Academic Chemistry
UNIT 7
PERIODIC TRENDS
& CHEMICAL
EQUATIONS
Name:
Class Period:
Test Date:
1
Chemistry Calendar
Monday
January 2nd
Tuesday
3rd
Wednesday
Thursday
Friday
4th
5th
6th
Notes #1: History
and Review of the
Periodic Table
Notes #2:
Definitions and
Trends
Graphing Activity
Teacher Prep Day
9th
Teacher Prep Day
10th
11th
12th
13th
QUIZ
Notes #3: Octet
Rule and
Valence e-‘s
Notes #4:
Bonding Review
and Intro to Rxns
Conservation of
Matter DEMO
17th
Notes #5: Rxns
(synthesis and
decomp)
Decomposition
DEMO
18th
Notes #5: Rxns
(Single and Double
Replacement)
CuCl2 + Al DEMO
Notes #5: Rxns
(RedOx and
Combustion)
Electrolysis Lab
19th
20th
Reactions Practice
Single
Replacement Rxn
Lab
Benchmark/
Test Review
Double
Replacement lab?
Benchmark/
Test Review
Double
Replacement lab?
TEST
16th
MLK Day
Student Holiday
2
NOTES #1: Textbook Activity
Use the textbook to fill in the notes about the following main ideas.
MAIN IDEA
NOTES

Dmitri Mendeleev



.
.
Henry Mosley
.
.
Periodic Law
.
.

His periodic table is arranged in order of increasing atomic number.

.

This is the way our current periodic table is still organized.
.
Metals
.
Moseley rearranged Mendeleev’s periodic table.


.


.
.
.
Alkaline Earth Metals
A periodic repetition of chemical and physical properties of the elements
when they are placed in order of increasing atomic number.
.
Row of the periodic table.
Periods have properties that change as you move across the row.
.
column of the periodic table
Groups have similar chemical and physical properties.
.



Most of the elements on the periodic table are metals.
Located to the left of the zigzag line.
All are solids at room temperature except Mercury (Hg). Gallium (Ga) is the
only metal that is a liquid at human body temperature.
.
.


.
.
.

Alkali Metals
.

.
Group
.
.
.
Period
Mendeleev organized elements according their chemical and physical
properties.
Arrangement allowed him to predict the existence and properties of three
missing elements (Ga, Sc, Ge)
Periodic: repeating pattern
Mendeleev’s periodic table is arranged in order of increasing atomic mass.
Traits: Soft, silvery metals that have a low density. They are the most
reactive metals and are not found by themselves in nature (they are always
in compounds). Explosive when exposed to air and water.
Uses: Used in soap, lights, and are present in the body (electrolytes).
Examples: Li, Na, K, Rb, Cs, Fr
.
.
 Traits: Harder, denser and stronger than the Alkali metals. Silvery in color.
They are the second most reactive metals and are not found by themselves
in nature (they are always in compounds).
 Uses: Gemstones, limestone, marble, alloys for aircraft, truck bodies,
fireworks and flares. Calcium needed for strong teeth and bones.
 Examples: Be, Mg, Ca, Sr, Ba, Ra
3
MAIN IDEA
NOTES
MAIN IDEA
NOTES

Transition Metals


.
.
Inner Transition Metals
.



.
.
Metalloids
.
.
Noble Gases
.
.
Nonmetals
.
.
Halogens
.

.
Traits: Shiny, silvery, reactive metals that have a high melting point.
Lanthanides are usually found in compounds. Actinides are all radioactive
(unstable). Many actinides are artificially produced. The only actinides
found in nature are Th, Pa, U and Np.
Uses: Lanthanides are used in tv tubes.
Examples: Ce, Pm, Eu, Am, U, Es
.


Traits: Properties of both metals and nonmetals. Semiconductors. Good
electrical conductors at high temperature, good insulators at low
temperature.
Uses: computer chips and glass
Examples: B, Si, Ge, As, Sb, Te, Po, At
.



Traits: No chemical reactivity (inert). All are gases.
Uses: Hot-air balloons
Examples: He, Ne, Ar, Kr, Xe, Rn
.

