NAME:
LAB PARTNER:
TYPES OF CHEMICAL REACTIONS (60 points)
Objectives:



Perform and make observations of different types of chemical reactions
Qualitatively identify products of these chemical reactions
Write balanced chemical equations for these chemical reactions
Background:
There are an infinite number of chemical reactions. Chemist have divided these into broad
classifications based on certain criteria. The most important classifications are; synthesis,
combustion, decomposition, single replacement, double replacement, neutralization and redox.
Note that some reactions will fall into more than one classification. For example, all single
replacement reactions are also redox reactions.
Reactions are described with chemical equations. A chemical equation is the symbolic
representation of the chemicals involved in the reaction. These chemical equations can be used
to describe both physical processes and chemical reactions. All chemical equations consist of
reactants, the starting materials, products, the ending materials, the state of matter that the
materials are in and a reaction arrow representing that a reaction has occurred.
reactants
products
The states of matter are symbolized by subscripts which follow each chemical. The most
common states of matter are solid, liquid, gas and aqueous. The first three should be selfexplanatory the final, aqueous, occurs when a substance is dissolved in water. Gasses and ions
are commonly found in an aqueous state. The symbols are as follows; solid (s), liquid (l), gas (g)
and aqueous (aq).
Here are two examples of chemical equations:
H2O (s)  H2O (l)
CH4 (g) + O2 (g)  CO2 (g) + H2O (g)
The first is the physical process of ice melting the second is the combustion of methane gas.
There is one other aspect necessary in writing a proper chemical equation, the equation must
have the same number of each type atom on each side of the reaction arrow. This is the Law of
Conservation of Matter, in other words you are not allowed to create nor destroy matter. The first
of the two above equations is fine, two hydrogens on each side and one oxygen on each side.
The second equation needs to be balanced. This can be accomplished by placing a 2 in front of
both the O2(g) and the H2O(g).
OVER
CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (g)
On each side we find one carbon, four hydrogens and four oxygens. This equation obeys the
conservation of matter and is said to be balanced.
Occasionally you may see something written above or below the arrow. This is normally a
catalyst, platinum and nickel are a common metal catalysts. Another possibility is that the
something is an environmental necessity, for example, heat or ultra violet light.
2 CO (g) + O2 (g)
Pt

2 CO2 (g)
The above chemical equation occurs in an automobiles catalytic converter, converting carbon
monoxide into carbon dioxide in the presence of a platinum catalyst. The platinum is the reason
the catalytic converter is expensive.
Synthesis reactions, two or more reactants combine to form one product. The general equation
is:
A+BC
A specific example is the tarnishing of a silver tea set:

 Ag2S (s)
Ag (s) + S (s) 
Decomposition reactions, one reactant breaks to form two or more products. The general
equation is:
A  B +C
A specific example is the electrolysis of water, electricity is causes this reaction:
2 H2O (l)
e


2 H2 (g) + O2 (g)
Single replacement reactions, one chemical replace another in a compound. The general
equation is:
A + BX  B + AX
A specific example is zinc metal replacing iron in the iron (III) oxide compound, ships use this
reaction to keep their hulls from rusting.
3 Zn (s) + Fe2O3 (s)

2 Fe (s) + 3 ZnO (s)
The reactivity of a substance depends on its ability to gain or lose electrons. It is possible to
arrange the elements into a series based upon their reactivity. Such a list is called an activity
series.
