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Science 10
Introduction
Science 10 Outline
• Resource text:
• Review of Science 10 Chemistry
• Science and Safety
• Classifying matter
• Atomic structure
• Chemical names and formula
• Chemical reactions
Evaluation:
• Course work
Tests & Quizzes, Labs 40%
Assignments and labs 40%
• Final Exam
20%
• Marks may be deducted for work handed
in late. Time factor will be addressed.
Lab Safety
• Common Sense!
• No fooling around
• Wear proper clothing
• Use equipment correctly
• No unauthorized experiments
• Hand chemicals with respect never taste
• Clean up - wash/return/turn off
TERMS
• IUPAC –
International Union of Pure
and Applied Chemistry
•MSDS –
Materials Safety Data
Sheet
WHMIS
Workplace Hazardous Materials Information System
Provides:
1) information
2) safety & training
3) labeling
WHMIS SYMBOLS
• Class A
• Compressed
Gas
• ie) acetylene
cylinder (welding)
• Class B
• Flammable
&
Combustible
Material
• ie) methane
• Class C
• Oxidizing
Material
• ie) Acid
• Class D1
• Materials causing
immediate &
serious toxic
effects
• ie) Acids & Bases
• Class D2
• Materials
causing other
toxic effects.
• ie) Heavy
metals (Pb, Hg)
• Class D3
• Biohazardous
Infectious material
• viruses, biological
weapons
• Class E
• Corrosive
Material
• ie) Acids &
Bases
• Class F
• Dangerously
Reactive
Material
• ie) sodium
Questioning
Defining Problems
Interpretation of Evidence
Ideas
Proposing
-Evolution
Information
-Evaluation of hypothesis and prediction
-Extension - sources of error
- applications
- new ideas
- Background
Processing Evidence
Experiments
-Analysis
Design
- charts
variables
- graphs
calculations
- Hypothesis
- Prediction
The Scientific
Process
Designing
- Investigative
- Identifying
- Pre-lab
Observing and Measuring
-Recording Data
-observing
- measuring
- estimating
Questioning
• Problem statement: What affect
does the (manipulated) have on
the (responding)
• Manipulated variable: changed
• Responding variable: reacts
• Controlled variable: stays the
same
Proposing Ideas
• Background info: Info you know
from class and experience
• Hypothesis: I hypothesize that
(answer to problem) because
(reason based on background)
• Prediction: statement using
values
Designing Experiments
• Overview: Design in 2-3
sentences.
• Experimental group vs
control group - one variable
changed
• Pre-lab calculations
• Materials & procedure:steps
Observing & Measuring Data
• Observations:
qualitative vs quantitative
Be specific & detailed
• Record Data: - in tables &
charts. Remember a title.
Processing Data
• Analysis: answer any questions,
complete calculations or graph
• Graphing: follow the rules
outlined in the booklet.
Remember a title & use pencil.
The 3 types of graphs are pie , bar
& linear(straight, sigmoid,
exponential, or curved).
Interpreting Data
• Evaluation:
1) Hypoth/pred/purp supported?
2) Experimental design?
3) Errors? - human or equipment
4) Conclusions
• Extension: Other related
experiments.
• Bibliography/References
Answers to Sig. Digs W/S
1. a. 3 2. a. 5.808 x 103 3.a. 60
b. 3
b. 6.3 x 10-5
b. 6200
c. 2
c. 5.30 x 102
c. 740
d. 2
d. 6.030 x 10-2 d. 0.0000000091
e. 3
e. 7.0 x 10-1
e. 0.003076
f. 2
f. 5.8 x 101
f. 4.3
Answers to Sig. Digs W/S
4. a. 400.9 g
b. 333.5 J
c. 83 m
d. 57 kg
e. 0.00247 g
f. 1.95 x 104 s
g. 4 x 105 g
h. 8.1 x 1010
i. 0.000018
Answers to Sig. Digs W/S
5.
