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Review Unit
1. Matter
a) Elements and Compounds
♣ All matter is made up of about 100 elements. Elements are pure substances that cannot
be broken down into simpler parts by ordinary chemical means. Each element is
composed of a fundamental particle called an atom. Each element has a unique atom and
is represented by a symbol (memorize the sheet “Some Common Symbols”).
♣ The elements with their symbols are organized into the Periodic Table.
♣ Elements combine to form compounds (atoms combine to form molecules).
Element 1
(atom)
+
Element 2
(atom)
→
Compound
(molecule)
The physical and chemical properties of a compound are completely different from the
elements it is made from.
+
→
two hydrogen atoms
+
one oxygen atom →
one water molecule
♣ Notice that the water molecule can only be made by joining together two hydrogen
atoms (symbol = H) with one oxygen atom (symbol = O). The formula for water will be
H2O.
Note: If there is no number after a symbol in a formula, assume it is a one.
Example: CaO means Ca1O1
♣ In summary, pure substances are organized this way:
Matter
Elements → Compounds
Particles
Atoms → Molecules
Naming
Symbols → Formulas
♣ Impure substances are called mixtures. They are made by mixing different types of
compounds and/or elements together. Only mixtures have a variable composition.
1b) Measurement
♣ Measurements are fundamental to the study of matter. There are two types of
measurements:
(i)
quantitative – describes matter using numbers.
(ii)
qualitative – describes matter in a descriptive, nonnumeric way.
♣ Measurements are reported in significant figures.
Significant Figures = all measured digits + one estimated digit
(exact)
(approximate)
2
43.52 km
Example:
known uncertain
(measured) (estimated)
♣ All of the digits in a measurement are significant unless it is a zero that is used as a
place marker.
for whole numbers
for decimal numbers
6700
0.003
place markers (not significant)
place markers (not significant)
♣ To identify significant digits, try this:
(i) for whole numbers:
67 490 =
count
(ii) for decimal numbers:
(iii) for mixed numbers:
4 sig. fig.
0.004 80 =
count
3 sig. fig.
680.420 = 6 sig. fig.
count everything
If the zeros in the number were actually measured, they must be shown to be significant.
This is done by marking them with a bar or a decimal point.
6700 = 2 sig. fig. (6.7 x 103)
6700 = 3 sig. fig. (6.70 x 103)
6700. = 4 sig. fig. (6.700 x 103)
♣ Remember that significant figures are used only with measurements. They are not
applied to defined quantities (100 cm = 1 m) or to pure numbers.
♣ The use of significant figures allows us to change the units of a measurement without
changing the accuracy of the measurement.
1.24 m
=
1240 mm
=
0.00124 km
3 sig. fig.
♣ All measurements use significant figures. When using an instrument:
- make sure you understand the scale first.
- Remember that the last digit of your measurement must be an estimate
(if the spacing is very small, the only practical estimate is either exactly on the line
[estimate a 0] or halfway between the lines [estimate a 5].)
- be sure the scale is read at eye level.
- the level of a liquid is read at the bottom of its meniscus.
- Vernier scales can be used to increase the accuracy of the estimated digit.
♣ When significant figures are used in calculations, the answer cannot show more
accuracy than the least accurate measurement. It must be rounded off to the proper digit.
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(1) Addition and Subtraction – round off to the highest place value (the least accurate
measurement).
6.4 cm + 8.425 cm + 2.81 cm = 17.635 cm = 17.6 cm
the least accurate measurement (tenths)
(2) Multiplication and Division – round off to the least number of
significant digits used (the least accurate measurement).
6.4 cm x 8.425 cm ÷ 2.81 cm = 19.188 612 1 cm = 19 cm
the least accurate measurement ( 2 sig. fig.)
2. Parts of an Atom
♦ The atom has two main parts:
(i)
the nucleus (or center) of the atom contains positively charged particles called
protons (p+) and a neutral particle with no charge called a neutron (n0). It
occupies very little space but contains the mass of the atom.
(ii)
The space in an atom is occupied by negative particles called electrons (e-)
traveling around the nucleus at regular intervals. Their have almost no mass.
♦ Atoms are different because their nuclei contain different numbers of protons.
Carbon → 6 p+
nitrogen → 7p+
The number of protons in the nucleus of an atom is called its atomic number.
Oxygen → atomic number = 8 → has 8 p+
♦ Almost all of the mass of an atom is in its nucleus. The mass number of an atom, then,
is the mass of the particles in the nucleus (protons and neutrons).
