
small particles is rooted in work of
scientists many centuries ago. One of
the earliest theories is credited to the theorist
Democritus (~400 BC), who stated that matter is made
up of infinite tiny, indivisible, constantly moving units.
This theory has evolved over recent history….
Historical Models of the Atom
Model of atom
Why model was
proposed
● to account for the
behaviour of
substances when
they interact with
each other
according to the
law of conservation
of mass (when
substances react,
the total mass of
substance before
and after are the
same)
Key features
● to account for the
existence of small
negative charges
(cathode ray
particles)
associated with the
atom
 gas discharge tube
-see figure 1.9 p13
 Thomson’s
experiment – see
figure 1.10 p15
● All materials produce
identical cathode ray
particles, which are much
smaller than even the
smallest atom, hydrogen.
● All atoms contain electrons
(small negative charges)
that are distributed
throughout a positively
charged solid.
 Electrons are extremely
tiny compared with the size
of an atom.
The Dalton Model
“Billiard Ball”
Dalton: 1766-1844
See figure 1.8 p13
The Thomson model
“Raison Bun”
Thomson: 1856-1940
See figure 1.11 p16
● All matter is composed of
small indivisible particles,
called atoms, that can be
neither created nor
destroyed.
● All atoms of the same
element are identical in
mass and size, but
different in mass and size
from atoms of other
elements.
● Atoms exist in an otherwise
empty space and are in
constant motion, during
which they may collide to
form new combinations
(compounds).
● Chemical reactions change
the way in which atoms
are grouped, but they do
not change the atoms
themselves.
Why model was
rejected or modified
● contained
inaccuracies
regarding the relative
masses of several
atoms
● could not account for
the electric nature of
matter
● could not account for
the production of
high-frequency
electromagnetic
radiation
● could not explain the
existence of alpha
particles (heavy,
positive charge) for
radioactive materials
Model of atom
The Rutherford model
“Solar System”
Rutherford: 1871-1937
See figure 1.14 p18
The Bohr Model
Bohr: 1885-1962
See figure 1.19 p20
Why model was
proposed
● to account for the
deflection and
reflection of alpha
particles from gold
foil
 Rutherford’s
experiment – see
figure 1.12 p17
 Explanation of the
experiment – see
figure 1.13 p17
 Isotopes – see
figure 1.15 p19
● to account for the
production of
bright-line spectra
in hot gases and
absorption spectra
in cold gases
 See figures
1.17,
1.18 p20
Key features
● At the centre of every
atom, there is a small,
positively charged nucleus.
● The nucleus accounts for
the majority of the mass
of the atom.
● Electrons are attracted to
the nucleus and orbit in a
cloud around the nucleus.
 A third subatomic
particle, neutron, with no
charge but a mass similar
to a proton exists in the
nucleus
● Electrons reside within
stable energy levels
(electron shells) at fixed
radii relative to the nucleus
of the atom. Transitions of
electrons to higher energy
levels require energy.
Transitions to lower energy
levels produce
electromagnetic radiation
(light and radio waves).
● Each energy level has a
fixed maximum number of
electrons that can reside
in it.
Why model was
rejected or modified
● could not account for
the emission spectra
of elements
● included orbiting
electrons that should
constantly radiate
energy and therefore
collapse into the
nucleus, but the
radiation and collapse
were not observed
● could not explain the
existence of
sublevels within
higher energy levels,
predicted by the
shape of the periodic
table
● could not fully explain
the nature of
electrons or the
nucleus
QUARKS
Scientists now believe that neutrons and protons are made up
of even smaller particles called quarks. They believe that
matter is made up of dozens more of these subatomic particles.
Features of a Simplified Modern Model of the Atom
(figure 1.20 p22)
 a tiny, dense nucleus that is surrounded by electrons (e-)
 nucleus contains protons (p+) and neutrons (no), called nucleons
(exception: H-1 nucleus contains one proton only)
 nucleus accounts for most of the mass of the atom
 e- exist at certain allowed energy levels
 p+ carry a positive charge, e- carry a negative charge and no
carry no charge. See table 1.2 p 22.