A graph to show the change in concentration of product
with time
Concentration of
Product (mmol/L)
12.00
10.00
8.00
6.00
Series1
4.00
2.00
0.00
0.00
1.00
2.00
3.00
4.00
Time (seconds)
Slope = 2.11 mmol/L/s
5.00
6.00
Review of
Science 10
Organization
MATTER
PURE SUBSTANCE
Compounds
Ionic
Molecular
MIXTURES
Elements
Metals
Metaloids
Solutions
HOMOGENOUS
Non metals
Mechanical Mixtures
HETEROGENOUS
Some Additional Definitions
• Matter - object that has mass &
occupies space
• Pure Substance - substance
that has no impurities
• Mixture - made of 2 or more
substances
• Compound - two or more
elements in fixed
proportions
• Element - chemical building
block; an atom table.
• Solution - a uniform mixture
where the parts are not visible
(Also called a homogeneous mix)
ie) salt water
The two parts to a solution are:
• Solvent-large dissolving partoften H2O
• Solute - smaller dissolved partsalt
• Mechanical mixture: nonuniform mixture where the
parts remain visible and
intact.
Also called heterogeneous
mixture
ie) oil and water
Properties of Matter
Physical
• Malleability
• Color
• Lustre
• Shape
• State
• Heat capacity • Bp
• Density
• Mp
• Ductility
• Solubility
Chemical Properties
• describe how a substance reacts
• 5 evidences of chemical change
1) precipitate
2) color change
3) odour formed
4) energy change
5) gas formed
• Examples of Physical Change
1) dissolving
2) phase change
3) dividing up
• Examples of chemical changes
- reaction where
1) a new gas is released
2) new solid is formed
3) new odour is produced
4) energy change occurs
5) new color appears
The Periodic Table
Metals, metalloids and nonmetals
Information on the table
• Name - 109 naturally occurring
elements
• Family or Group - 18 columns,
each with similar properties
• Period - 7 rows, each with same
electron orbitals
• Atomic number - # or protons
• Atomic molar mass - ave mass
• physical properties - bp, mp, etc
An Introduction to
Chemistry
Atomic Structure
• An atom is the smallest part of an element
that retains the properties of that element.
• The modern atomic theory was developed
based on the work of the many scientists.
• Greek philosophers in the centuries before
Christ’s birth believed that all matter was
composed of tiny, indivisible particles that
they called atoms.
• Their ideas prevailed for almost 2000 years
until the scientific revolution of the 19th
century.
Dalton’s Atomic Theory
• Dalton’s proposal was based on many
experimental observations by him and
other scientists when elements react and
compounds are formed.
• Dalton’s Atomic Theory states:
• Atoms are tiny, indivisible particles of
elements
• All elements are composed of atoms
• Atoms of the same element are identical,
but the atoms of different elements are
different
• Atoms can combine in fixed ratios to form
compounds.
• Chemical
reactions
occur
by
a
rearrangement of atoms NOT by changing
atoms.
• Dalton basically said that an atom is a solid
sphere similar to a billiard ball:
J. J. Thompson (1897)
• Thompson used a device called a cathode ray
tube. This is essentially a vacuum tube
containing a gas at low pressure.
• He found evidence that a stream of negatively
charged particles was produced from the atoms
in the tube. Thompson called these negative
particles electrons.
• Shortly after electrons were discovered, scientists
carried out experiments with hydrogen and
discovered positively charged particles that they
called protons.
• Thompson’s atomic theory uses the
following model:
• Thompson said that an atom is a sphere
with embedded electrons, similar to a
raisin bun.
Positive sphere
Negative electrons
Rutherford’s Atomic Theory
• Rutherford experimented using gold foil
and small positively particles called alpha
particles.
• He shot the alpha particles at a thin piece
of gold foil. In his experiments, he found
that the majority of the alpha particles
passed straight through the gold foil, but
some bounced back.