Oxygen → mass number = 16 → has 16 p+ + n0
A Simplified Lithium Atom
= protons
= neutrons
= electrons
atomic number = 3
mass number = 7
atomic symbol = 73Li
Nucleus
Energy levels
♦ The first 20 electrons are placed in four energy levels ( 2e-, 8e-, 8e-, 2e-) . Their
arrangements can be shown using energy level diagrams.
Example: show the energy level diagram for potassium (19K)
K 19p+ 2e- 8e- 8e- 1e♦ Note: (i) Since atoms have no overall charge, then
number of p+ = number of e(ii) number of n0 = mass number – atomic number
(iii) Electrons travel around the nucleus in increasing intervals called energy levels.
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3. Periodic Table
◘ The modern periodic table arranges the elements in order of increasing atomic number.
◘ Metals are separated from nonmetals by the “staircase line”.
metals - shiny, malleable, ductile, conductors of heat and electricity.
◘ The columns are families (groups) of elements having similar chemical properties.
Some of these families have names to memorize (see sheet).
◘ The rows are called periods. The elements gradually change from metallic to
nonmetallic from left to right.
The number of the row tells us the number of energy levels for electrons in that atom.
◘ The Group A elements are called representative elements.
The Group B elements are called transition elements.
◘ Note that the number given under the symbol is not the mass number of that particular
atom. It is an averaged mass number called the atomic mass and will be used later.
4. Ions
a) Creating Ions
♣ We know that atoms are electrically neutral because they have equal numbers of
protons (p+) and electrons (e-).
♣ When atoms join together, though, they can lose or gain electrons. This causes the
charges in the atom to become unbalanced. Ions are atoms that have a charge.
atom
electron
ion
metal (M)
loses one
M+
nonmetal (N)
gains one
N-
♣ Some atoms will gain or lose more than one electron:
- Atoms of metals (M) tend to lose one, two or three electrons and form positive ions
(cations).
M - 1e- = M+
M - 2e- = M2+
M - 3e- = M3+
Example Mg - 2- = Mg2+ (called a magnesium ion).
- Atoms of non-metals (N) tend to gain one, two or three electrons and form negative
ions (anions).
N + 1e- = NN + 2e- = N2N + 3e- = N3Example
O + 2e- = O2- (called an oxide ion).
Note that anion names end in “-ide”.
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4b) Predicting Ionic Charges
▪ It is possible to predict the ion that an atom will form:
(i) Representative Elements - these are the Group A elements in the periodic table. The
ionic charge for these elements can be determined by their position in the periodic table.
1A
M+
2A
M2+
Let M = metal and N = non-metal
3A
4A
3+
transition
M
metals
5A
N3-
6A
N2-
7A
N-
8A
(ii) Transition Elements - most transition (Group B) elements are metals that can form
more than one positive ion.
iron = Fe2+ and Fe3+
lead = Pb2+ and Pb4+
Since it is difficult to predict these ions, you will be given a list of them to work with (see
“Common Ions” sheet).
Two methods of naming these ions are used:
(a) Stock System - a Roman numeral is used in brackets after the name
to indicate the charge.
Pb2+ = lead(II)
(b) Classical System - the classical name of the element is used with the
endings “-ous” (for the lower charge) and “-ic” (for the higher
charge).
If
the
symbol
is
Fe2+
Fe3+
Pb2+
Pb4+
then
its
name
is
iron(II)
iron(III)
lead(II)
lead(IV)
and its
classical
name
was
ferrous
ferric
plumbous
plumbic
▪ Try This Predict the ion that will be formed from these atoms:
1) phosphorus
6) lead(IV)
11) chromium(III)
2) aluminum
7) oxygen
12) calcium
3) cobalt(III)
8) magnesium
13) sulfur
4) nitrogen
9) lithium
14) potassium
5) mercury(II)
10) chlorine
15) tin(IV)
4c) Polyatomic Ions
♠ Poly = many
♠ Polyatomic ions are tightly bound groups of atoms that behave as a unit and carry a
charge.
Sulfate Ion (SO42-)
One sulfur atom and four
oxygen atoms forming one
strongly bonded unit with
two extra electrons.
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♠ There are many of these ions. Since we cannot predict their charges, we will be given a
list of them to work with (see “Common Ions” sheet).
♠ Most polyatomic groups contain oxygen and their names end in “-ate”.