.
.
Traits: Poor electrical conductors but are good insulators. They do not have
luster and are brittle in the solid form. Non-ductile. Exist in various phases
(states) of matter: solids, liquids, or gases.
.
.
 Fluorine and chlorine are yellow-green gases. Bromine is a dark-red liquid.
Iodine is purple-black crystalline solid
 Most reactive group of nonmetals.
 Uses: gemstones, bones, teeth, DNA, medicines, organic compounds, Xerox
copying, bleach, antiseptic
4
MAIN IDEA
Transition Metals
Traits: Hard, silvery, solid metals that have a high density. High luster
(shiny), malleable, ductile and good conductors of heat and electricity. They
are less reactive than alkali and alkaline earth metals and many are found by
themselves in nature.
Uses: Paints, steel, coins, electrical wiring, plumbing, jewelry, construction.
Examples: Fe, W, Cu, Ag, Au, Pt, Ni
.
NOTES
.
NOTES #2: Definitions & Periodic Trends
WARM-UP:
1. As you move from left to right across the periodic table, the number of valence electrons
increases
Atomic mass definition:
the weighted average of the masses (protons + neutrons) of the isotopes of an element.
TREND:
Atomic mass increases left to right
REASON:
.
.
Atomic mass
Adding more protons and
neutrons!
increases
.
Atomic radius definition:
one half the distance between the nuclei in a molecule consisting of identical atoms.
TREND:
Atomic radii decreases from left to right
REASON:
.
.
Atomic radii
increases
.
Shielding Effect:
More protons to the right, increases
nuclear charge on electrons pulling them
closer to the nucleus. Therefore atomic
radii decreases.
Principal energy levels increase moving
down.
Electrons in outer shells are repelled by electrons occupying the inner shells. They are not as
strongly attracted to the positive nucleus as a result.
5
.
Ionization Energy definition:
the energy required to remove an electron from a gaseous atom. First ionization energy- energy required
to remove the outermost electron.
TREND:
Ionization energy increases left to right
.
.
Ionization energy
decreases
.
REASON:
Left to right – Nuclear charge
increases & shielding effect is constant
resulting in greater attraction of the
nucleus for the electron.
Moving down – Size of atoms increase so
the outermost electron is further from
the nucleus resulting in less attractive
forces.
Electronegativity definition:
Tendency for the atoms of an element to attract electrons when they are chemically combined with
atoms of another element
TREND:
Electronegativity increases left to right
REASON:
.
Electronegativity
decreases
.
.
Left to right – the octet rule. More
likely to attract electrons when
combined if closer to 8.
Moving down – valence electrons get
farther away from the nucleus, so
there are less attractive forces because
of distance and shielding effect.
Reactivity definition:
Reactivity refers to how likely or vigorously an atom is to react with other substances. This is usually
determined by ionization
energy and electronegativity because it is the transfer/interaction of
electrons that is the basis of chemical reactions.
Trend for nonmetals:
6
Reactivity increases
Reactivity decreases
Reactivity decreases
Reactivity increases
Trend for metals:
ASSESSMENT: Periodic Trends
Matching (#1-4):
D 1. electronegativity
B 2. shielding effect
C 3. ionization energy
A 4. atomic radius
A. The distance from the nucleus to the outermost electron of an atom.
B. The reduction of the attractive force between a positively charged
nucleus and its outermost electrons due to the cancellation of some of
the positive charge by the negative charge of the inner electrons.
C. The amount of energy needed to remove an outer electron from a
specific atom or ion in its ground state (in the gas phase).
D. The tendency of atom to attract electrons to itself when bonded to
another atom.
Which has the larger radius?
I
Which has a LOWER ionization energy?
I or Xe
As 11.
As or Se
Si 6.
Si or P
Al 12.
B or Al
Br 7.
F or Br
S 13.
S or Cl
P 8.
P or S
Be 14.
Be or B
Tl 9.
Al or Tl
At 15.
O or At
K 16.
K or Kr
5.
Ar 10. Ne or Ar
Which is MORE reactive?
Which has the HIGHER electronegativity?
Fr 17.
Fr or K
Ge 22.
Ga or Ge
Ra 18.
Be or Ra
Br 23.
Se or Br
F 19.
I or F
Se 24.
As or Se
O 20.
O or Se
N 25.
N or P
He or Ne
O 26.
O or I
K 27.
K or Cs
N/A 21.
7