• Based on this evidence, he concluded that
there is a very dense region within the gold
foil that was causing the deflection of the
alpha particles.
• Since
like
charges
repel,
Rutherford concluded that this
small dense region must be
positively charged.
• He called this region the
nucleus of the atom and
suggested that it was surrounded
by
mostly
empty
space
containing the electrons. The
positively charged particles in the
nucleus he called protons.
• Later experiments (Chadwick,
1932) showed that the nucleus
also contains dense, neutral
particles
that
were
called
neutrons.
Bohr’s Atomic Theory
• Bohr postulated a new revised atomic theory
stating that electrons have specific energies and
can only occupy certain areas around the
nucleus
• Bohr’s atomic theory is summarized as follows:
• Electrons are found at different energy levels
around the nucleus
• Electrons closest to the nucleus have the least
amount of energy
• The higher the energy level of an electron, the
further it is from the nucleus
• Electrons can move from one energy level
to another by gaining or losing energy, but
cannot exist between energy levels
• Each energy level can hold a maximum
number of electrons:
o The first energy level has a maximum of 2
electrons
o The second energy level has a maximum of 8
electrons
o The third energy level has a maximum of 8
electrons
• An atom with the maximum number of
electrons in its outermost energy level is
stable and is therefore unreactive
• Bohr’s atomic theory enables us to explain
a great deal about the properties of
elements
and
how
they
behave,
particularly how and why they react in the
way they do.
• It also explains the groupings of elements
in the periodic table and periodicity.
• It is the arrangement of electrons at
different energy levels that determines all
these things, especially the electrons in the
outermost energy level.
• The Bohr model will be used primarily in
this course to explain chemical properties
of matter.
Some Bohr atom models:
Quantum Mechanical Theory
of Atomic Structure
• Although the Bohr model is useful in allowing us
to explain matter and its properties, it is really
an over simplification.
• In reality the structure of an atom is much more
complex and mathematical, particularly with
regard to the position of electrons in their orbits
around the nucleus.
• This is referred to as Quantum mechanics.
Quantum Theory
Structure of an atom
• All atoms are composed of protons,
neutrons and electrons.
• The nucleus of an atom is composed of
positively charged protons and neutrons,
which are neutral. This contributes to the
mass of the atom.
• The electrons are found revolving in
regions of space around the nucleus.
• The nucleus is very small in comparison
to the area occupied by the electrons.
• The usual analogy is that of an ant,
representing the nucleus, at the centre of
a football field, which represents the total
size of the atom.
Summary of Atomic
Structure
Atomic number
• This number is assigned based on number of
protons in the nucleus.
• fluorine has 9 protons in its nuclei, the atomic
number is 9
• oxygen has 8 protons in its nuclei, the atomic
number is 8
• Atoms are electrically neutral, which means the
number of electrons is equal to the number of
protons.
• Therefore, the atomic number of an element is
equal to both the number of protons and the
number of electrons.
Mass Number
• Is the total number of protons and
neutrons in the nucleus.
• Each proton and neutron has the same
mass; therefore in a carbon atom with 6
protons and 6 neutrons, the mass number
is the sum of the protons and neutrons,
which are 12.
• For Be the atomic number is 4, so there
are 4 protons in the nucleus. But, if there
are 5 neutrons the mass number is 9.
• Since the mass of each sub-atomic particle
is very small, they are assigned arbitrary
units.
• The mass of each proton and neutron is
1 atomic mass unit (amu).
• The mass of an electron is so small as to
be negligible and is not considered in the
mass of an atom.
Isotopes:
• Atoms that have the same number of protons
but may contain a different number of neutrons.
• Carbon has a number of naturally occurring
isotopes, for example:
•
carbon -13
carbon -14
mass number = 13
# of protons is 6
# of neutrons is 7
mass number = 14
# of protons is 6
# of neutrons is 8
• We represent isotopes in the following way, this
is carbon – 13
13
C
6
• Here 13 is the mass number and 6 is the
atomic number
Atomic mass
• As given on the periodic table is the
average mass of all naturally occurring
isotopes in a pure sample of that element.