 A neutral atom always has equal numbers of e- and p+
If a nucleus were the size of a baseball, the electron orbitals
around the nucleus would take up the space of an arena!
Nuclear Notation
Elements can have 2 or more isotopes, where the number
of neutrons is different. Scientists use the following
notations to describe a specific isotope:
29
14
Si
Silicon-29
or
 top number represents the mass
= total number of protons + neutrons
 bottom represents the atomic number
= total number of protons in the nucleus
These two numbers can be used to determine the number
of neutrons:
mass number – atomic number = # neutrons
so with Silicon-29, the number of neutrons is 29-14=15!
Do Practice Problems #5-8 p23 + #4 p24
Do Check Your Understanding #2, 3, 5 p24
1.3
Electrons and the Formation of Compounds p25
During the 1800’s, a Russian chemist, Dmetri Mendeleev,
examined 62 elements. He developed a table of these elements
based upon the fact that they had repeating properties. The
table was called a “periodic table” for this reason. He further
predicted the existence and properties of unknown elements and
left spaces on his periodic table for them.
Examine our modern periodic table on p482. It displays
the known elements in a format that follows various
patterns and trends:
 Elements are arranged according to increasing
atomic number (number of protons in the nucleus)
 Period – horizontal row
 Inner-transitional elements
top period = lanthanoids (fits in period 6)
bottom period = actinoids (fits in period 7)
 Group or family – vertical column, each element in a
group has similar properties
 Group 1 = alkali metals (react with water)
 Group 2 = alkali earth metals (react with
oxygen)
 Group 17 = halogens (very reactive)
 Group 18 = noble or inert gases (not reactive)
 Staircase line –bordered by metalloids, this line
divides metals on the left from non-metals found on
the right. See table 1.3 p25 to examine the
properties of metals, non-metals and metalloids.
Patterns of electron arrangement in periods & groups
Recall that Neils Bohr inferred that electrons orbit the
nucleus of the atom in fixed energy levels, and each level
can only hold a certain maximum number of electrons.
The first can hold 2 electrons, the second can hold 8, and
the third can hold 8. We can use the patterns in periodic
table to predict the number of energy levels of an
elements and the number of electrons in each energy
level. Examine figure 1.22B p26.
 period number shows the number of energy levels
used in each element (eg. period 2 elements have 2
energy levels)
 group number describes how many electrons are
found in the valence or outermost energy level (eg.
lithium is in group 1 and has 1 valence electron)
 for groups 13 – 18, we use the last number to
designate the number of valence electrons (eg.
elements in group 16 have 6 valence electrons)
 electrons fill the first orbital before they can
occupy the second, and fill the second before they
can occupy the third
 when the valence level is full, it is referred to as a
stable octet since there are 8 electrons occupying
the orbital (unless it is the first level)
The diagram representing the element beryllium looks
like this: __2e-__
Try the diagram for fluorine:
__2e-__
4p+
Do Practice Problems #9-12 p27 and BLM 1-4.
Electron Dot Diagrams, a.k.a. Lewis Dot Diagrams
American chemist G.N. Lewis invented these structures
to visualize and track electrons during bond formation.
To draw:
1) write its chemical symbol
2) surround it by the number of dots that represent the
atom’s valence electrons
3) if an atom has more than 4 valence electrons, the
additional electrons are shown paired with others
4) note that elements in the same group will have identical
dot diagrams since they have the same number of
valence eDo BLM 1-6 and then BLM 1-5.
Formation of Ions
Any atom or group of atoms that either loses or gains eand thus carries either a positive or negative charge is
called an ion.
 Cation - positive charge, has fewer e- than p+
- metallic atoms lose electrons, eg. Na+
- remember: cats have paws (pos)
 Anion - negative charge, has more e- than p+
- non-metallic atoms sometimes gain electrons
from other atoms, eg. Cl- but they can also
share electrons
See p29 figure 1.24 and 1.25 for the 3 ways to represent
the formation of ions.