ACTIVITY: GRAPHING PERIODIC TRENDS
PRE-LAB DISCUSSION:
The Periodic Table is arranged according to the Periodic Law. The Periodic Law states that when elements are arranged
in order of increasing atomic number, their physical and chemical properties show a periodic pattern. Students can
discover these patterns by examining the changes in properties of elements on the Periodic Table. The properties that
will be examined in this lesson are: atomic radius AND first ionization energy.
PURPOSE: To understand trends of the periodic table and practice methods of graphing.
PROCEDURE: Graph the following information according to the steps described.
Symbol
Period 2
Li
Be
B
C
N
O
F
Ne
Period 3
Na
Mg
Al
Si
P
S
Cl
Ar
Period 4
K
Ca
Rb
Sr
Cs
Ba
Atomic Number
Atomic Radius
(Picometers)
First Ionization
(Energy-Joules)
3
4
5
6
7
8
9
10
1.23
0.89
0.80
0.77
0.70
0.66
0.64
0.67
124
215
191
260
335
314
402
497
11
12
13
14
15
16
17
18
1.57
1.36
1.25
1.17
1.10
1.04
0.99
0.98
119
176
138
188
242
239
299
363
19
20
37
38
55
56
2.03
1.74
2.16
1.91
2.35
1.98
100
141
96
131
90
120
CALCULATIONS AND GRAPHS:
DIRECTIONS: You will construct FOUR graphs, please read through the directions for each before beginning.
Graph 1
For elements 3-20 make a graph of atomic radius as a function of atomic number. Plot atomic number on the X
axis and atomic radius on the Y axis. After creating the graph, use a colored pen or pencil to draw a vertical
line that represents that beginning of each period (horizontal row on the periodic table).
8
Graph 2
For elements in Group 1 (Alkali metals), make a graph of atomic radius as a function of atomic number. Make a
second line on this same graph that will represent Group 2 (Alkaline Earth Metals). Use a periodic table to
determine which elements are members of Group 1 and which elements are members of Group 2.
Graph 3
For elements 3-20, make a graph of the energy required to remove the easiest electron (first ionization energy)
as a function of atomic number. Plot atomic number on the X axis and energy required on the Y axis.
After creating the graph, use a colored pen or pencil to draw a vertical line that represents that beginning of
each period (horizontal row on the periodic table).
Graph 4
For elements of Group 1 (Alkali metals), make a graph of the energy required to remove the easiest electron
(first ionization energy) as a function of atomic number. On the same graph make a second line to represent
Group 2 (Alkaline Earth Metals). Use a periodic table to determine which elements are members of
Group 1 and which elements are members of Group 2.
QUESTIONS FOR DISCUSSION:
1. What happens to the atomic radius as the atomic number increases across a period? Down a group?
2. What happens to the energy needed to remove an electron as the atomic number increases across a period?
Down a group?
3. Why does atomic radius change as it does?
4. Why does the energy required to remove an electron change as it does?
9
Graph 1
GRAPH 2
10
Graph 3
Graph 4
11
NOTES #3: The Octet Rule and Valence Electrons
WARM-UP:
1. How many valence electrons are in the outer shell of each of the following:
a. Cl - 7
c. Al - 3
b. P – 5
d. Mg - 2
The Octet Rule:
1. All atoms want 8 valence electrons (except for H and He – they want 2 valence electrons.) Atoms will gain, lose
or share electrons in order to get to 8 valence electrons.
o The noble gases in group 8A (or 18) already have 8 electrons in the outer shell and are the most stable.
2. Electrons have a NEGATIVE charge.
o If an atom loses electrons, it becomes more positive. These atoms are called Cations.
o If an atom gains electrons it becomes more negative. These atoms are called Anions.
3. Valence electrons can be found by using the periodic table
1. How many valence electrons does Ca have? 2 What ion does it form? Ca2+
2. How many valence electrons does Sr have? 2
What ion does it form? Sr2+
3. How many valence electrons does Li have? 1
What ion does it form? Li+
4. How many valence electrons does Si have? 4
What ion does it form? Si4-
5. How many valence electrons does Cl have? 7
What ion does it form? Cl-
6. How many valence electrons does H have? 1
What ion does it form? H+
***Do not forget that if it ends in “ate” or “ite” it is a polyatomic ion and IS NOT found on the periodic table***
12
Bohr Models:

Electrons orbit in orbitals or more precisely – clouds. But for simplicity purposes, scientists draw electrons in
shells.
•
1st shell can fit up to 2 electrons. Think about the electron configurations. This first shell is 1s.
1st Period/Row
•
2nd and 3rd shells can each fit up to 8 electrons. These are the 2s and 2p shells or the 3s and 3p shells,
respectively.
2nd Period/Row
3rd Period/Row

The diagrams above are Bohr models. The show all of an atom’s electrons, but can take a while to draw –
especially for larger atoms. Chemists prefer to draw Lewis Dot Structures, which only show valence electrons
and the outer shell and are therefore quicker to draw and easier to read.
Lewis Dot Structures Review:

EXAMPLES:
Directions:
1. Start by writing the atomic symbol.
2. Determine the number of VALENCE electrons.
3. Draw dots for each of the valence electrons on each
of the four sides of the symbol.
a. Put one dot on each side before you start to pair
the electrons
Practice: Draw the Lewis Dot Structures for the following atoms
1.
He
3. Mg
2.
Cl
4. N
13
Notes #4: Bonding Review and Intro to Reactions
In a chemical reaction one or more bonds are formed, broken and reformed to create new substances. The original
substances are referred to as the reactants and the resulting substances are the products.
Reactants

Products
In the following examples, see if you can identify the reactants and the products as we review the types of bonds that
can be formed.

Ionic bond: The donating/receiving of electrons to form a bond. This bond is between metals and nonmetals
or metals and polyatomic ions.
Na + Cl

→
Na Cl
Sodium DONATES one electron to
chlorine leaving both atoms with full
outer shells thus the octet rule is fulfilled.
Covalent bond: The sharing of electrons to form a bond. This bond forms between nonmetals and nonmetals.
Cl + Cl
→
One Chlorine atom SHARES an electron
with the second Chlorine atom thus
fulfilling the octet rule.
Cl Cl
Practice: Identify the following types of bonds (Ionic or Covalent)
1.
CCl4
2.
Covalent
7. MgO
Ionic
MgCl Ionic
8. CH4
covalent
3. Na2NO3 Ionic
9. CaS
ionic
4. C2H6O2 Covalent
10. HCl
covalent
5. H2O
11. LiBr
ionic
12. K2SO3
ionic
Covalent
6. NH4Cl Covalent
14
Chemical Reactions
Symbols:
+

(s)
(l)
(g)
(aq)

Pt
Explanation
with, and; separates reactants/products
yields, gives, produces, reacts to form
solid state
liquid state
gas state
aqueous state; dissolved in water
heat used in reaction
catalyst in reaction (ex. Platinum)
Reversible reaction
Skeletal equations show no relative amounts of reactants or products.
1.
Mg (s) + H2SO4 (l)  MgSO4 (aq) + H2
Word equations explain reactions using names.
2.
Solid tetraphosphorus decoxide reacts with water to produce phosphoric acid.
Balancing Chemical Reactions
1.
Count the number of each element for both sides of arrow.
2. Add (coefficient) whole number at beginning of each compound to balance elements.
3. Each element must have same number on both sides of arrow. All in smallest ratio.
4. If number is a subscript, multiply by coefficient to tell how many of each.
5. There are 7 diatomic molecules which when not combined with another element, written as below.
H2
N2
O2
F2
Cl2
Br2
I2
6. When writing formulas from names, apply rules such as crisscross.
Guided Practice: Balance these equations:

1.
1 F2 + 2 KCl
2.
2 Al +
3 CuSO4
3.
Aluminum
+
Oxygen 
3 O2
->
4 Al
+
2 KF +

1 Cl2 **How many of each element do we have? Are both sides equal?
1 Al2(SO4)3 +
3 Cu
Aluminum oxide
2 Al2O3
15
**Write out the equation using nomenclature and
criss-cross and then balance.
Quick Nomenclature Review:
 Ionic Compounds: Name the metal unchanged, Name the nonmetal with “ide” on the end.
 Covalent Compounds: Use the prefixes listed below to name the compounds
1*
Mono