• It is not a whole number.
• It must not be used to determine the
number of neutrons in an atom.
Valence electrons
• Are the outermost electrons of an atom,
occupying the outer energy level.
• For the representative elements the
number of valence electrons is the same
as the last digit of the group number.
• The number of energy levels that contain
electrons is the same as the period (row)
number for that element in the periodic
table.
Elements and Isotopes
Chemical Names and
Formula
• The names of compounds and their
formula follows the rules laid down by
the International Union of Pure and
Applied Chemistry (IUPAC)
Atoms and Ions
• Ions are atoms or groups of atoms that gain or
lose electrons to become:
• Cations are positively charged ions formed by
atoms losing electrons.
• Formed from metallic elements, E.g:
• Na+
• Mg2+
(lost 1e-)
(lost 2e-)
• Names of cations remain the same as that of the
element with ion added,
• e.g. sodium ion, magnesium ion
• Anions are negatively charged ions formed
by atoms gaining electrons.
• Formed from non-metallic elements,e.g.
• Cl- (gained 1e-)
• O2- (gained 2e-)
• Names of anions have
changed to –ide. E.g.
• chloride ion, oxide ion
their
ending
• When ions are formed the number of protons
remains the same while the number of electrons
changes so that the electron arrangement is the
same as the nearest noble gas.
• This is more stable than the arrangement of the
electrons in the atom.
• Sodium atoms have 11 protons and so 11
electrons.
• When an electron is lost, its ions have only 10
electrons with a net charge of 1+, i.e. Na+.
• These electrons are arranged like those in neon,
the nearest noble gas.
Atoms and Ions
Activity - Complete the table:
Symbol of
element
Change in
electrons
Formula of
ion
Name of
ion
Ca
2 lost
Ca2+
calcium ion
F
1 gained
F-
fluoride
ion
Al
3 lost
Al3+
aluminium
ion
Se
2 gained
Se2-
selenide
ion
Compounds
• Compounds
are
pure
substances
consisting of more than one type of atom.
• Formed when atoms of two or more
elements combine chemically.
• They do so in fixed proportions by mass,
since they combine in fixed ratios.
Molecular Compounds
• Tend to have low boiling and melting
points, therefore are liquids or gases at
room temperature.
• Compounds of two or more non-metallic
elements are composed of molecules.
• Molecules are neutral groups of atoms
that act together as a single unit due to
the atoms being bound together by
chemical bonds.
Ionic Compounds
• Are composed of positive ions (cations) and
negative ions (anions) arranged in a 3D pattern
(lattice) to form crystals
• In a crystal, each cation is surrounded by anions
and visa versa
• Since the negative and positive charges (ions)
are balanced, ionic compounds are electrically
neutral.
• Tend to have high melting points and are
therefore crystalline solids at room temperature
• Formed from the combination of metal and nonmetallic elements
Chemical Formula
• Show the kind of atoms and the number of
them present in a Representative Unit of
the compound.
• A complete chemical formula should also
show the state of the material at SATP.
(Standard Atmosphere Temperature and Pressure)
Molecular Formula
• Chemical formula of molecular compounds.
• The number of each kind of atom is shown
by a subscript after the symbol of the
element, e.g. H2O(l)
• There is no indication of the shape or
structure of the molecule.
Ionic Formula
• Chemical formula of an ionic compound.
• The formula shows the ions contained in
the compound, placing the symbol of the
metal first.
• Ionic compounds do not exist as distinct
units, unlike molecules they are composed
of fixed ratios of ions to provide a neutral
compound.
• Formula of ionic compounds do not
represent molecules but show the lowest
whole number ratio of ions in the
compound,
• e.g. sodium chloride has a ratio of 1:1
sodium to chloride ions so the formula is
NaCl(s)
• This is known as a formula unit of an ionic
compound.
