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Science 1206
Unit 3 – Chemical Reactions
Text: Chapters 5-8
Chemistry:
the study of matter; its properties and its changes.
Matter:
anything that has mass and takes up space (energy is not matter).
Three states of matter: solid, liquid and gas.
Lab Safety:
Workplace Hazardous Materials Information System (WHMIS):
Three parts:
 Labeling
 Worker education and training
 MSDS (Material safety data sheets)
WHMIS Safety Symbols – labeling (see p.658):
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MSDS – Material Safety Data Sheets:
Provides workers and emergency personnel with the proper procedures for handling and
working with a particular substance.
Classification of Matter
Classification of matter as pure substances or mixtures:
Matter can be classified into 2 groups:
1. Pure substances
-
have constant composition; all the particles that make up the
substance are the same.
Two divisions:
a) Elements:
 The simplest form of matter that can exist under normal
conditions.
 Composed of only one kind of atom. Eg. Carbon, magnesium
 Cannot be broken into simpler substances by chemical means
 Combine to form other substances
b) Compounds:
 Substances composed of 2 or more different kinds of atoms. Eg
H2O
 Can be broken down into simpler substances by chemical means
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2. Mixtures
Have variable composition; composed of 2 or more pure substances.
Two groups:
a) Homogeneous mixture (solutions)
 Have only1 visible component/phase
 eg. Tap water. Air, sugar solution (sugar + water)
b) Heterogeneous mixture (mechanical mixture):
 Have 2 or more visible components
 Eg, sand in water, vegetable soup, tossed salad.
Properties of Matter – 2 Types
1. Physical Property
 Characteristics of matter, used to identify substances
 Eg. State at room temperature, boiling point (found on periodic table),
color, mass, electrical conductivity
2. Chemical Property
 Characteristic of matter that can be observed when matter undergoes a
change in composition;
 Describes “how it reacts”
 Eg. Propane reacts with oxygen to produce carbon dioxide and water
Changes of Matter – 2 Types
1. Physical Change:
 A change in the size or form of a substance that does not change its
composition.
 Eg. Cutting paper, melting ice.
2. Chemical Change:
 A chemical reaction; a change in which at least 1 or more “new”
substances (products) are formed.
 The products have different properties from the starting substances
(reactants)
 Eg. Rust, burning, cooking
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Example: Rusting
Chemical equation for rusting:
Fe(s) + O2(g)  Fe2O3(s)
The rust produced has completely different properties from iron and oxygen.
Evidence of Chemical Change





Change in color
Energy produced (heat and/or light)
Gas produced (bubbles)
New solid produced (called a precipitate)
Difficult to reverse
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Elements and the Periodic Table
All elements are classifies as metals or nonmetals, depending on their properties.
Property
Metals
Nonmetals
Lustre
Malleability
Conductivity of heat and
electricity
State at room temperature
shiny
Malleable (bendable)
Good conductors
(electrolyte)
All solids except mercury
(liquid)
Reactivity with acid
Location (periodic table)
Mostly yes
Left of staircase lone
Dull
Brittle
Poor or nonconductors
(non-electrolyte)
Most are gases, some are
solids and bromine is a
liquid
no
Right of staircase line
Metalloids (semimetals):
 Elements that have some properties of metals and some of nonmetals.
 Include all elements on either side of the staircase line except AL and At.
 Also includes one form of C, called graphite, which is dull and brittle (nonmetal)
but is a good conductor of electricity (metal).
Chemical Families (groups)
Groups of elements in the same vertical column that have similar physical and chemical
properties.