Formation of Ionic Compounds
Ions do not form by themselves. Instead, as metallic and
non-metallic atoms collide with one another, their valence
electrons interact. The metal will lose its valence
electrons (the cation) and an adjacent non-metal will gain
those electrons (becoming the anion); this is a transfer
of electrons. The two ions formed are opposite in charge
and are greatly attracted to each other, forming a very
strong ionic bond. The rearrangement of electrons allows
each ion a full valence orbital (like its nearest noble gas)
and therefore leads to a greater stability and lower
energy level for the ionic compound.
Important notes about ionic compounds:
 ions arrange themselves in a regular repeating pattern
called a crystal lattice (see table salt + figure 1.28,
p31)
 a binary ionic compound is formed from only 2
elements.
 Why does Na become Na1+ and Mg become Mg2+?
Sodium (Na) has only 1 valence electron, so when it is
lost, the resulting charge is 1+, however magnesium
(Mg) has 2 valence electrons so its ion has a charge of
2+
Do BLM 1-7 to practice isotope and ion formation
Formation of Molecular Compounds
When non-metals react with one another, electrons are
NOT transferred since both atoms tend to gain electrons
to fill their valence energy level. Instead, the valence
electrons are shared. These compounds are referred to
as molecular and the sharing of electrons forms a bond
called covalent (co = share, valent = valence shell)
Important notes about molecular compounds:
 a diatomic molecule is composed of only 2 atoms and
a polyatomic molecule is composed of many atoms
(see and memorize table 1.5 p 31)
 molecules do NOT form a crystal lattice shape
It is important to note that the physical and chemical
properties of a compound is different from those of
the individual elements that make up the compound.
Do Investigation 1-A p 33 + Questions 1-6 p 35
Do Check Your Understanding #1,3,6,7 p36
Do Chapter 1 Review #2,4,5,7,10,11,12,15,16,17,19,20 p38
Chapter 2: Names, Formulas and Properties
pp 40-81
2.1 Chemical Names and Formulas p 42
The IUPAC (International Union of Pure and Applied Chemistry)
was founded in 1919 and has developed a method to name
chemicals based on their composition. The systematic name of a
compound allows you to write the chemical’s formula and predict
some of its properties.
Binary Molecular Compounds
Naming Rules: see table 2.2 p 44
1st name:
use the name of the element found
furthest to the left on the periodic table
2nd name:
add the suffix “ide” to the name of the
second element (eg. oxygen becomes oxide)
**attach a prefix to each name to indicate how many of each
atom is present in one molecule of the compound (see and
memorize the prefixes in table 2.1 p 43)
**note that the prefix “mono” is not generally used for
the first element
**when hydrogen is the first element in the formula, the
compound is molecular BUT prefixes are NOT used (see p54 –
Hydrogen Compounds). More later.
**memorize: diatomic elements (N2, O2, H2, F2, Cl2, Br2, I2, At2,) + P4 + S8
**memorize:
O3 – ozone
NH3 – ammonia
CH4 – methane
C6H12O6 – glucose
CH3OH – methanol
C12H22O11 - sucrose
Do Practice Problems p44
C2H5OH – ethanol
H2O2 – hydrogen peroxide
Binary Ionic Compounds
Naming Rules:
1st name:
2nd name:
use the name of the metal element
add the suffix “ide” to the name of
the non-metal element
**the chemical formula is the simplest whole ratio of
each type of ion in the compound and it represents the
smallest repeating unit (formula unit) within a crystal
lattice (fig 1.28 p31)
**prefixes are NOT used in naming ionic compounds
Study Model Problem 1 and figure 2.3 p45
Do Practice Problems p45
Stock System:
As you look at your periodic table, you may notice that
many transitional elements have more than one possible
ionic charge.
Study Model Problem 2 and do Practice Problems #9,10 p46
A German chemist, Alfred Stock, developed a way to
indicate which cation is in a compound. The Stock
System uses a roman numeral to indicate the charge on
the metal or cation.