• Example:
• Magnesium chloride is composed of Mg2+ ions
and Cl- ions.
• To make this compound electrically neutral, two
chloride ions must be present to balance each
2+ magnesium ion.
• The ratio of Mg2+ to Cl- ions is 1:2, therefore a
formula unit of magnesium chloride is MgCl2(s)
Ionic Charges on the Elements
• To be able to state what compounds are
formed we need to know what types of ions
are formed from different elements, i.e. their
ionic charges.
• For the representative elements the charge
can be easily determined from the periodic
table.
Metals:
• In groups 1, 2 and 3 lose electrons to form
positive ions (cations).
• The charge on the ion is equal to the group
number and equal to the number of
valance electrons.
• For example: magnesium:
2+
2Mg
Mg
Non - Metals:
• Here the opposite is true, these elements
form negative ions (anions) by gaining
electrons.
• The charge can be determined by
subtracting the last digit of the group
number from 8.
• For example: oxygen:
2-
2O
O
• Elements in Group 14:
• These elements do not easily form ions,
instead, they share electrons in covalent
bonds
• Elements in Group 18:
• The noble gases have a full outer energy
level and therefore are very stable and do
not form ions.
Polyatomic Ions
(complex ions)
• Polyatomic ions are groups of atoms that
together act as a single unit with an overall
charge,
• e.g. SO42- is a group of 1 sulfur and 4 oxygen
atoms with an overall charge of 2• The names and formula of polyatomic ions are
provided on your periodic table. You should learn
be able to be able to recognize these complex
ions.
Binary Ionic Compounds
• Are composed of two mono-atomic ions,
one cation (metal) and one anion (nonmetal).
• The name of the cation is kept the same
as the name of element and always
comes first, while the ending of the name
of the anion is changed to -ide (note
neither are capitalized).
• E.gs. sodium chloride; calcium oxide
• The formula of binary ionic compounds must
show that charges balance and must be in the
lowest possible whole number ratio:
• Sodium chloride is formed from the combination
of Na+ and Cl- ions. Thus it becomes NaCl(s)
• The lowest common multiple of the ion charges
can be used to determine the correct ratio of ions
in the formula:
• Aluminium oxide is from Al3+ and O2• The lowest common multiple is 6.
• 3 goes into 6 twice and so there must be
two aluminium ions in the formula.
• 2 goes into 6 three times meaning that
there will be 3 oxide ions.
• Therefore, the formula is Al2O3(s)
• The criss/cross method is the simplest
way to write formulae of ionic compounds.
• Here the numerical charge from each ion
is crossed over and used as the subscript
for the other ion:
• E.g. iron (III) oxide is from Fe3+ and O2Criss/crossing the two charges gives
Fe2O3(s)
Molecular Compounds
• These compounds are combinations of
non-metallic
elements.
They
are
composed of molecules and therefore ionic
charges are NOT a factor.
• When two non-metallic elements combine
there are often different possible
combinations
producing
different
compounds.
• These different compounds must be
distinguished between since they have
different chemical and physical properties.
• Prefixes are used to show how many
atoms of each element are present in the
molecule.
• In the formula these prefixes are
represented by a subscript number after
the symbol for the element.
• The name of the first element remains the
same (after any prefix that is necessary)
followed by the second element’s name,
the ending of which is changed to -ide.
• Generally the name goes:
• prefix-element name, prefix-element root -ide.
• Often mono- is dropped as a prefix to the
first element.
• Writing formula of molecular compounds
is simply done by using the prefix to
determine the subscript number of that
element. For example:
Reactants
Product formula
Product name
C(s) and S8(s)
CS2(l)
carbon disulphide
N2(g) and I2(s)
NI3(s)
nitrogen triiodide
N2(g) and O2(g)
N2O(g)
dinitrogen monoxide
P4(s) and O2(g)
P4O10(s)
tetraphosphorous
decaoxide
Molecular Elements:
• Most elements are found as single atoms,
i.e. they are mono-atomic.