Group 1 (IA) – alkali metals
Group 2 (IIA) – alkaline earth metals
Group 17 (VIIA) – halogens
Group 18 (VIIIA) – noble gases
Rows on the bottom:
 Lanthanide series
 Actinide series
Properties of chemical families:
Alkali metals:
-
-
group 1, IA
Eg. Sodium, lithium
Show metallic properties
Highly reactive, especially with water; reactivity increases down the group
Cs and Fr are the most reactive metals.
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Alkaline earth metals:
-
Group 2, IIA
Eg. Calcium, magnesium
show metallic properties
less reactive than alkali metals; reactivity increases down the group
Halogens:
-
Group 17, VIIA
Eg. Chlorine, fluorine
show nonmetallic properties
reactivity decreases going down the group; F is the most reactive nonmetal
react with metals to produce salts (ionic compounds)
react with hydrogen to form compounds that dissolve in water to form acids.
Noble Gases:
-
Group 18, VIIIA
Helium, neon, argon
show nonmetallic properties
extremely low chemical reactivity
Hydrogen:
-
the lightest and most abundant element
doesn’t really belong to any group (can be group 1 or 17)
it sometimes behaves like an alkali metal (group 1) and at other times as an acid.
The ATOM
-
the basic building block of all matter
the smallest particle of an element that retains the properties of that element
electrically neutral: the number of positive charges (protons) equals the number of
negative charges (electrons)
composed of 3 subatomic particles: protons (p+), electrons (e-) and neutrons (no)
Particle
Symbol
Relative
charge
Mass (g)
Location
Proton
Neutron
electron
p+
eno
1+
0
1-
1.67 x 10-24
1.67 x 10-24
9.11 x 10-28
Nucleus
Nucleus
orbital
Models of the atom – see handout
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a. Atomic nucleus is located at the core of an atom and it contains protons and
neutrons.
[i] Proton: A subatomic particle that carries a positive one (+1) electrical charge and
has a mass value of 1 a.m.u. (atomic mass units)
The number of protons in the atomic nucleus determines an element's Atomic
Number.
[ii] Neutrons:A subatomic particle that has no electrical charge (0) and has a mass
value of 1 a.m.u.
Atomic Mass is the sum of the number of protons and neutrons found in an atom of
any element.
b. Energy levels are regions of subatomic space that lie outside the atomic nucleus.
Electrons are found in the energy levels. The energy levels are also called quantum
levels.
[i] Electrons
These subatomic particles carry a negative one (-1) electrical charge.
In a neutral atom of any element the number of protons equals the number of
electrons.
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Atomic number:
- identifies the element
-
equal to the number of protons in the nucleus
since atoms are electrically neutral, # of protons = # of electrons
Mass number:
-
# of protons + # of neutrons
# neutrons = atomic mass (rounded off) – atomic #
Protons and neutrons account for most of the mass of the atom
Electron Energy Diagrams for Atoms (Bohr diagrams):
 an energy level represents a specific value of energy of an electron and




corresponds to a general location around the nucleus
the number of occupied energy levels in any atom is normally the same as
the period number in which an atom appears
for the first 3 energy levels, the maximum number of electrons that can be
present are 2, 8 and8 in order of increasing energy (increasing distance
from the nucleus)
a lower energy level is filled with electrons to its maximum level before
the next level is started.
The electrons in the highest (outermost) occupied energy level are called
valence electrons. Number of valence electrons is the same as the group
number for group A elements (1,2,13-18).
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Stable Atoms

The outermost energy level of any atom is called the valence level. The
electrons found in this level are called valence electrons. The number of
valence electrons an atom has and the way they are arranged determines
whether or not it will form a compound and the type of compound it
will become a part of.
 have low chemical reactivity
 include noble gases, all of which have 8 valence electrons (except He,
which has 2) stability is a function of having a full complement of
valence electrons. Atoms that do not have full electrons energy levels are
unstable and must gain, lose or share electrons to become stable.
 other atoms can become more stable by reacting and changing the number
of their electrons; thereby attaining the same stable electron configuration
(structure) of the nearest noble gas:
 atoms can follow one of 2 rules:
a) Octet rule atoms attempt to obtain 8 valence electrons
includes most atoms
b) Duet rule atoms attempt to obtain 2 valence electrons
includes H, Li and Be
Atoms can achieve a stable octet or duet by forming ions.
Ions
-
an atom or groups of atoms that have a positive or negative charge, due to the loss
or gain of one or more electrons.
Single atoms form simple ions (monatomic ions); groups of atoms form complex
ions (polyatomic ions)
Example:
sodium metal and chlorine gas react to produce NaCl, a very stable and
unreactive substance, compared to Na or Cl. The sodium atom loses 1
electron to the chlorine atom so both of their outer levels are filled. In
doing so, the atoms form ions of opposite charge.
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How ionic bonds form:
Begin with atoms of two different elements that do not have 8 electrons in their
outer most energy level.
The sodium atom (Na ) donates 1 electron becoming a positively sodium ion ( Na+).
The chlorine atom (Cl) accepts the donated electron becoming a negatively charged
chloride ion (Cl-).
This chemical reaction produces table salt (NaCl).
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Two types of ions:
a) Cations:




Ions form when atoms lose electrons
Metals form cations
Have a positive charge because they have more protons than electrons
Name stays the same (sodium atom written as sodium ion)
b) Anions:




Ions form when atoms gain electrons
Nonmetals form anions
Have a negative charge because they have more electrons than protons
Name changes to “-ide”. (chlorine atom becomes chloride ion)
Note:
1. Both cations and anions are more stable than the atoms from which they form
since these ions have the same stable electron structure/configuration as the
nearest noble gas.
2. Boron, carbon and silicon do not tend to form ions (they share electrons with
other atoms)
3. The noble gases do not form ions since they are already stable (have filled
orbitals)
4. Hydrogen can form a cation or an anion:
 Cation: H+, hydrogen ion has 1 proton but no electrons
 Anion: H-, hydride ion has 1 proton and 2 electrons
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Types of Compounds
1.
Ionic Compounds:
involve the transfer of electron(s) between 2 oppositely charged ions
(cation and anion)
metal and a nonmetal or a combination involving a complex ion
forms an ionic bond
exists as an ionic crystal lattice (not individual mlecules
-
known as a formula unit (eg. A formula unit of salt, not a molecule)
Formula Unit: a chemical formula showing the simplest whole number ratio of cations
to anions in an ionic compound.
2.
Molecular compounds
-
involve the sharing of electrons between nonmetals
forms a covalent bond
exists as individual molecules
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Properties of Ionic and Molecular Compounds
1.
State at room temperature:
all ionic compounds are solids
-
2.
Conductivity of solution:
-
3.
molecular compounds may be a solid, liquid or a gas
ionic compounds conduct electricity (they are electrolytes)
molecular compounds do not conduct electricity (non-electrolytes)
Solubility in water:
ionic compounds are soluble, to varying degrees (some better than others)
-
and form colored or colorless solutions.
molecular compounds may or may not be soluble
Nomenclature
Chemical nomenclature is the systematic naming of chemical compounds.
Compounds can be divided into two basic categories, those which are true binary
compounds (they contain only two types of elements), and those which contain more than
two different types of elements.
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Rules for Naming ionic compounds:
Identify the type of ions:
A.
Monoatomic or simple ions
-
B.
Single atoms that have lost or gained one or more electrons
Form binary ionic compounds (2 simple ions)
Eg. Sodium + chlorine
Na+ ClPolyatomic or complex ions
-
cations or anions composed of a group of atoms with a net positive or negative
charge.
-
eg.
Nitrate NO3Ammonium
C.
NH4+
Multivalent ions
-
certain transition metals can form more than one type of ion, each with a different
charge.
-
the one written on top is the more common ion
-
eg. Cu2+ or Cu+
D.
Hydrates
-
ionic compounds that contain water in their structure
eg. CuSO4 H2O
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Naming Ionic compounds; (4 types)
1. Binary Ionic compounds



Consist of cations and anions
Cations are written first, anions are second (name changes to “-ide”)
The total charge must be zero
Rules:
i. write the symbols for the ions involved
eg. Silver and chlorine
Ag+ and Clii. determine the lowest whole number ratio of ions which will
provide an overall net charge of zero
Ag1+ Cl1- becomes AgCl (silver chloride)
iii. do not write charges in your final answer
iv. if one of your charges is odd, you can use the criss-cross method
Example:
potassium and oxygen:
Potassium Oxygen -
aluminum and sulfur
K+
O2-
aluminum sulfur -
K2+ O1- becomes K2O
Potassium oxide
Al3+
S2-
Al23+ S32- becomes Al2S3
aluminum sulfide
Naming Ionic compounds:
 Name the cation (positive ion) by writing the full name of the metallic element
 Name the anion by abbreviating the nonmetallic element to “ide” (chlorine to
chloride.
 Practice:
NaCl
BaCl2
Al2O3
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2.
Polyatomic/Complex Ions
Complex ions are groups of atoms that are made stable by sharing electrons and which
then become even more stable by gaining (usually) or losing electrons. Unlike neutral
molecules, complex ions carry an electric charge and do not exist by themselves. An
ionic bond is formed by the attraction of a positive simple ion to a negative complex ion
or of a positive complex ion (NH4+) to a negative simple or complex ion. The total
positive charge in the formula must be equal to the total negative charge.
Polyatomic ion: atoms of 2 or more elements covalently bonded together with an overall
charge
Nitrate NO3Ammonium NH4+
eg.
Rules:




Don’t change the ending of a polyatomic ion!
Name the cation, then name the anion
Balance the charges
If you need more than 1 complex ions, use brackets for that group
Example:
1. Sodium ions and carbonate ions bond ionically to form an ionic compound.
Na+
CO32- 
Na2CO3 sodium carbonate
2. Ammonium ions and hydrogen phosphate ions bond ionically to form an ionic
compound.
NH4+ HPO4-  (NH4)2HPO4 ammonium phosphate
3. magnesium ions and hydroxide ions bond ionically.
Mg 2+
Practice
NaNO3
Al2(SO4)3
Mg(OH)2
NaCH3COO
OH-  Mg(OH) 2
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3.
Multivalent
-
Certain transition metals can form more than one type of ion, each with a
different charge.
The transition metals have various electron configurations that will make
them stable.
the one written on top is the more common ion.
eg. Cu2+ - copper (II)
Cu+ - copper (I)
Fe3+ - iron (III)
Fe2+ - iron (II)
-
-
-
To name a multivalent metal compound, you must specify the
charge on the ion as a roman numeral in brackets (Stock naming
system).
Eg.
Iron (ii) oxide
FeO
Iron (iii) oxide
Fe2O3
Do not write charges as a roman numeral if the metal is not
multivalent.
Practice:
CuSO4
PbO
uranium (vi) oxide
uranium (iv) oxide
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4. Hydrates
-
A number of ionic compounds called hydrates produce water when they
decompose upon heating.
eg. CuSO4 5H2O copper (II) sulfate pentahydrate
-
-
When the formula of a hydrated compound is written, the number of water
molecules is also included. For example, copper (II) sulfate pentahydrate
is written as CuSO4  5 H2O, meaning 5 molecules of water are bonded
within the ionic crystal for every one formula unit of CuSO4.
Rules for naming:
a) Name the ionic part of the formula first
b) Name the water part second, using a prefix system for the number of
water molecules
c) Add prefix to “hydrate”
d) Prefixes:
mono 1
Di
2
Tri
3
Tetra 4
Penta 5
Hexa 6
Hepta 7
Octa 8
Nona 9
Deca 10
Example:
Barium hydroxide hexahydrate
Ba(OH) 2  6 H2O
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MOLECULAR COMPOUNDS
 Binary molecular compounds form between 2 non-metals
 Covalent bonds: shared electrons
 Molecular formula: shows number and kind of atoms in a molecule
Rules for naming molecular compounds:
 Write the name of the first element of the formula in full
 Second element ends with “-ide”
 Use prefixes to specify number of atoms of each element in the molecule; use the
same prefixes used for hydrates
 No charges used in formula
 The prefix “mono-” is optional on the first name
 Name the following:
 NO
 CO2
 N4O9
 N6O
nitrogen monoxide
carbon dioxide
tetranitrogen nonaoxide
hexanitrogen monoxide
 Write formulas for the following:
 Boron trifluoride
 Sulfur hexafluoride
 Nitrogen monoxide
 Phosphorous pentachloride
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The following trivial names for molecular compounds must be
memorized:
Name
Formula
H2O
Propane
Formula
C3H8
Hydrogen
peroxide
Glucose
H2O2
Ammonia
NH3
C6H12O6
Methane
CH4
Sucrose
C12H22O11
Methanol
CH3OH
Ozone
O3
Ethanol
C2H5OH
Water
Name
Molecular Elements:
Some non-metals do not exist as single atoms in nature. Some are diatomic
(containing 2 atoms) and some are polyatomic (more than 2 atoms):
Element
Formula
Element
Formula
Phosphorus
P4(s)
Nitrogen
N2(g)
Sulfur
S8(s)
Fluorine
F2(g)
chlorine
Cl2(g)
Iodine
I2(g)
Bromine
Br2(l)
oxygen
O2(g)
hydrogen
H2(g)
Note:
P.S. Clem Brown Has No Friends In Ottawa.
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ACIDS
They must contain hydrogen (H+)
They must be dissolved in water (aqueous, aq); the formula will always contain the
subscript aq.
Rules for naming acids:
Hydrogen is always the positive ion for an acid
Ending
1. -ide
Acid Name
begins with hydro, ends with -ic and acid
2. –ite
ends with –ous and acid
3. –ate
ends with –ic and acid
1. hydrogen _____ide becomes hydro____ic acid
Ex: hydrogen chloride becomes hydrochloric acid
HCl
HCl(aq)
2. hydrogen _____ate becomes______ic acid
Ex: hydrogen sulfate becomes sulfuric acid
H2SO4
H2SO4(aq)
3. hydrogen _____ite becomes____ous acid
Ex: hydrogen nitrite becomes nitrous acid
HNO2
HNO2(aq)
NOTE: when sulf is the root we add “ur” and when phosph is the root we
add “or” to make it sound better.
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Chemical Reactions
A chemical reaction represents a chemical change.
A+B
Reactants