Study Model Problem 3 and do Practice Problems #11,12 p47
Do BLM 2-1 (and BLM 2-5 to practice the crossover method)
Polyatomic Ions
These ions consist of two or more different atoms joined
together by covalent bonds. As a group, the
bonded atoms have either a positive or negative
charge. See the polyatomic ion table on the back
of the periodic table, and become familiar with
the common polyatomic ions in table 2.3 p51.
Naming Rules:
1st name:
name the cation
2nd name:
name the anion
**parentheses must be placed around the polyatomic ion
if more than one are found in a chemical formula
Study Model Problems 4,5 p51 and do Practice Problems p52
Getting to know the polyatomic ions:
Many polyatomic anions contain oxygen and those that
contain the same elements are part of a “family”.
Anion name
Per----ate
----ate
----ite
Hypo----ite
# of oxygen
1 more
--------------1 less
2 less
Do Practice Problems 17-19 p53
Example
pernitrate
nitrate
nitrite
hyponitrite
Formula
NO4NO3NO2NO1-
Do BLM 2-3 (and Check Your Understanding p55 for practice)
2.2 Explaining Properties of Substances p56
Physical properties – observable and measurable qualities
of a substance, such as melting and boiling points,
conductivity, appearance, state (see table 2.7 p56)
Chemical properties – reactive properties of a substance
Property
Ionic compound
State at room temp.
Melting point
Solid (hard, brittle,
Solid, liquid or gas
crystal lattice shape)
High
Low
Attraction between
molecules
Strong because of
opposing charges
Weak
Conductivity when
solid?
No, molecules held
rigidly in crystal
lattice shape
Yes, crystal lattice
has broken down
= electrolyte
No
Conductivity when
dissolved in water or
melted?
Lab safety demo
Read together and sign laboratory contract
Review lab write-up guidelines
Molecular
compound
Sometimes, but
these compounds
tend to be
non-electrolytes
Do Investigation 2-B p58 (use BLM 2-2) and Q #1-6 p59
Toxic Properties
There are many chemicals found in society
today with properties that produce
pronounced physical and psychological
effects. Some of these chemicals are
regulated or even banned by government and
others are not. Understanding chemistry
can help people make informed decisions
about personal use of such chemicals.
Examine table 2.8 p 60 to familiarize
yourself with effects of the chemicals
ethanol and nicotine.
Do Check Your Understanding #1,3,5 p62
2.3 Properties of Acids and Bases p 63
Both acids and bases have several properties in common.
These chemicals can be both dangerous and useful.
Acidic and basic solutions are electrolytes, so they
contain freely moving ions and thus conduct electricity.
See figure 2.7 p63 for examples of household acids and
bases.
Property
Arrhenius
definition
(a Swedish
chemist)
Formula
Electrolyte?
Taste
Feel
Uses
ACID
 releases H+
(hydrogen ions ) in
a water solution
HB(aq) ->H+(aq) + B-(aq)
Yes
Sour (like lemons)
Will burn skin
 removes rust
 refine metal
Less than 7
pH
Indicators:
Litmus
Red
Phenolphthalein Colorless
BASE
 releases OH(hydroxide ions) in a
water solution
MOH(aq) ->M+(aq) + OH-(aq)
Yes
Bitter (like tonic water)
Will burn skin, but bases
feel slippery (soap)
 cleaning products
 medical drugs
More than 7
Blue
Pink
For important acids and bases and their uses see tables 2.9 and 2.10 p64.
Notes:
 pH indicators are made of chemicals that changes color in
response to the concentration of hydrogen or hydroxide
ions.
 pH meters and probes can also be used. These instruments
measure the electrical property of the solutions, which
relates to the solution’s pH. They are more precise but
more expensive.
 When acids and bases react, they neutralize each other, or
they lose their acid and base properties. The resulting
solution is a salt dissolved into water.
 Strong acids react with metals to produce hydrogen gas.
 Strong acid = high conductivity, weak acid = weak
conductivity
Naming Acids:
Name the compound as though it were an ionic compound.
Use the following rules to complete the acidic name:
1.