Activity Series for Metals and Non-metals
Single Replacement Reactions
Name
Symbol
Decreasing activity
↓
Metals
Lithium
Sodium
Potassium
Rubidium
Barium
Strontium
Calcium
Li +
Na+
K+
Rb+
Ba+2
Sr+2
Ca+2
Magnesium
Aluminum
Manganese
Zinc
Chromium
Iron
Cadmium
Cobalt
Nickel
Tin
Lead
Mg+2
Al+3
Mn
Zn+2
Cr+3
Fe
Cd+2
Co+2
Ni+2
Sn
Pb
Hydrogen
H2
Antimony
Bismuth
Copper
Mercury
Silver
Platinum
Gold
Sb
Bi
Cu
Hg
Ag+1
Pt
Au
reacts with water and
acids
reacts with acids
replacing hydrogen
fairly unreactive
Non-metals
Fluorine
Chlorine
Bromine
Iodine
F2
Cl2
Br2
I2
OVER
While there are many types of single replacement reactions, this lab will focus on two primary
types. In one type, a more active metal, or halogen, replaces a less active metal or halogen from
solution. An example of this is:
Zn (s) + CuSO4 (aq)

Cu (s) + ZnSO4 (aq)
The other type of single replacement reaction involves the replacement of hydrogen from an acid
by a metal. Consider the reaction between zinc and hydrochloric acid:
Zn (s) + 2 HCl (aq)

H2 (g) + ZnCl2 (aq)
The zinc metal is active enough to replace the hydrogen from the acid. Bubbles of hydrogen gas
can be seen rising to the surface, and the piece of zinc is consumed. On the other hand, if the
less reactive metal, copper, is placed into a hydrochloric acid solution, no reaction will take
place.
Double replacement reactions, chemicals in each of two compounds switch compounds. The
general equation is:
AB + CD  AD + CB
A specific example is allows for the removal of toxic barium from a water source by adding a
compound containing sulfate.
BaCl2 (aq) + Na2SO4 (aq)  BaSO4 (s) + 2 NaCl (aq)
Certain substances are not very soluble in water. Frequently such substances are generated in a
reaction vessel by the addition of various other substances that are themselves soluble. For
example, silver chloride is not soluble in water. If an aqueous solution of silver nitrate (very
soluble) is mixed with an aqueous solution of sodium chloride (very soluble), the combination of
silver ions from one solution and the chloride ions from the other solution produces silver
chloride, which then forms a precipitate that settles to the to the bottom of the container. The
solution that remains above the precipitate of silver chloride is a solution of sodium nitrate. The
silver ions and the sodium ions have switched partners:
AgNO3 (aq) + NaCl (aq)  AgCl (s) + NaNO3 (aq)
The silver ion and the sodium ion have replaced each other in this process; this sort of reaction is
referred to as a double replacement reaction.
To clarify what is really happening in such reactions, it is often more instructive to write the
reaction in its net ionic form. In a net ionic equation for a precipitation reaction, only the ions
involved in actually forming the precipitate are shown; the other ions contained in the original
reagents used are called spectator ions and are still present in the solution after the precipitate
has been formed. For example, the net ionic reaction for the previous reaction is:
Ag+ (aq) + Cl- (aq)  AgCl (s)
The net ionic reaction is especially instructive because it implies that any solution containing
silver ions should react with any solution containing chloride ions since it is these ionic species
that are really reacting. For example, if dilute hydrochloric acid, HCl, were added to a solution
of silver nitrate, a precipitate would be expected to form. The reactions of silver nitrate with
sodium chloride and of silver nitrate and hydrochloric acid are exactly the same when the net
ionic reaction is considered.
Neutralization reactions occur when the hydrogen ion from an aqueous acid combines with the
hydroxide ion from an aqueous base to produce water and an aqueous salt. For example,
HNO3 (aq) + NaOH (aq)  H2O (l) + NaNO3 (aq)
The net ionic equation for neutralization reactions is the same and is typical of the reaction
between acids and bases in aqueous solution:
H+ (aq) + OH- (aq)  H2O (l)
Predicting Products
The products of a chemical reaction may often be predicted by applying known facts about
common reaction types. While there are hundreds of different ”kinds” of chemical reactions,
only four general types of reactions will be considered; single displacement, double
displacement, decomposition, and synthesis. The following sections give examples of these
general types.
Synthesis
In a synthesis reaction two or more simple substances (compounds and/or elements) are
combined to form one new and more complex substance. Here the general form is
Element + element  compound
or
compound + compound  compound
A + X  AX
The following are some general types of synthesis reactions.