An “-ite” ending means one less oxygen than the “-ate” ending.
sulfate = SO42sulfite = SO32-
chlorate = ClO3chlorite = ClO2-
♠ A few poyatomic groups do not contain oxygen. Their names end
in “-ide”.
cyanide = CN5. Formulas
a) Creating Formulas
♥ Ionic compounds are created when one kind of positive cation joins with one kind of
negative anion. The positive ion is always written first.
♥ The sum of the ionic charges in the formula must be zero. This is done by adjusting the
numbers of each ion.
Li+ + O2- → Li2O
(+1) (-2) (sum=0)
Note that the formula for the compound (Li2O) shows no charge. It is neutral overall.
♥How to Balance Ionic Charges
- When the charges are opposite numbers, their sum will automatically equal zero.
Ca2+ + O2- → CaO
(+2) (-2) (sum=0)
- When the charges are different, the number of the ions must be adjusted until the
sum of the charges becomes zero.
Al3+ + Cl- → AlCl3
(+3) (-1) (sum=0)
One convenient way of balancing charges is to cross-exchange the digits.
Al3+ + O2- → Al3+ 2O2- 3
♥Note that when you want to show more than one polyatomic group, you must put
brackets around the group.
Ca2+ + NO3- → Ca(NO3)2
(+2)
(-1)
(sum=0)
5b) Nomenclature
♦ Let M = Metal N = Nonmetal P = Polyatomic Group H = Hydrogen
(1) Ionic Compounds - always begin with a metal
- use the “Common Ions” sheet.
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M
+
N
=
M
+
P
=
M (stock number)
N -ide
if necessary
M ( ) P -ate (or -ite)
NaCl = sodium chloride
FeClO3 = iron(III) chlorate
(2) Molecular Compounds - no metals present
- use prefixes (not the “Common Ions” sheet).
N1 +
N2
=
prefix N1
prefix
except mono
N2 -ide
1 = mono
2 = di
3 = tri
4 = tetra
5 = penta
6 = hexa
7 = hepta
8 = octa
9 = nona
10 = deca
CO = carbon monoxide
P2O5 = diphosphorus pentoxide
(3) Acids - always begin with hydrogen
- there are two types.
a) Binary Acids contain only two elements
H
+
N
=
hydro
N -ic
acid .
HCl = hydrochloric acid
b) Oxo-acids contain oxygen in a polyatomic group
ite
ous
H + P
ate = P
ic
acid .
H2SO4 = sulfuric acid
H2SO3 = sulfurous acid
Nomenclature Summary
Formula begins with
Metal
Ionic substance
Use the
“Common Ions”
Sheet
Nonmetal
Hydrogen
Molecular Substance
Use Prefixes
Acid
binary
name begins
with “hydro-“
oxo“ate-ic
ite-ous”
acids
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6. Reactions
a) Balancing Equations
♦ A chemical equation is an expression for a chemical reaction. It is a quantitative
statement indicating the number of moles of each reactant and of each product.
Reactants → Products
♦ In chemical equations, matter must be conserved. The number of atoms of each kind
on the reactant side must equal those on the product side.
Consider this reaction: carbon + oxygen gas → carbon monoxide
C + O2 → CO
+
However,
2C + O2 → 2CO
+
→
(Matter is not conserved)
→
(Matter is conserved)
This equation is balanced.
♦ All equations must be balanced.
Balanced equations have 1. the chemical facts
2. correct formulas
3. atoms conserved
♦ Hints: (1) Try balancing O and H last
(2) Never change a formula to make the equation balance. Change coefficients, not
subscripts!
2CO2
coefficient
subscript
(3) Equations that will not balance probably contain an incorrect formula
All chemical reaction involves changes to substances. The starting chemicals (reactants)
are chemically changed into new chemicals (products)
Reactants → Products
Chemical equations must obey the Law of Conservation of Matter. Each side of the
chemical equation must contain equal amounts of each element.
Consider this equation:
Carbon + oxygen gas → carbon monoxide gas
→
CO
C
+ O2
→
Since the amounts of each element are not balanced, we use coefficients in front to the
formulas to make it balanced.
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2C
+ O2
→
→
2CO
Note: i. the coefficients tell us the amount of each chemical: 2 H2O
means 2 mol or 2 molecules of water
ii. Equations are balanced only wit coefficients.
DO NOT CHANGE THE FORMULAS!