2
Di-
3
Tri-
4
Tetra-
5
Penta-
6
Hexa-
7
Hepta-
8
Octa-
9
Nona-
10
Deca-
Compounds with Polyatomic Ions: Write both the metal and the polyatomic ions AS IS .
Acids: Use the chart below
ION TYPE
Polyatomic
Monatomic
ION ENDING
-ite
-ate
-ide
ACID NAME BEGINNING
NO hydro- beginning
NO hydro- beginning
hydro- beginning
ACID ENDING
-ous
-ic
-ic
Independent Practice:
1. 2 C2H6 + 7 O2  4 CO2 + 6 H2O
2.
Magnesium Nitrate + Calcium Hydroxide  Calcium Nitrate + Magnesium Hydroxide **use polyatomic Ion list.
Mg (NO3)2 + Ca(OH)2 → Mg(OH)2 + Ca(NO3)2
Balanced
3. _____ H2 + _____ Cl2  2 HCl
4. Magnesium Chloride + Sodium Iodide  Magnesium Iodide + Sodium Chloride
MgCl2 + 2 NaI → MgI2 + 2 NaCl
5. 2 HgO  2 Hg + _____ O2
6. Aluminum Carbide + Sodium Oxide  Aluminum Oxide + Sodium Carbide
Al4C3 + 6 Na2O → 2 Al2O3 + 3 Na4C
7. _____ Ca + 2 H2O  _____ Ca(OH)2 + _____ H2
8. _____ CH4 + 2 O2  _____ CO2 + 2 H2O
9. _____ Na2O2 + _____ H2SO4  _____ Na2SO4 + _____ H2O2
balanced
10. Copper (II) Carbonate  Copper (II) Oxide + Carbon Dioxide
CuCO3 → CuO + CO2
Balanced
11. Silicon dioxide + Hydrofluoric Acid  Silicon Tetrafluoride + water
SiO2 + 4 HF → SiF4 +2 H2O
12. Sulfuric Acid  Sulfur Trioxide + Water
H2SO4 → SO3 + H20
16
Balanced
Notes #5: Reaction Types
COMBINATION (synthesis) Reactions

Two or more reactants form a SINGLE product. Use crisscross rules to form compounds.
1. 4 Al (s) +
3 O2 (g) 
2 Al2O3
2. 2 S (s) +
3 O2 (g) 
2 SO3
3. 2 Cu (s) +
_____ Cu (s)
_____ S (s)

+
_____ S (s)
4. 2 Fe (s) +
5. 4 Fe (s) +
_____ O2 (g)
3 O2 (g) 
_____ Cu2S *

_____ CuS *

2 FeO *
2 Fe2O3 *
*Elements with multiple charges can form two different compounds depending if limited or in excess.

Reactions between WATER and NONMETAL OXIDES usually give an ACID. (H + polyatomic)
6. _____ SO3 (g) + _____H2O (l)  _____H2SO4
Balanced
7. _____ N2O5 (g) + _____H2O (l)  2 HNO3

Reactions between WATER and METAL OXIDES usually give a BASE. (Metal + OH)
8. _____ CaO (s) + _____H2O (l) 
_____Ca(OH)2 Balanced
9. _____ Na2O (s) + _____H2O (l) 
2 Na(OH)
10. _____MgO (s) + _____H2O (l) 
_____Mg(OH)2 Balanced
DECOMPOSITION Reactions

Single reactant broken down into TWO OR MORE simpler products.

1. CaCO3 (s)

CaO
+
CO2
+
O2
4 Ag
+
O2
H2
+
O2
2.
2 H2O (l) + electricity  2 H2
3.
2 Ag2O (s)
4.
H2O2 (l)


Metal Carbonate decomposes to METAL OXIDE and CO2.
5.
Nickel (II) carbonate
NiCO3 → NiO + CO2

Nickel(II)oxide + _____
Balanced
17
Metal Hydroxide decomposes to METAL OXIDE and H2O.
6.
Calcium hydroxide

Calcium Oxide and Water
Ca(OH)2 → CaO + H2O Balanced

Metal Chlorate decomposes to METAL CHLORIDE and O2.
7.
Potassium chlorate (s)