1. Combination of elements.
Fe(s) + S(1)  FeS(s)
2Na(s) + Cl2(g)  2NaCl(s)
OVER
2. Combination of an acid anhydride with water to give an acid.
SO2(g) + H2O(1)  H2SO3(aq)
N2O3(g) + H2O(1)  2HNO2(aq)
CO2(g) + H2O(1) H2CO3(aq)
P2O5(s) + 3H2O(l) 2H3PO4(aq)
3. Combination of a basic anhydride or a metallic oxide with water to form a base.
Na2O(s) + H2O(1)  2NaOH(aq)
CaO(s) + H2O(1)  Ca(OH)2(aq)
BaO(s) + H2O(1)  Ba(OH)2(aq)
4. Combination of the metal of a basic oxide with the nonmetal of an acidic oxide to form a salt.
CO2(g) + Na2O(s)  Na2CO3(s)
P2O5(s) + 3BaO(s)  Ba3(PO4)2(s)
SO2(g) + MgO(s)  MgSO3(s)
Decomposition
When energy in the form of heat, electricity, light, or mechanical shock is supplied, a
compound may decompose to form simpler compounds and/or elements. The general form for
this type of reaction is
compound  two or more substances
AX  A + X
The following are some general types of decomposition reactions.
1. If some acids are heated, they decompose to form water and an acidic oxide.
H2SO3(aq)  SO2(g) + H2O(1)
H2CO3(aq)  CO2(g) + H2O(1)
2. When some metallic hydroxides are heated, they decompose to form a metallic oxide and
water.
Ca(OH)2(s)  CaO(s) + H2O(g)
2Fe(OH)3(s)  Fe2O3(s) + 3H2O(g)
3. Some metallic carbonates decompose to form a metallic oxide and carbon dioxide when
heated.
Li2CO3(s)  Li2O(s) + CO2(g)
CaCO3(s)  CaO(s) + CO2(g)
4. Metallic chlorates decompose to form metallic chlorides and oxygen when heated.
2KC103(s)  2KC1(s) +302(g)
Ni(C103)2(s)  NiC12,(s) + 302(g)
5. Most metallic oxides are stable, but a few decompose when heated.
2HgO(s)  2Hg(1) + O2,(g)
2Ag2O(s)  4Ag(s) + O2(g)
6. Some compounds cannot be decomposed by heat, but can be decomposed into their elements
by electricity.
2H2O(1)  2H2(g) + O2(g)
2NaC1(1) 2Na(1) + Cl2,(g)
MgCl2(1)  Mg(1) + Cl2(g)
Single Displacement
One metallic element displaces another metallic element in a compound, or a nonmetallic
element displaces another nonmetallic element in a compound. A single displacement has the
general form
element + compound  element + compound
A+BX  AX+B
X+BY  BX+Y
The following are some general types of single displacement reactions.
1. An active metal will displace the metallic ion in a compound of a less active metal.
Fe(s) + Cu(NO3)2(aq)  Fe(NO3)2(aq) + Cu(s)
Pb(s) + 2AgC2H3O2(aq)  Pb(C2H3O2)2,(aq) + 2Ag(s)
2. Some active metals such as sodium and calcium will react with water to give a metallic
hydroxide and hydrogen gas.
2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)
Ca(s) + 2H2O(1)  Ca(OH)2(aq) + H2(g)
3. Active metals such as zinc, iron, and aluminum will displace the hydrogen in acids to give
a salt and hydrogen gas.
Zn(s) + 2HC1(aq)  ZnCl2(aq) + H2(g)
Fe(s) + H2SO4(aq)  FeSO4(aq) + H2(g)
OVER
4. Halogens (which are active nonmetals) will displace less active halogens.
Cl2(g) + 2NaBr(aq)  2NaC1(aq) + Br2(aq)
Br2(g) + 2KI(aq)  2KBr(aq) + I2(g)
For example, use the Activity Chart to either predict the products or determine that a reaction
does not take place:
(a) magnesium is added to a solution of iron (III) chloride:
First write the formulas and physical states of the reactants:
Mg (s) + FeCl3 (aq)
Then predict the products based on the Activity Chart:
Mg (s) + FeCl3 (aq)  MgCl2 + Fe
Balance the equation:
3 Mg (s) + 2 FeCl3 (aq)  3 MgCl2 + 2 Fe
Insert the physical states:
3 Mg (s) + 2 FeCl3 (aq)  3 MgCl2 (aq) + 2 Fe (s)
Double Displacement
The positive and negative ions of two compounds are interchanged. The form of these
reactions is easy to recognize,
compound + compound  compound + compound
AX + BY  AY + BX
The following are some general types of double displacement reactions.