• All metals are like this and some nonmetals, however, some elements are found
as diatomic molecules and even
polyatomic molecules.
• This arrangement allows them to be more
stable.
• Hydrogen,
nitrogen,
oxygen
and
the
halogens are all diatomic. Their formulas
are written H2(g), O2(g), Cl2(g) etc.
• Sulfur (S8(s)) and phosphorus (P4(s)) are
polyatomic.
Acids and Bases
• Acids have special names.
• Acids are defined as compounds that release
hydrogen ions when dissolved in water.
• Acids and Bases will be reviewed and discussed
in Science 30.
Chemical Reactions
• Chemical reactions always involve
reactants changing to products.
• This involves the rearrangement of the
atoms of the reactants into new
substances as the products.
• Particles of the reactants must collide for
reactions to occur.
• Reactants must have a certain minimum
energy for reactions to occur.
• Chemical reactions do not cause atoms to
be created or destroyed.
• Changes in energy are always associated
with chemical reactions due to the
breaking and/or formation of chemical
bonds.
 Other evidence for chemical reactions
includes production of a solid, gas or odor
or a change in colour.
Features of Chemical
Reactions
Chemical Equations:
• Chemical equations are used to represent
what happens in a chemical reaction.
• They always indicate:
Reactants  Products
• They always show that matter is
conserved i.e. that no matter is created or
destroyed.
Word Equations:
• Show the names of the chemicals
involved in the reaction and the products:
E.g.
• iron (III) + oxygen  iron (III) oxide
Balance Chemical
Equations:
• Show the quantities of the reactants and
products involved in terms of the number of
atoms of each substance.
• Use a symbolic representation of the
substances as their chemical formulas.
• Have the same number of each type of atom in
the reactants and the products, i.e. they are
balanced in accordance with the law of
conservation of matter.
• Show the states of matter of the reactants and
products.
• E.g. 4Fe(s) + 3O2(g)  2Fe2O3(s)
Rules for Balancing Chemical Equations
1.
Determine the correct chemical formula for
all reactants and products.
2. Write the formulas for the reactants on the
left and those for the products on the right
and separate them with an arrow. Place a plus
sign between two or more reactants or
products.
3. Indicate the state of each of the reactants
and products.
4. Count the number of atoms of each element
on both the reactant and product sides of
the equation. Polyatomic ions that remain
unchanged on both sides of the equation are
counted as a single unit.
Rules for Balancing Chemical Equations
5.
•
•
•
•
6.
7.
8.
Balance the number of atoms of each element on
both sides of the equation using coefficients.
Place the coefficient in front of the chemical formula;
This multiplies all of the atoms in the formula by the value
of the coefficient (when there is no coefficient it is
taken to be 1)
Coefficient must be whole numbers
Start with elements other than hydrogen and oxygen
When balancing hydrocarbon combustion, start with
carbon always, followed by hydrogen and finally oxygen
Check that the number of each atom or polyatomic
ion is the same on both sides of the equation
Make sure that the coefficients are the lowest
possible ratio
Practice, practice, practice!!!
Interpreting balanced
chemical equations
• So far we have considered that we are
working with individual atoms, ions or
molecules of the reactants and products.
• However, these entities are much too
small to see and so observable changes
in chemical reactions must involve very
large numbers of particles.
• A convenient way to communicate these
enormous numbers has been developed by
scientist.
• The term mole is used to represent these large
numbers in a convenient way. Just like one
dozen always represents 12 eggs or donuts, one
mole represents a certain number of particles.
• This number is called Avogadro’s Number and
it is a rather unusual number, 6.02 x 1023.
• One mole of any substance is always
6.02 x 1023 particles of that substance.