2.
3.
If the ionic name was:
hydrogen _______ide
The acidic name is:
hydro_____ic acid
eg. H2S = hydrogen sulfide = hydrosulfuric acid
hydrogen _______ate
______ic acid
eg. HClO3 = hydrogen chlorate = chloric acid
hydrogen _______ ite
______ ous acid
eg. H2SO3 = hydrogen sulfite = sulfurous acid
Notes:
 Acids are aqueous solutions so the proper way to
write the formula of an acid is with an (aq)
subscript. H2S, hydrosulfuric acid, can also be
called aqueous hydrogen sulfide.
 When trying to figure out the formula of an acid
from its name, remember that you must have enough
hydrogen ions to balance the negative charge of the
anion.
 If the polyatomic ion that is part of the acid ends in
a ---COO-, the hydrogen ion bonds at this end to
become ---COOH (this is called a carboxylic acid).
Do Practice Problems p70
Do Find Out Activity p 67 and BLM 2-6
Do Check Your Understanding #3,4,5,6,8 p71
2.4 Why Water is Weird p 72
Water is essential to life. The cells of most living
organisms are 80-95% water, and virtually all
chemical reactions for life occur in aqueous
solutions. But water also has many other unique
properties:
Occurs in all 3 states on Earth
 Ice is less dense than liquid
 Its melting and boiling points are
higher than similar substances (see table
2.14 p72)
 It has a very high surface tension
Do Investigation 2-D p73 (use BLM 2-9 to record results) and
do Q #1-3,6 p74

A Molecular View
The above characteristics of water can be explained by
several features of water.
1.
Shape – bent (see figure 2.13 p75)
2. O-H bonds – each bond is made up of
2 electrons that are more strongly
attracted by the oxygen atom so that
it gains a slightly negative charge as
the hydrogen gains a slightly positive
charge, making it a polar or dipole
molecule.
These properties encourage individual water molecules to
attract each other. The negative oxygen attracts the
positive hydrogen and forms a hydrogen bond. This kind
of bond is an intermolecular bond, one between molecules
(versus intramolecular bonds, within molecules, like
covalent bonds which are stronger).
The hydrogen bond explains many of water’s unique
properties:
 its melting and boiling points are high
-more energy has to be added to water in order to
break these intermolecular bonds and allow ice to
melt or water to vaporize
 it takes much energy just to raise the
temperature of water
-the strong attraction between water molecules
must be overcome to increase the average speed
of the water molecules and thus heat the water
 it has a concave meniscus and shows capillary
action
-its polar molecules are attracted to the sides of
the container
 when it is solid water, it floats in liquid water
-hydrogen bonds force water molecules farther
apart as it freezes and leaves the ice less dense
than liquid water
 it has a high surface tension / cohesion
-hydrogen bonds pull water molecules close
together
Do Check Your Understanding #1,2,4,6 p78
Do Chapter 2 Review #1,2,4,6,8,9,10,13,14,17,18,19,22,25
p80
Chapter 3: Chemical Reactions pp 82137
3.1 Recognizing and Describing
Chemical Reactions p 84
A chemical reaction, or chemical change, has taken place
when one or more substances have changed to form
different substances. The reactants are those
substances that undergo a chemical change and the
products are the new substances formed.
Evidence of a Chemical Reaction: (see table 3.1 p84)
 Odor change
 Color change
 Formation of a gas (bubbles)
 Formation of a solid (called a precipitate)
 Energy change such as:
 Temperature increase or decrease
 Light produced
 Sound produced
 Electricity produced
*energy changes can also accompany physical changes,
which are simply changes in state
Do BLM 3-1
Solubility and Chemical Reactions
Recall that when most ionic compounds are dissolved in
water, the cations and anions separate from each other
and move about freely. Also remember that some ionic
compounds do not dissolve well in water. If
two ionic compounds are mixed, their free
cations and anions may react with each other
to produce new substances. If the new
compound does not dissolve in water, a precipitate is
formed, and this indicated a reaction has occurred.