Potassium Chloride + oxygen gas
2 KClO3 → 2 KCl + 3 O2

Some Acids decompose to NONMETAL OXIDE & H2O.
8. Carbonic Acid →
H2CO3 → CO2 + H2O
________________________________ + H2O
Balanced
SINGLE REPLACEMENT Reactions

An element in reactant replaces a* similar element.
o A + BC

B + AC
A = Metal
o D + BC

C + BD
D = Nonmetal
1.
_____ Zn (s) +_____ H2SO4 (aq) _____ ZnSO4 + _____ H2
2. _____ K (s) +
2 H2O (l)
_____ KO2 + 2 H2
3.
_____ Sn (s) + _____ NaNO3 (aq) 
4.
_____ Cl2 (g) + 2 NaBr (aq)  2 NaCl + _____ Br2
5.
2 Al (s) +
6.
_____ KCl +_____ I2  No Reaction
No Reaction
3 H2SO4 (aq) _____ Al2(SO4)3 + 3 H2
7. Would copper replace chromium in a reaction? No
8. Would calcium replace mercury in a reaction? YES
9. What element would platinum be able to replace in a reaction.
Gold
Activity Series
Metal
Lithium
Potassium
Barium
Calcium
Sodium
Magnesium
Aluminum
Manganese
Zinc
Chromium
Iron
Cobalt
Nickel
Tin
Lead
(Hydrogen)
Copper
Mercury
Silver
Platinum
Gold
You can use the activity
series to predict whether
or not certain reactions
will occur. A specific
metal can replace any
metal listed below it that
is in a compound. It
cannot replace any metal
listed above it. For
example, copper atoms
replace silver atoms in a
solution of silver nitrate.
However, if you place a
silver wire in aqueous
copper (II) nitrate, the
silver atoms will not
replace the copper. Silver
is listed below copper in
the activity series, so no
reaction occurs. The
letters NR (no reaction)
are commonly used to
indicate that a reaction
will not occur.
**This applies to the
nonmetal halogens as
well. They increase as
you move up the family.
18
DOUBLE REPLACEMENT Reactions

Exchange of positive ions between two compounds.
o A+B- + C+D- 
A+D- + C+Bo Products are usually a precipitate, gas or molecular compound such as water.

Formation of a gas:
1.

_____ FeS (s) + 2 HCl (aq)
_____ FeCl2 +_____ H2S
Acid-Base Reactions
o When acids and bases react, they generally produce Salt and water. The common salt used to flavor foods
such as french fries or scrambled eggs is sodium chloride, NaCl, and is a product of some acid-base
reactions. However, there are many other salts used in chemistry that are produced from acid-base
reactions.
o
Examples:
HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)
acid
base
salt
water
H2SO4 (aq) + 2KOH (aq) → K2SO4 (aq) + 2H2O (l)
acid
base
salt
water
2.

HCl (aq) + Ca(OH)2 (aq) → CaCl2 (aq) + H2O (l)
3. _____ H2SO4 (aq) + 2 NaOH (aq)
_____ Na2SO4 +2 H2O
4. 2 HCl (aq) + _____ Mg(OH)2 (aq)
 _____ MgCl2 + 2 H2O
Precipitation Reactions: When two solutions of ionic compounds are mixed, a product that can form is an insoluble
salt called a precipitate. The precipitate is a solid that falls out of solution and is indicated as a solid on the product
side of a chemical equation. We use a solubility chart to determine what compounds will combine to form a
precipitate.
o
Example:
AgNO3 (aq) + NaCl (aq) → AgCl (s) + NaNO3 (aq)
Precipitate
5.
_____ Na2SO4(aq) + _____ Ba(NO3)2(aq) → _____ BaSO4(s) + 2 NaNO3(aq)
6.
_____ K2CO3 (aq) + _____ Sr(NO3)2 (aq) → _____ SrCO3 (s) + 2 KNO3 (aq)
7. _____ BaCl2 (aq) _____ + K2CO3 (aq) _____ BaCO3 + 2 KCl
8. _____ SrBr2 (aq)_____ + (NH4)2 CO3 (aq)_____ SrCO3 +2 NH4Br
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COMBUSTION Reactions