Try This: Balance these equations
a) H2 + O2
→ H2O
b) Ca + O2
→ CaO
c) Mg(ClO3)2 → MgCl2 + O2
d) Ca + H2O
→ Ca(OH)2 + H2
e) HNO3 + Ba(OH)2 → Ba(NO3)2 + H2O
Symbol
+
→
=
⇔
(s)
↓
(l)
(aq)
(g)
↑
→

Pt
→
heat
Symbols used in Equations
Explanation
Used to separate two reactants or two products
“Yields”, separates reactants from products
An alternative to →
Used in place of a → for reversible reactions
Designates a reactant or product in the solid state; placed after the
formula
Alternative to (s); used only for a solid product (precipitate)
Designates a reactant or product in the liquid state; placed after the
formula
Designates an aqueous solution; the substance is dissolved in water
Designates a reactant or product in the gaseous state; placed after the
formula
Alternative to (g); used only for a gaseous product
Indicates that heat is supplied to the reaction
A formula written above or below the yield sign indicates its use as a
catalyst (in this example, platinum)
6 b) Types of Reactions
a) Combination Reactions
- Two or more reactants combine to form a single product. R
Note:
i. Two elements must form a binary compound.
S + O2 → SO2
H2 + Cl2 → 2HCl
ii.
Non-metal oxides in water form an acid.
SO3 + H2O
→ H2SO4
CO2 + H2O
→ H2CO3
+
S
→
RS
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Try This:
a) C + O2
→
b) aluminum + oxygen gas →
c) copper (I) + sulfur
→
d) N2O5 + H2O →
b) Decomposition Reactions (Opposite of Combination)
- A single reactant is broken down into two or more products: RS
→
R
+ S
decomposes
Note: i. Binary compound
2H2O →
2HgO →
into
→
element
+
element
2H2 + O2
2Hg + O2
ii. for more complicated compounds, try this:
If the compound contains
Example
C + O
NiCO3
H + O
NH4NO3
O
KClO3
NiCO3 heat
→

heat
NH4NO3 →

heat
2KClO3 →

One product might be
CO2
H2O
O2
NiO + CO2
N2O + 2H2O
2KCl + 3O2
Try This:
a) H2O →
b) mercury (II) oxide heat
→

c) magnesium carbonate heat
→

c) Single Replacement Reactions
- The atoms of one element can replace the atoms of a different element that is part of an
ionic compound.
For Metals:
M1 + M2X → M1X + M2
Note: that the first metal (M1) can only replace the second metal (M2) if it is more
reactive, as listed in the activity series of metals.
Examples:
i. Ca + 2NaOH → Ca(OH)2 + 2Na
ii. Mg + NaOH → no reaction (Mg is not more reactive)
For Non-Metals: (mostly the halogen gases) N1 + XN2 → XN1 + N2
Note: that the first non-metal (N1) can only replace the second non-metal (N2) if it is
more reactive.
Examples:
i. Cl2
+
BaBr2
→
BaCl2
+
Br2
ii. I2 + BaBr2
reactive)
Try This:
a) Fe + CuSO4 →
b) Mg + HCl →
c) Cl2 + KI →
→
no reaction (I is not more
d) Double Replacement Reactions
- This involves the reaction of two compounds to form two new
compounds.
RS + TU → RU + TS
2KI + Pb(NO3)2 → PbI2 + 2KNO3
Note:
i. These reactions involve ionic compounds or acids in aqueous
solution.
ii.
The two reactants just exchange their positive ions.
iii. Two compounds will not always react. There are three “driving
forces” which cause these reactions to occur:
a) a molecular compound can be formed
b) a gas can be formed
c) a precipitate can be formed
- One special type of a double replacement reaction is a neutralization
reaction:
water
Acid + Hydroxide → salt +
HCl +
NaOH
→ NaCl + H2O
Try This:
a) Barium chloride + sodium sulfate →
BaCl2
+
Na2SO4 →
b) Silver nitrate + sodium chromate →
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Activity Series of
the Elements
Metals Non-metals
Li
F
Rb
Cl
K
Br
Na
I
Sr
S
Ba
Ca
Mg
Al
MN
Zn
Cr
Fe
Cd
Co
Ni
Sn
Pb
H
Sb
Bi
As
Cu
Hg
Ag
Pt
Au
c) Hydrochloric acid + calcium hydroxide →
HCl
+
Ca(OH)2 →
Reactants
E1 + E2
C
E1 + C1
C1 + C2
SUMMARY
Let E= element and C = Compound
Reaction Type
Products
Combination
C
Decomposition
E1 + E2
Single replacement
E2 + C2
Double replacement C3 + C4
General Equation
R + S → RS
RS → R + S
T + RS → TS + R
RS + TU→RU + TS
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Other Equation Terms
a) Heat
- Exothermic reactions are accompanied by the release of heat into their surroundings.