1. A reaction between an acid and a base yields a salt and water. Such a reaction is a
neutralization reaction.
2KOH(aq) + H2SO4(aq)  K2SO4 (aq) +H2O(l)
Ca(OH)2(aq) + 2HNO3(aq)  Ca(NO3)2(aq) + 2H2O(1)
2. Reaction of a salt with an acid forms a salt of the acid and a second acid which is volatile.
2KNO3(aq) + H2SO4(aq)  K2SO4(aq) + 2HNO3(g)
FeS(c) + 2HC1(aq)  FeCI2(aq) + H2S(g)
2NaC1(aq) + H2SO4(aq)  Na2SO4(aq) + 2HCl(g)
2a. This same reaction of a salt with an acid or base may yield a compound which can be
decomposed.
CaCO3(aq) + 2HC1(aq)  CaCl2(aq) + H2CO3(aq)
H2CO3(aq)  CO2(g) + H2O(1)
or
CaCO3(s) + 2HC1(aq)  CaCl2(aq) + CO 2(g) + H2O(1)
K2SO3(aq) + 2HNO3(aq)  2KNO3(aq) + SO2(g) + H2O(1)
NH4C1(aq) + NaOH(aq)  NaC1(aq) + NH2(g) + H2O(1)
3. Reaction of some soluble salts produces an insoluble salt and a soluble salt.
AgNO3(aq) + NaC1(aq)  AgC1(s) + NaNO3(aq)
Na2SO4(aq) + Ba(NO3)2(aq)  BaSO4(s) + 2NaNO3(aq)
CuSO4(aq) + Na2S(aq)  CuS(s) + Na2SO4(aq)
Ca(C2H3O2)2(aq) + (NH4)2CO3(aq)  CaCO3(s) + 2NH4C2H3O2(aq)
The solid precipitate can be predicted using the solubility rules/chart:
Solubility Chart
I - Insoluble, S - Soluble, SS - slightly soluble
Acetate
Bromide
Carbonate
Chlorate
Chloride
Chromate
Cyanide
Fluoride
Hydroxide
Iodide
Nitrate
Oxide
Phosphate
Silicate
Sulfate
Sulfide
Sulfite
Ag
I
I
I
S
I
I
I
S
I
I
S
I
I
I
SS
I
I
Al
S
S
I
S
S
S
S
I
S
S
I
I
I
S
I
I
Ba
S
S
I
S
S
I
S
I
S
S
S
S
I
I
I
I
I
Bi Ca
S S
S
I I
S
S S
S S
I I
S I
I SS
S S
S S
I I
I I
I I
S SS
I I
I I
Cd
S
S
I
S
S
S
S
S
I
S
S
I
I
I
S
I
I
Co
S
S
I
S
S
S
I
S
I
S
S
I
I
I
S
I
I
Cr
S
S
I
S
S
S
I
S
I
S
S
I
I
I
S
I
I
Cu
S
S
I
S
S
S
I
I
I
S
S
I
I
I
S
I
I
Fe
S
S
I
S
S
S
I
I
I
S
S
I
I
I
S
I
I
H
S
S
S
S
S
S
S
S
S
S
S
S
S
S
S
S
S
Hg
S
I
I
S
I
I
S
OVER
I
I
S
I
I
I
I
I
I
K Mg Mn Na NH4
S S S S S
S S S S S
S I
I S S
S S S S S
S S S S S
S S S S S
S S
S S
S I
I S S
S I
I S S
S S S S S
S S S S S
S I
I S SS
S I
I S S
S I
I S S
S S S S S
S I
I S S
S I
I S S
Ni
S
S
I
S
S
S
I
I
I
S
S
I
I
I
S
I
I
Pb
S
I
I
S
I
I
I
I
I
I
S
I
I
I
I
I
I
Sn
S
S
I
S
S
S
S
I
S
I
I
I
S
I
I
Sr
S
S
I
S
S
I
S
I
SS
S
S
S
I
I
I
I
I
Zn
S
S
I
S
S
S
I
I
I
S
S
I
I
I
S
I
I
General Solubility Trends:








All compounds of the ammonium ion (NH4+), and of the Alkali metal (Group IA) cations, are
soluble.
