.3 - l
Experiment 3
Observation of Reactions in Aqueous Solutions
Objectives
1 . To observe and describe the changes which occur when aqueous solutions of two different
substances are mixed and a chemical reaction (change) occurs.
2.
To analyze the observations and, where
similarly.
possibJe~ to
classify them into groups which behave
3. To write a general statement for each group in objective 2 which describes the change characteristic
of the group.
4. To apply the general statements in objective 3 t.o predict the observable changes for other
experiments in which aqueous solutions of two different substances are mixed.
Discussion
A Evidence of a Chemical Reaction when Solutions are
1'fi..~ed
Aqueous solutions are prepared by dissolving a substance (called the solute) in water (called the
solvent). When two aqueous solutions are mixed, it is possible that a chemical reaction involving
the two solutes may occur. The observable evidence for a chemical change may be:
1. A change in the color of the solution.
2. The formation of an insoluble solid substance (precipitate. ppt.).
3. The appearance of bubbles (effervescence) due to the formation of an insoluble gaseous
substance.
4. A change in temperature. The mixture may be warmer (exothermic change) or cooler
(endothermic change) than the original solutions.
5. A change in the odor given off by the solutions.
B. Making Observations
There are three important periods of time when making observations of a chemical reaction.
1. Observe and record the characteristic properties of each reactant before mixing.
2. Observe the changes as they occur when the reactants are mixed.
3. Observe and record the characteristic properties of each product once all chemical change has
ceased.
When making observations with respect to solutions or precipitates, there are two common
misconceptions to bear in mind.
1. A true solution will transmit light and therefore appear clear. This is not to be confused with
colorless. Some solutions are colored but never-the-less clear.
2. Depending on particle size and density a precipitate may appear as a cloudiness or as a solid
which quickly settles to the bottom.
For example, milk appears white but is not clear and therefore contains a precipitate; the solid
particles are so small that they do not settle to the bottom. White solutions are not possible
since most white solutes dissolve to form a colorless solution.
C.
1\nalysis of Observations
To analyze your observations in terms of the compounds involved in the reactions, consider the
following:
1. All compounds used as reactants in this experiment are water soluble.
2. The types of compounds used in this experiment can be described as composed of a positive
part and a negative part (i.e., positive and negative ions).
3. To understand the aqueous solution reactions in this experiment the positive and negative
parts of each compound should be considered separately.
4. Positive parts of the solutes are either hydrogen, ammonium or a metal; they are the first part
of the formula or name. Consider the following examples with the positive parts circled. The
other part of each formula is the negative part.
~oride
~hydroxide
@DH
The name and formula of each reactant will be provided with each reaction which you carry out.
In doing your analysis, mentally split each formula into two parts as described above and look
over all your observations for general rules which relate part$ mixed together with observed
reaction results.
Procedure
1. Examine the solutions of the substances listed by name in the second column and by formula in the
third column of the observation table. Record the color of each solution in the fifth column.
2. .For each numbered reaction (first column) mix together 1 mL portions of each reagent using small
test tubes. For the reagent, phenolphthalein, use only 1 drop. In reaction 25 use only 1 drop of
ammonia. Stir each mi.nure thoroughly with a clean glass stirring rod.
3. E.icamine each reaction mixture and record your observations in the sixth column.
4. Wash glassware well (final rinse with distilled water) before using again.
Note: The label on the reagent bottle containing ammonia, NH3 , as a solute may
actually say ammonium hydroxide, NH40H.
Glossary of Terms for this Experiment
Given below are some useful definitions of terms that may be used to describe reactions,.reactan.ts and
products in this experiment.
Descriptive Terms
precipitate: a solid which appears in a solution when two solutions are mixed. It is the result of a
chemical reaction, when a solution of sulfuric acid is added to a solution of barium chloride, a
precipitate appears. Precipitates are described as flocculent, milky, creamy or heavy.
flocculent: describes a precipitate which has the appearance of woonfoating in the liquid, e.g. a
precipitate of aluminum hydroxide is flocculent.
milkv: describes a liquid with a white precipitate which gives the liquid the appearance of milk. The
precipitate is very light, when carbon dioxide is passed into lime water, a light precipitate of calcium
carbonate turns the lime water into a milky liquid.
creamv: describes a white precipitate which is heavier than the precipitate forming a milky mixture;
the precipitate still floats in the liquid, e.g. silver chloride forms a creamy precipitate.
heavv: describes a precipitate which sinks to the bottom of the liquid, e.g. barium sulfate forms a
heavy precipitate.
clear: describes a liquid which is transparent, e.g. water is a clear liquid. A clear liquid can be colored
or colorless. e.g. tea is a clear, brown liquid; kerosene is a clear, colorless liquid.
translucent: describes a solid object, material or substance which allows light to pass through it, but a
person can not see clearly through it. For example, waxed paper is translucent, but not transparent;
milk is a translucent liquid.
opaoue: describes an object, material or substance which does not allow light to pass through it, e.g.
leather and paper are opaque; mercury is an opaque liquid.
colored: describes a material or substance which has color, but without naming the color; e.g. a colored
solution can be brown, blue, green, black, etc. A material or substance can be described as white or
colored~ e.g. milk is a white liquid and lead (II) sulfide is formed as a black precipitate which is
considered to be a colored precipitate.
colorless: describes a material or substance which has no color, e.g. water is colorless; air is colorless.