CH4 + 2O2 → CO2 + 2H2O + energy
The release of energy is shown as a product in the equation.
We can show this in an energy level diagram:
- Endothermic reactions take place only when heat is continuously to the reactants.
CaCO3 + energy → CaO + CO2
The heat is much like a reactant – without it, the reaction cannot take place.
Remember that the unit for energy is the Joule (J).
b) Physical State
- Chemical reactions often depend on the physical state of the chemicals involved. This
information can be included in an equation by using these symbols:
(s) = solid (l) = liquid (g) = gas (a) = aqueous (dissolved in water)
For example, water may be H2O(s), H2O(l) or H2O(g).
- Double replacement reactions generally take place between ionic compounds in
aqueous solution: PbCl2(aq) + Na2CrO4(aq)
These two chemicals will now react if one of the products will form a precipitate that is
insoluble in water. PbCl2(aq) + Na2CrO4(aq) → PbCrO4(s) + 2NaCl(aq)
The precipitate PbCrO4(s) may now be removed from the solution by filtration.
- To write a balanced equation that includes physical states, then, we must be able to
predict if a chemical is soluble in water (aq) or insoluble in water (s). To do this, we will
use a solubility chart.
Note: “s” on a solubility chart = soluble and “(s)” in a equation = solid (insoluble)
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Try This: Write a balanced equation for each reaction in solution. Include the physical
state symbols.
1. Ba(OH)2 + Na3PO4 →
5. (NH4)S + CuCl2 →
2. AgNO3 + K2S →
6. NaOH + CuSO4 →
3. BaCl2 + Na3PO4 →
7. Na2S + Pb(NO3)2 →
4. K2CrO4 + Pb(NO3)2 →
8. PbBr2 + K2CrO4 →
Ionic Reactions
a) Aqueous Solutions
- The particles of a substance, when dissolving, separate from each other and disperse
throughout the solution.
+
water molecule (
)
- When NaCl is dissolved in water, the resulting solution contains ions:
NaCl(s) → Na+(aq) + Cl-(aq). The water simply causes the Na+ and Cl- ions to become
separated. The process of separating the ions in a solid is called dissociation.
- When a crystal contains more than one ion at a particular type, all of the ions will
dissociate.
CaCl2(s) → Ca2+(aq) + 2Cl-(aq)
Al2(SO4)3(s) → 2Al3+(aq) + 3SO42-(aq)
- The presence of ions in a solution permits the solution to conduct electricity; the more
ions present, the better the ability to conduct.
- Many soluble substances, such as HCl(g), are made up of molecules, not ions. These
substances, though, create ions when dissolved in water:
HCl(g) → H+(aq) + Cl-(aq)
This process of splitting a molecule into ions is called ionization.
- Precipitation is the reverse of the dissociation process.
Ag+(aq) + Cl-(aq) → AgCl(s)
b) Ionic Equations
A reaction between two chemicals dissolved in water is shown three ways:
i. Molecular Equation - the reaction is written as if all of the chemicals consist of
molecules:
BaCl2(aq) + Na2SO4(aq) → BaSO4(s) + 2 NaCl(aq)
ii. Total Ionic Equation - the equation is written to show that the dissolved
chemicals are actually in the forms of ions.
Ba2+(aq) + 2 Cl-(aq) + 2 Na+(aq) + SO42-(aq) → BaSO4(s) + 2 Na+(aq) + 2 Cl-(aq)
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iii. Net Ionic Equation - the equation is written to show only the ions or molecules
that are actively involved in the reaction.
Ba2+(aq) + SO42-(aq) → BaSO4(s)
The Na+ and Cl- ions were eliminated since they did not take part in the reaction. Ions
like this that occur in the same form on both sides of the equation are called spectator
ions.
Note: i. Ionic equation should be written from molecular equations.
ii. All solution equation must contain the physical state symbols.
iii. Ionic Equations are balanced in terms of mass and charge.
Try This: Write the total and net ionic equations for the following:
a) AgNO3 + K2CO3 →
b) Na3PO4 + MgSO4 →
c) SnSO4 + Na2S
→
d) FeBr2 + KOH
→
→
e) SnCl2 + K3PO4