• Essentially, a mole represents a certain number
of something just like a dozen does.
• Now we can interpret chemical equations a little
differently.
• For the following reaction we can talk about
individual particles, dozens of particles or
Avogadro’s number of particles:
• 4Fe(s)
+
3O2(g)

2Fe2O3(s)
4 atoms
4 dozen atoms
3 molecules
3 dozen molecules
2 formula units
2 dozen form u’ts
24.08 x 1023 atoms
4 mol of atoms
18.02 x 1023 molecules
3 mol of molecules
12.04 x 1023form u’ts
2 mol of form units
• The mole concept becomes even more useful
when we find that we can relate moles to the
mass or volume of reacting substances.
• We will deal with the mole concept in much
more detail later.
Types of Chemical Reaction
• To be able to predict the outcome of a
chemical reaction and write a complete
equation, you must be able to recognize
the type of reaction that is occurring.
There are five basic reaction types:
1.
2.
3.
4.
Formation
Decomposition
Single replacement
Double replacement (neutralization reactions
are a special kind of double replacement)
5. Combustion
Types of Chemical Reaction
Formation reaction
• This is a reaction of two or more elements
forming an ionic or molecular compound.
• E.g. magnesium reacts with oxygen to form
magnesium oxide.
•
2 Mg(s) +
O2(g)  2 MgO(s)
• Summary:element A + element B  compound AB
Decomposition reaction
• This is a breakdown of a compound into simpler
compounds or into the component elements.
• E.g.1 calcium carbonate decomposes into
calcium oxide and carbon dioxide.
•
CaCO3(s)  CaO(s) + CO2(g)
• E.g.2 water decomposes into its constituent
elements by using electricity, electrolysis
•
2 H2O(l)  2 H2(g) + O2(g)
• Summary: compound AB  element A + element B
Single replacement reaction
• This is a reaction of an element with a
compound to produce a new element and a
new compound.
• Basically the two elements switch around.
• Occurs in aqueous solutions, usually involving
replacement of metal ions.
• E.g. copper solid reacts with a silver nitrate
solution.
• Cu(s)
+
2AgNO3(aq)  2Ag(s)
+ Cu(NO3)2(aq)
• Summary: element A + compound BC  element B + compoundAC
• This type of reaction can also happen with
halogens.
• E.g.
Cl2(g) + 2NaBr(aq)  Br2(l) + 2NaCl(aq)
Double replacement Reaction
• This is a reaction between two ionic compounds
in solution to produce two new compounds.
• Occurs in aqueous solution.
• Generally, the new compounds are produced
by switching the partners of the negative ions
that are found in the reactants.
• E.g. calcium chloride solution is mixed with a
sodium carbonate solution .
CaCl2(aq) + Na2CO3(aq)  CaCO3(s)
+
2NaCl(aq)
compound AB + compound CD  compound AD + compound CB
• A specific double replacement reaction called a
neutralization reaction occurs when an acid
reacts with a base to produce water and an
ionic salt.
E.g. Hydrochloric acid reacts with potassium
hydroxide.
•
HCl(aq) + KOH(aq) 
KCl(aq) +
H2O(l)
acid
+ base

salt
+
water
• Double replacement reactions always produce
either a solid or a molecular compound.
Combustion reactions
• This reaction type involves oxygen as a
reactant.
• It is an exothermic reaction.
• The products being produced in a
combustion reaction must be determined
by the composition of the substance being
burned.
CHEMICAL REACTIONS
•
If the substance that
combustion contains:
is
undergoing
i. carbon, then CO2(g) is produced.
ii. hydrogen, then H2O(g) is produced.
iii. carbon and hydrogen, then CO2(g) & H2O(g) is
produced. This is a hydrocarbon Combustion.
iv. sulfur, then SO2(g) is produced.
v. a metal, then the oxide of the metal with the most
common ion charge is produced.
Endothermic and
Exothermic Reactions
• End of Review