Colorless is the opposite of colored. To contrast white and colorless: the paper of this handout is white;
glass in a window is colorless.
smoke: a system of solid particles suspended in a gaseous medium such as air.
effervescence: to give off bubbles of gas in solution.
Below are acceptable abbreviations for some of the terms defined above. You may use these
abbreviations when describing your observations.
Term
precipitate
flocculent
milh.7
creamy
clear colorless
Abbr.
ppt
Term
heavy
floe
clear
milky
translucent
opaque
crm
solution
ccs
Abbr.
hvy
clear
trans
opq
clear colorless gas
Term
smoke
efferescence
solution
gelatinous
Abbr.
smk
eff
soln
gel
ccg
Describing colors
Colors are described with the appropriate name, and modifying words that help describe the intensity,
and shading. For example, a solution might be described as a light sky blue, or another solution might
be described as a deep royal blue. Below are acceptable abbreviations for various colors that you might
observe during this experiment. You may use these abbreviations when describing your observations.
Term
red
Abbr.
red
orange
orng
Term
green
blue
Abbr.
grn
bl
Term
yellow
white
Abbr.
ylw
wht
Term
black
Abbr.
blk
.3 - 4
In ten ti on ally blank
lO*
9*
8*
7*
6*
5
4
3
2
1
H.x:n
No.
*Save the pmducts for flllther testing.
<Nfli)2COs
BaC12
barium chloride
ammonium carbonate
K2C0 3
BaCl 2
Nai~~0 3
BaCl2
potassium carbonate
barium chloride
sodium carbonate
barium chloride
(N1J.i)~04
BaCl2
baiium chlmide
ammonium sulfate
Na2S04
BaCl 2
barium chl01ide
sodium sulfate
HiS0 4
BaCl 2
barium chlmide
su1fmic acid
K 2C0 3
-
ammonium carbonate
barium chloride
polusHiwn carbonate
bmium chloride
sodium carbonate
barium chlmide
ammonium sulfate
barium chl01ide
sodium nulfate
barium chlotide
sulfwic acid
barium chloride
potassium carbonate
sulfuric acid
f!iS0 4
sulfwic acid
potassium carbonate
sodium carbonate
Na 2COs
suHlnic acid
Hp0 4
su]fmic acid
sodium carbonate
potassium carbonate
H.I~Oa
hydrochlmic acid
HCJ
hydrogen chloride
Name-·-··----·····
Observations of Solutions Before
l\lixing
_.. __ _
Obse1vations dming Heaction and of
Products after Reaction
Sta ti on Number ____ . _____ -· ···-- Date
potassium carbonate
sodium carbonate
hydmchlmic add
Names of Solutions
Na 2C0 3
HCl
Fo1mulas of
Solutes
sodium carbonate
hydrogen chloride
(Where appmpriate)
Names ofSo]ules
Expeiimen t 3
Observation of Reactions in Aqueous Solution
01
6:>
CQ
I
:o
hyd mgen chloride
HCJ
__ .,. .....
HCI
hydmgen chlotide
Rxn 10* products
Rxn 9* products
-----
Rxn 9* products
hydrochlotic acid
Rxn lO* pmducts
hydrochlmic acid
hydmchlmic acid
HCI
Rxn 8* products
-----
Rxn 8* products
hydmgen chl01ide
hydrochlmic acid
llCI
H.x.n 7* pruducts
-----
Rxn 7* products
hydrogen chlotide
hydrochloric acid
HCl
hydrogen chl01ide
H.xn 6* products
Names of Solutions
-----
Fonn ulas of
Solutes
H.xn 6* products
Names of Solutes
(Where apprnpriate)
Observations of Solutions Before
Mixing
Observations during Heaction and of
Prnduct.s after Reaction
19§
18§
17§
16§
hyd1uchlolic acid
HCI
KOH
hydrogen ch1otide
potassium hydroxide
hydmchlmic acid
Nlfs
HCl
NI\
H}30 4
ammonia
hydrogen chlotide
ammonia
sulfuric acid
sulfiuic acid
ammonia
ammonia
~'30,.
sulfwic acid
suHlnic acid
potassium hydroxide
sodium hydroxide
Na OH
sodium hydroxide
§For reactions 16 - 19 do the folJowing for each pair of reactants. First make observations on t.he fo1lowing solutions - sodium hydroxide,
potassium hydroxide, and ammonia, then add one drop of phenolphthalein to each of the solutions. Now make observations on the resulting
solution. Next make observations on each acid. then mix the reagents together. Finally make observations on the reaction and products.