All nitrates and actetates are soluble.
All chlorides, bromides, and iodides are soluble EXCEPT those of silver, lead, and mercury(I).
All sulfates are soluble EXCEPT those of silver, lead, mercury(I), barium, strontium, and
calcium.
All carbonates, sulfites, and phosphates are insoluble EXCEPT those of ammonium and alkali
metal (Group I) cations.
All hydroxides are insoluble EXCEPT those of ammonium, barium, and alkali metal (Group I)
cations.
All sulfides are insoluble EXCEPT those of ammonium, Alkali metal (Group I) cations, and
Alkali earth metal (Group II) cations.
All oxides are insoluble EXCEPT those of calcium, barium, and Alkali metal (Group I) cations;
these soluble ones actually react with the water to form hydroxides.
For example:
Predict the products of a reaction between aqueous rubidium phosphate and aqueous
titanium (IV) nitrate.
First write the formulas of the reactants:
Rb3PO4 (aq) + Ti(NO3)4 (aq)
Then predict the products based on the cations trading places with one another:
Rb3PO4 (aq) + Ti(NO3)4 (aq)  Ti3(PO4)4 + RbNO3
Balance the equation:
4 Rb3PO4 (aq) + 3 Ti(NO3)4 (aq)  Ti3(PO4)4 + 12 RbNO3
Finally, based on the solubility rules, insert the physical states of the products. The rules state:
“All sulfates are soluble EXCEPT those of silver, lead, mercury(I), barium, strontium, and calcium.”
And “All nitrates and actetates are soluble.” Therefore, the Ti3(PO4)4 compound will be a solid and the
RbNO3 will be aqueous.
4 Rb3PO4 (aq) + 3 Ti(NO3)4 (aq)  Ti3(PO4)4 (s) + 12 RbNO3 (aq)
Materials
1 M NaCl
1 M NaOH
1 M NaNO3
6 M HCl
2 M HCl
.1 M Cu(NO3)2
.1 M Fe(NO3)3
.1 M AgNO3
6 M NH4OH
CuCO3
Mossy zinc
Copper wool
Magnesium ribbon
Steel wool
Equipment
Pipets, thin stem
Distilled water and wash bottle
Balance
Paper towels
Crucible tongs
Test tubes
Wood splints
Evaporating dish
Beakers
Reaction plate, 48-well plate
Toothpicks
Bunsen burner
Microspatula
Test tube holder
Test tube rack
Matches
Safety Precautions
Copper(II) nitrate and silver nitrate solutions are slightly toxic by ingestion. Silver nitrate and
sodium hydroxide solutions are skin and eye irritants; silver nitrate will also stain skin and
clothing. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles.
Wash hand thoroughly if you contact any of the above chemicals.
Wash hands thoroughly with soap and water before leaving the lab.
OVER
Laboratory Procedure:
For each reaction you perform, make observations and predict the products formed. Below is a
section titled results, write the balanced chemical equations for these reactions here. If no
observable reaction occurred write no reaction or NR.
This lab will use a flaming or smoldering piece of wood as a qualitative test for either CO2(g) or
O2(g). If the fire goes out, you have CO2 being generated, if the fire flames up, you have O2 being
generated.
1. With your lab partner, go to a chemical reaction station.
2. Record, in your data table, your observations of the appearance of the reactants before you
perform the reaction.