C+D
Products
Note:
Reactants are your starting materials
The arrow means produces
Products are the new substances formed
A chemical equation is a shorthand way of representing what experimental evidence
indicates happens in a chemical reaction and must show the substances involved and the
correct number of each atom or ion involved.
A chemical equation only involves a rearrangement of the atoms involved. It does not
produce new atoms, only new substances. This is according to the Law of conservation of
matter (or mass) – matter is neither created or destroyed in a chemical change.
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Example:
Paragraph form:
iron reacts with oxygen to produce iron (iii) oxide.
Word equation:
iron + oxygen  iron (iii) oxide
Skeleton Chemical equation:
Fe(s) + O2(g)  Fe2O3(s)
Balanced chemical equation:
4Fe(s) + 3O2(g)  2Fe2O3(s)
How to write equations:
Step 1:
Step 2:
Write the correct formula and state (s, l, g or aq) for each reactant and
product
Balance the atoms without changing the formula of the compounds.
Change the number in front of the formula, called the coefficient, to show
the correct number of atoms involved. Do not show “1” in the balancing.
Generally, start by balancing the atom of which there is the greatest
number. Find the lowest common multiple of the numbers of reactants and
product atoms. Then balance the rest of the atoms. Leave the elements
until the end.
Example:
Write the equation for magnesium reacting with oxygen. (Reminder of
P.S.Clem Brown Has No friends in Ottawa).
Magnesium + oxygen  magnesium oxide
Mg(s) + O2(g)  MgO(s)
2Mg(s) + O2(g)  2MgO(s)
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Types of Reactions:
1.
Simple composition/formation
-
2 or more elements combine to form a compound
Element 1 + element 2  Compound
Eg.
Aluminum + oxygen  aluminum oxide
4 Al(s) + 3 O2(g)  2 Al2O3(s)
2.
Simple Decomposition
-
3.
4.
compound  element + element
eg.
Sodium chloride  sodium + chlorine
2 NaCl(s)  2 Na(s) + __Cl2(g)
Single replacement
-
element 1 + compound 1  element 2 + compound 2
eg.
Magnesium reacts with hydrochloric acid.
Mg(s) + 2 HCl (aq)  H2(g) + MgCl2(aq)
-
Use the solubility chart to predict the states of the products if at least one
of your reactants is aq.
Double replacement
-
Compound 1 + compound 2  compound 3 + compound 4
Eg.
Potassium iodide reacts with lead (ii) nitrate).
KI(aq) + Pb(NO3)2(aq)  PbI2(s) + KNO3(aq)
 A special type of double replacement is a neutralization reaction,
where an acid joins with a base to produce water and a “salt”
Acid + base  water + “salt”
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5.
Combustion
 Hydrocarbon + oxygen  carbon dioxide + water
 Eg. Burning propane
 C3H8(g) + O2(g)  CO2(g) + H20(g)
6.
Other
The reaction is classified as “other” if it is not one of the above.