Oxygen reacts with hydrocarbon to produce CARBON DIOXIDE and WATER for complete combustion.
(Incomplete combustion results when NOT enough oxygen is available)
1. 2 C6H6 + 15 O2  12 CO2 + 6 H2O
2.
2 C8H18 + 25 O2  16 CO2 + 18 H2O
3. _____ C6H12O6 + 6 O2  6 CO2 + 6 H2O
4. 2 C10H22 + 31 O2  20 CO2 + 22 H2O
OXIDATION-REDUCTION (Redox) Reactions

Electrons are transferred from one atom to another through the change in oxidation numbers.
Oxidation
Reduction
o
*OIL RIG = ox is loss, red is gain
Loss of eGain of e-

Rules:
o
o
o
o

Assign Charges (oxidation #) to all atoms in equation.
Identify atoms that are oxidized & atoms reduced.
Determine the change in charge (oxidation #) for atoms that are oxidized or reduced.
Draw line connecting atoms oxidized or reduced & write the net change in oxidation #.
Oxidation numbers are a convenient way of identifying redox reactions and also indicating which element is
oxidized and which is reduced. Here's an example - the reaction between sodium metal and chlorine gas:
2 Na + Cl2 → 2 NaCl

It is often useful to write the oxidation number for every element, in every compound, above the element in the
equation. Thus for our reaction we have:
0
2 Na
o

0
+
Cl2
+1 -1
→
2 NaCl
Be sure to note that the balancing coefficients in the equation (the "2" in front of Na and in front of
NaCl) do not affect the value of the oxidation numbers. We'll return to these coefficients soon.
Since no electrons are transferred in acid-base reactions or precipitation reactions, acid-base reactions and
precipitation reactions are not redox reactions.
20


A chart is a useful way for us to summarize the changes in oxidation number for each element:
element
Initial ox. #
Na
0
Cl
0
Final ox #
change in electrons (e-)
oxidized or reduced
→
+1
lost 1 e-
oxidized
→
-1
gain 1 e-
reduced
We see several important things in our table 


Since oxidation numbers did change, this was a redox reaction
Na's oxidation number increased - from 0 on the reactant side to +1 on the product side. An element
becomes more positive by losing electrons. **Loss of electrons is Oxidation (LEO)
Cl's oxidation number decreased, from 0 to -1, as chlorine gained electrons. **Gain of electrons is
Reduction (GER)
Practice: For the following redox reactions, write what elements are oxidized and reduced:
1.
2.
3.
2 Fe
+

3 Cl2
element
Initial ox. #
Fe
0
Cl
0
2 KBr +

Cl2
2 FeCl3
Final ox #
change in electrons (e-)
oxidized or reduced
→
+1
lost 1 e-
oxidized
→
-1
gain 1 e-
reduced
Final ox #
change in electrons (e-)
oxidized or reduced
2 KCl + Br2
element
Initial ox. #
K
+1
→
+1
No change
No change
Br
-1
→
0
Lost 1 e-
Oxidized
Cl
0
→
-1
Gained 1 e-
Reduced
Final ox #
change in electrons (e-)
oxidized or reduced
N2 (g) +
3 H2 (g)