Do Investigation 3-A p86 (use BLM 3-3 to record)
and do Q#1-6 p87
Predicting Solubility
Table 3.2 p88 outlines which ionic compounds are soluble
in water and which have a low solubility. To determine
whether a compound is soluble in water or not, follow
these steps:
1.
locate one of the ions found in the compound on
the top row of the chart
2. look down that column of the table to find the
second ion
3. if the 2nd ion is found in the “high solubility”
category, the compound is soluble in water and
should be given a subscript of (aq) (aqueous)
OR
If the 2nd ion is found in the “low solubility”
category, the compound is NOT soluble in water
and should be given the subscript (s) (solid)
It will form a precipitate in water.
Study Model Problem 1 p 89
Do Practice Problems 1-4 p90 and BLM 3-2
Chemical Reactions and Energy Changes
Energy changes accompany any chemical reaction.
Endothermic – energy is added to start
the reaction
Eg. Photosynthesis – energy from
the sun is required to start the process
of converting carbon dioxide and water
into food (glucose) and oxygen in plants
Exothermic – release of energy
during the reaction
Eg. Cellular respiration –
the opposite reaction of
photosynthesis, where glucose
and oxygen are converted into
carbon dioxide and useful energy
in both plants and animals
Do BLM 3-4 for reinforcement
Law of Conservation of Energy
This law states that energy can be converted from one
form into another, but the
total energy of the universe remains constant.
Energy cannot be created or destroyed.
Energy is required to break chemical bonds, and energy is
released when chemical bonds are formed.
Endothermic reaction – more energy is required to
break the bonds of the reactants than is released when
the products’ bonds are formed
Exothermic reaction – less energy is required to
break the bonds of the reactant than is released when
product bonds are formed
See figure 3.8 p 92.
Do Check Your Understanding #1,3,4,5 p93
3.2 Representing Chemical Reactions p94
Closed system – an exchange of energy
between the system and its surroundings
occurs, but NOT an exchange of matter
Example – a terrarium
- the earth?
Open system – an exchange of both energy and
matter between a system and its surroundings
Example – a flower pot
Isolated system – exchange of neither energy nor
matter between a system and its surroundings
Example (theoretical) – a thermos
Law of Conservation of Mass
Developed by Antoine Lavoisier (1743-1794)
This law states that during a chemical reaction, the total
mass of the reacting substances (reactants) is always
equal to the total mass of the resulting substances
(products).
Do Investigation 3-B p 95 and Q# 1-10 p96
Writing Balanced Chemical Equations
Reaction Description: hydrogen gas reacts with oxygen
gas to produce liquid water
Word equation:
Skeleton equation:
hydrogen + oxygen -> water
H2(g)
+
O2(g) -> H2O(l)
Balanced equation:
2H2(g) +
O2(g) ->2H2O(l)
*shows atoms are conserved
*coefficients (numbers in front of chemical
formulas) show how many of each compound is there
(note there is 1 molecule of oxygen gas, but 2 atoms of
oxygen)
Q: how many atoms of hydrogen are reacting? ______
*subscripts (letters below and to the right of each
compound) show what state each compound is in
How the balance equations:
1.
When writing the correct skeleton
equation, check your formulas to
ensure they are correct. Include
all states of matter for each
reactant and product.
2. Balance the atom or polyatomic ion
present in the greatest number. Find the lowest
common multiple to obtain whole number
coefficients to balance.
3. Repeat step 2 to balance each remaining atom/ion.
HINTS:
 O2(g) – when O2(g) is present in the
reaction, balance it last
 When single elements are present on one
side of the equations, balance those
elements last (ie. Na, Mg)
 As long as complex ions are found intact
on both sides of the equation, treat the
complex ion as one group
4.
Check once again to make sure the equation is
balanced by creating a table of atoms on each side
of the equation.
Study Model Problem 1-3 p99-100
Do Practice Problems 5-8 p101
Do Check Your Understanding #2-5 p102
Do Worksheets p2 (left side)
3.3 Types of Chemical Reactions p103
Do Find Out Activity p103 using BLM 3-5 to record.