3. Follow the instructions at each station to complete the chemical reaction.
4. Record your observations both during and after the reaction.
5. Clean all equipment and dispose of the chemicals properly.
6. Move to another station and repeat steps (1) – (5).
Station #1:
Magnesium metal and oxygen gas
a. Acquire a piece of magnesium metal approximately 2 cm long, scrub it will steel wool to
return its luster.
b. Light the Bunsen burner.
c. Using crucible tongs, hold the magnesium ribbon in the burner flame until it ignites.
d. Once ignited, place it in the evaporating dish to burn.
e. Note and record its appearance.
Station #2:
Hydrogen Chloride and Ammonia
a. In a fume hood, add 5 drops of 6 M HCl, hydrochloric acid, to a test tube.
b. In a fume hood, add 5 drops of 6 M NH4OH, ammonium hydroxide, to a test tube.
c. Remove the test tubes from the hood. Bring them together so the test tube mouths are as close
as possible to each other and tilted toward each other.
d. Record any observations.
e. Return the test tubes to the “used test tube beaker” in the fume hood
Station #3:
Copper (II) Carbonate
a. Measure as close to 1.000 gram of CuCO3, copper(II) carbonate onto a piece of paper and
place it into a clean, dry test tube.
b. Heat the test tube and contents gently, moving the test tube in and out of the flame from the
top of the solid to the bottom of the solid.
c. Using the glowing splint test, check periodically for the presence of an evolving gas as you
heat the test tube.
d. After a couple of minutes of gentle heating, increase the flame and heat vigorously for no
more than 10 minutes.
e. You should not quit heating until you get a negative test for the presence of this gas.
Allow five minutes of cooling before you remove the test tube and mass it along with its
contents.
Station #4:
Silver Nitrate and Copper
a. Put 10 drops of 0.10 M AgNO3, silver nitrate into a test tube.
b. Insert a short piece of copper wool into the test tube.
c. Allow the reaction to take place for at least 30 minutes.
d. Check up on this several times during the 30 minutes and record ALL changes.
Station #5:
HCl and Zinc
a. Add 5ml of 6M HCl to a clean dry test tube.
b. Affix the test tube to a ring stand.
c. Catch a wood splint on fire.
d. Add a small piece of Zn metal to the test tube, wait 10 seconds.
e. Carefully insert the burning splint into the test tube, expect a sound and pop.
f. Make observations of this second test tube.
Station #6:
Copper Nitrate and Zinc
a. Add 10 drops of 0.10 M Cu(NO3)2, copper(II) nitrate to a test tube.
b. Add a small piece of zinc metal into the test tube.
c. Allow the reaction to take place for at least 30 minutes.
d. Check up on this several times during the 30 minutes and record ALL changes.
OVER
Station #7: Well-Plate Reactions:
Place the well plate on a paper towel and number the well plate wells as follows:
Number one set of wells 1 – 3, then skip a well and number the next set 4 – 6, then skip a well
and number the last set 7 – 9.
1. Add 4 drops of 1 M NaOH, sodium hydroxide to each of wells 1 - 3.
a. Add 4 drops of 0.10 M Cu(NO3)2, copper(II)nitrate, to well 1.
b. Add 4 drops of 0.10 M AgNO3, silver nitrate to well 2.
c. Add 4 drops of 0.10 M Fe(NO3)3, iron(III) nitrate, to well 3.
Agitate each test tube to mix the contents and record the behavior of each mixture.
2. Add 4 drops of 1 M NaCl, sodium chloride, to each of wells 4 - 6.
a. Add 4 drops of 0.10 M Cu(NO3)2, copper(II)nitrate, to well 4.
b. Add 4 drops of 0.10 M AgNO3, silver nitrate to well 5.
c. Add 4 drops of 0.10 M Fe(NO3)3, iron(III) nitrate, to well 6.
Agitate each test tube to mix the contents and record the behavior of each mixture.
3. Put 4 drops of 1 M NaNO3, sodium nitrate, to each of wells 7 - 9.
a. Add 4 drops of 0.10 M Cu(NO3)2, copper(II)nitrate, to well 7.
b. Add 4 drops of 0.10 M AgNO3, silver nitrate to well 8.
c. Add 4 drops of 0.10 M Fe(NO3)3, iron(III) nitrate, to well 9.