2 NH3 (g)
element
Initial ox. #
N
0
→
-3
Gained 3 e-
Reduced
H
0
→
+1
Lost 1 e-
oxidized
21
4. 2 Mg (s) +

O2 (g)
2 MgO (s)
element
Initial ox. #
Mg
0
O
0
Final ox #
change in electrons (e-)
oxidized or reduced
→
2+
Lost 2 e-
Oxidized
→
2-
Gained 2 e-
reduced
Reducing agent: loss = oxidized
Oxidizing agent: gain = reduced
Academic Chemistry Unit 7 Test Review
1. How is the modern periodic table arranged?
2. Rows on the periodic table are called
. Columns are called
.
3. Be able to identify Metals, Nonmetals, and Metalloids on the periodic table.
4. Know the names for groups 1, 2, 3-12, 17 and 18.
5. Atoms across a period are arranged according to
.
6. What are two ways in which atoms within a group are similar to one another?
7. How did Mendeleev arrange the periodic table?
8. How did Mosley arrange the periodic table?
9. Name 3 elements in the same period as nitrogen:
10. Name 3 elements in the same group as oxygen:
11. What is an anion?
12. Are anions larger or smaller than their neutral atoms?
13. What is a cation?
14. Are cations larger or smaller than their neutral atoms?
22
15. What is the octet rule?
16. Elements in group 2A tend to form ions with what charge?
17. Elements in group 7A tend to form ions with what charge?
Describe the following trends across a period and down a group:
Electronegativity Explain Reason for Trend
Definition:
Label & Describe the trends
First Ionization
Energy
Definition:
Explain Reason for Trend
Label & Describe the trends
Atomic Radii
Definition:
Explain Reason for Trend
Label & Describe the trends
23
Ionic Radii
Definition:
Explain Reason for Trend
Label & Describe the trends
18. ____ NaBr + ____ H3PO4  ____ Na3PO4 + ____ HBr
Type of rxn: ______________________
19. ____ C2H4 + ____ O2  ____ CO2 + ____ H2O
Type of rxn: ______________________
20. ____ Mg + ____ Fe2O3  ____ Fe + ____ MgO
Type of rxn: ______________________
21. ____ PbSO4  ____ PbSO3 + ____ O2
Type of rxn: ______________________
22. ____ H2SO4 + ____ NH4OH  ____ H2O + ____ (NH4)2SO4
Type of rxn: ______________________
23. __HCl (aq) + __NaOH (aq) → __NaCl (aq) + __H2O (l)
Type of rxn: ______________________
24. __Pb(NO3)2 (aq) + __H2SO4 (aq) → __PbSO4 (s) + __HNO3 (aq)
Type of rxn: ______________________
25. __HNO3 (aq) + __KOH (aq) → __KNO3 (aq) + __H2O (l)
Type of rxn: ______________________
26. __Al (s) + __HCl (aq) → __AlCl3 (aq) + __H2 (g)
Type of rxn: ______________________
27. __I2O5 (s) + __CO (g) → __I2 (s) + __CO2 (g)
Type of rxn: ______________________
28. __H2SO4 (aq) + __KOH (aq) → __H2O (l) + __K2SO4 (aq)
Type of rxn: ______________________
29. __Cu(s) + __HNO3(aq) → __Cu(NO3)2(aq) + __NO2(g) + __H2O(l)
Type of rxn: ______________________
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30. __Na3PO4 (aq) + __FeCl3 (aq) → __NaCl (aq) + __FePO4 (s)
Type of rxn: ______________________
31. __Pb(C2H3O2)2 (aq) + __HCl (aq) → __PbCl2 (s) + __HC2H3O2 (aq)
Type of rxn: ______________________
32. __Bi(OH)3(s) + __Na2SnO2(aq) → __Bi(s) + __Na2SnO3(aq) + __H2O(l)
Type of rxn: ______________________
33. __HNO3 (aq) + __Mg(OH)2 (aq) → __H2O (l) + __Mg(NO3)2 (aq)
Type of rxn: ______________________
34. __Cl2 (g) + __H2O (l) → __HCl (aq) + __HClO (aq)
Type of rxn: ______________________
35. __Na3PO4 (aq) + __Pb(NO3)2 (aq) → __NaNO3 (aq) + __Pb3(PO4)2 (s)
Type of rxn: ______________________
36. __NaOH (aq) + __CaCl2 (aq) → __NaCl (aq) + __Ca(OH)2 (s)
Type of rxn: ______________________
37. __Na2CO3 (aq) + __Ca(NO3)2 (aq) → __CaCO3 (s) + __NaNO3 (aq)
Type of rxn: ______________________
38. __S (s) + __HNO3 (aq) → __SO2 (g) + __NO (g) + __H2O(l)
Type of rxn: ______________________
39. __AgNO3 (aq) + __H2S (aq) → __Ag2S (s) + __HNO3 (aq)
Type of rxn: ______________________
40. __HCl (aq) + __LiOH (aq) → __H2O (l) + __LiCl (aq)
Type of rxn: ______________________
41. __KMnO4(aq) + __HCl(aq) → __MnCl2(aq) + __Cl2(g) + __H2O(l) + KCl(aq) Type of rxn: ______________________
42. __(NH4)2S (aq) + __Co(NO3)2 (aq) → __CoS (s) + __NH4NO3 (aq)
Type of rxn: ______________________
43. __H3PO4 (aq) + __Ca(OH)2 (aq) → __H2O (l) + __Ca3(PO4)2 (aq)
Type of rxn: ______________________
25