Do the appropriate sections of BLM 3-6 as equations are
learned
1.
Formation Reaction: two or more reactants combine
to form a new product
Eg. X + Y -> XY
See figure 3.15 and 3.16 p104.
2.
Decomposition Reaction: one compound breaks down
into two or more simpler compounds/elements
Eg. XY -> X + Y
See figure 3.17 and 3.18 p 104.
3.
Single-Replacement Reaction: one element takes
the place of another element in a compound
Eg.1.
A + BX -> AX + B
Eg.2.
AX + Y -> AY + X
*many of these reactions are between a metal and a compound
(see p105)
*sometimes the reaction is between a halogen and a halogencontaining compound (see p105)
4.
Double-Replacement Reaction: the cations of two
different compounds exchange places, forming two
new compounds
Eg. WX + YZ -> WZ + YX
*many of these equations result
in a precipitate forming
or the formation of water (p105)
5. Reactions involving Carbon
Compounds:
The study of carbon-containing compounds is called
organic chemistry. A hydrocarbon is a
compound that contains only hydrogen and
carbon atoms (eg. C2H6). These compounds are
retrieved by refining crude oil and natural gas.
Approximately 95% of these hydrocarbons are
burned as fuels in exothermic combustion
reactions to create thermal energy to warm
buildings and provide energy for transportation.
Complete Combustion: a hydrocarbon
reacts with oxygen gas (or burns) to
create carbon dioxide gas, water vapor
and thermal energy. See figure 3.20 and
3.21 p108.
Incomplete Combustion: when oxygen is in poor supply,
the products are carbon dioxide gas, water vapor, carbon
monoxide, carbon (soot) and less thermal energy than
a complete combustion reaction.
*carbon monoxide is a colorless, odorless, highly
toxic gas that, when breathed in, strongly binds to red
blood cells instead of oxygen and can lead to death
Study the Tools of Science p109 and learn the chemical tests
for hydrogen, oxygen, carbon dioxide and water
Do Check Your Understanding #1-9 p114
Do the rest of your Worksheet
3.4 The Mole p116
Since atoms and many compounds are very tiny, scientists group them into an
extremely large number called a mole.
The mole:
 A mole is defined as the amount of substance that contains
as many elementary entities (atoms or
compounds) as exactly 12 g of carbon-12
(the most common isotope).
 One mole of a substance contains 6.022
141 99x1023 particles of that substance
 6.02 x 1023 is called Avogadro’s number
(see Did You Know p116)
The unit mol is short for the German word “Molekulargewicht”,
which is literally translated to, you guessed it, molecular weight!
Atomic Molar Mass (M)
 this is a weighted average of 1 mol of all the
naturally occurring isotopes of that element
 it is listed on the periodic table
 units are in g/mol
Molar Mass of a Compound (M)
 this is the mass of 1 mol of any pure substance
 units are in g/mol
 it is determined by using the formula for that
compound (see eg. p120)
Do Practice Problems p 120
Do WS p.2,4
Review significant digits worksheet
Converting Between Mass and Moles
Scientists often convert between the mass (g) of a
sample and the number of moles (mol). The following
formula is used:
n = _m_
M
n = number of moles of substance
m = mass of substance
M = molar mass of substance
Study Model Problem 1 and 2 p121 and 122
Do Practice Problems p 122 and 123
Do BLM 3-10
The Mole and the Law of Conservation of Mass
Study the chemical reaction at the top of p124. The
coefficients explain the number of molecules and/or the
number of moles of each substance needed for the
reaction to run. Now view table 3.5 at the bottom of the
page. Notice how molecules, moles and mass are all
related in the chemical equation. Also notice that the
mass of the reactants is always equal to the mass of the
products.
Do
Do
Do
Do
BLM 3-10 + rest of Worksheet
Check Your Understanding #2-6 p125
Chapter 3 Review #1-5,7,9-10,12(a-b),13(a-c),14-19
Unit 1 Review #1-38,39-43(a-c for each),46