Agitate each test tube to mix the contents and record the behavior of each mixture.
Rinse out the used well plate in the sink and place it in the white tub next to the drying
oven.
DATA – OBSERVATIONS – RESULTS (3 points per station/well plate well = 45 points)
Reactant Formulas:
Observations:
Balanced
Chemical
Equation:
Station Reactant Formulas:
Observations:
Station
1
Word
Equation:
2
Word
Equation:
Balanced
Chemical
Equation:
Station Reactant Formulas:
Observations:
3
Word
Equation:
Balanced
Chemical
Equation:
Station Reactant Formulas:
Observations:
4
Word
Equation:
Balanced
Chemical
Equation:
OVER
Reactant Formulas:
Observations:
Balanced
Chemical
Equation:
Station Reactant Formulas:
Observations:
Station
5
Word
Equation:
6
Word
Equation:
Balanced
Chemical
Equation:
Station 7 Reactant Formulas:
Well #1
Observations:
Word
Equation:
Balanced
Chemical
Equation:
Station 7 Reactant Formulas:
Well #2
Word
Equation:
Balanced
Chemical
Equation:
Observations:
Station 7 Reactant Formulas:
Well #3
Observations:
Word
Equation:
Balanced
Chemical
Equation:
Station 7 Reactant Formulas:
Well #4
Observations:
Word
Equation:
Balanced
Chemical
Equation:
Station 7 Reactant Formulas:
Well #5
Observations:
Word
Equation:
Balanced
Chemical
Equation:
Station 7 Reactant Formulas:
Well #6
Observations:
Word
Equation:
Balanced
Chemical
Equation:
OVER
Station 7 Reactant Formulas:
Well #7
Observations:
Word
Equation:
Balanced
Chemical
Equation:
Station 7 Reactant Formulas:
Well #8
Observations:
Word
Equation:
Balanced
Chemical
Equation:
Station 7 Reactant Formulas:
Well #9
Observations:
Word
Equation:
Balanced
Chemical
Equation:
Concluding Questions: Classification of Chemical Reactions (1 point each = 15 points)
1. In the following complete ionic equation, which ions are spectator ions? How do you know?
Cu2+(aq) + 2NO3-(aq) + 2Na+(aq) + S2-(aq)  CuS(s) + 2NO3-(aq) + 2Na+(aq)
2. Suppose you had an unknown solution that contained either dissolved NaCl, or NaNO3. You
added AgNO3 to the unknown solution and observed that no solid forms. What can be
concluded about the unknown solution? (use the solubility rules)
3. Suppose you tried to carry out a double-replacement reaction by mixing together equal
volumes of a solution that contained dissolved NaF and a solution that contained dissolved
NaCl. Would you expect a reaction, if so write the balanced chemical equation?
4. In the single replacement reaction between zinc and HCl, what reaction resulted from
inserting the lit splint into the test tube mouth?
5. Using the Activity Series, predict whether a reaction will occur for the following
combinations. If a reaction occurs, finish writing the balanced equation; if not, then indicate
as such with the letters NR.
(a) Zn (s) + NaCl (aq) →
(b) Cu (s) + KI (aq) →
(c) Ca (s) + NiCl2 (aq) →
(d) Cl2 (g) + 2 NaF (aq) →
(e) Zn (s) + HNO3 (aq) →
OVER
6.
Using the solubility table/rules, predict whether a reaction will occur. If a reaction occurs,
complete the balanced equation; if not, then indicate as such with the letters NR.
(a) 2 KI (aq) + Pb(NO3)2 (aq) →
(b) 2 NH4Cl (aq) + CaS (aq) →
(c) AgNO3 (aq) + NaOH (aq) →
(d) CuSO4 (aq) + Ba(NO3)2 (aq) →
(e) K2CO3 (aq) + 2 LiOH (aq) →
7. Let’s say you were doing an experiment on two different chemicals in two different test
tubes. One chemical is giving off oxygen gas or hydrogen gas the other is giving off carbon
dioxide gas. What one test could you do in each test tube that would prove which is giving
of the oxygen or hydrogen and which is giving off the carbon dioxide?