15
14
13
12
lJ
Rxn
No.
Names of So1utes
(Where appropriate)
Fo1mulas of
Solutes
Names of Solutions
Observations of Solutions Before
lVIixing
26
25*
24
23
221'
2lt
20-j"
* Save the products for further testing.
Nl-Js
ammonia
NI~
ammonia (1 dmp)
·---·
CuS0 4
copper(II) sulfate
Rxn 25* products
KOH
CuS0 4
copper(!!) sulfate
potassium hydroxide
Na OH
(1
drop)
ammonia
Rxn 25* pl'oducts
ammonia
copper(Il) sulfate
potassium hydmxide
c·opper<Jl) sulfate
sodium hydroxide
copper(Il) sulfate
cuso ..
copper(ll) sulfate
sodium hydroxide
ammonia
NI-Ja
iron(III) chloride
ammonia
FeCl 3
KOH
potassium hydroxide
iron(III) chloride
imn(IlI) chloride
FeCl3
iron(fl I) chloride
potassium hydmxi<le
so<lhm1 hydroxide
imn<III) chloride
Na OH
FeCl 3
sodiwn hydmxide
iron(JI!) chloride
t ff no precipitate is evident in 20, 21 or 22, add more Na OH, KOf·L or Nf~ solution (as required).
No.
Hxn
Observations dming Heaction and of
Pmducts aft.er Reaction
-:1
(.!.J
.S-8
Intentionally blank
Exp. 3 “Reactions in Aqueous Solutions”
Analysis of Observations
Reaction
Numbers
1, 2, 3, and 4
Common Ions
Common Observation(s)
Cation
Anion
H+
CO 3 2−
bubbles appear when the two solutions
are mixed
3-10
Applications of Analysis
1. Predict and describe what you would observe if solutions containing the following solutes are
mixed. Hint: use your experience in the lab.
a. HN0 3 and Rb 2C0 3
b. BaBr2 and N a~C0 3
c. Ba(N03h and
K~0 4
d. Fe(Cl0 3) 3 and
N&
(Hint: NF:I.PfD
e. Ca(OH>2 and phenolphthalein
f. CuCl2 and LiOH
2. You are given a test tube which must contain 3 mL of either a) pure water, b) a soluble carbonate
salt, c) a soluble sulfate salt, or d) a mixture ofboth soluble carbonate and soluble sulfate salts.
First; Using what you have learned, develop a general test procedure to determine the presence or
absence of a sulfate or a carbonate ion in the solution.
Describe your test.
Second; Can your first test be used to predict if you have a) pure water? Yes
No
Explain your response.
Third; If the test tube does not have 1 a) pure water, how can you proceed to determine which ion or
ions you were given?
6-1
Experiment 6
Common Chemical Reactions
Objectives
1.
To observe and study common types of chemical reactions.
2.
To relate observed chemical reactions to chemical equation writing .
.3.
To develop a replacement or activity series for selected elements.
Discussion
By now you have had substantial opportunity to observe chemical reactions. The most easily detected
changes involve changes in color, odor, physical state or temperature as product forms. As a reaction occurs,
characteristics of reactants disappear, and the characteristics of one or more new substances appear. In
this experiment you will have a chance to observe many reactions and then explain what is happening in
terms of the chemicals involved. A chemical equation is written with the hope of connecting a chemical
equation on paper to a corresponding chemical. reaction in the laboratory. Ultimately you want t.o he able to
look at a chemical equation and make a substantial number of predictions as to what would be observed if
the reaction was actually cani.ed out in the laboratory.
Chemical equations must always be balanced, and the physical state of each substance must be shown.
Use (g) for gas, (J) for liquid, (s) for solid and (aq) for aqueous (i.e., dissolved in water).
Procedure
The solutions used in this experiment are 0.1 M except when otherwise indicated. When a reaction involves
mixing two solutions, the mixture should be stirred. Silver nitrate should be used with care: it stains
everything (skin, clothes, etc.).
A.
B.
Combination Type of Reaction (Demonstration)
1.
Hold a short piece of magnesium ribbon at one end with crucible tongs and ignite the other
end using the laboratory burner.
2.
Catch the ash from the above reaction on a watch glass. Crush the ash· on the watchglass
with a stirring rod. Add a few drops of distilled water and stir. Test the resulting mixture
with both red and blue litmus paper, then add a drop of phenolphthalein.
3.
Add a small amount of CaO (calcium oxide) to a few milliliters of distilled water in a test
tube. Test the resulting mixture with both red and blue litmus paper, then add a drop of
phenolphthalein.
4.
Hold a small piece of charcoal (carbon) with crucible tongs and burn it completely in a burner
flame. Heat another piece until it glows and then drop it into a wide-mouth bottle
containing oxygen.
5.
Open bottles of concentrated hydrochloric acid, HQ(aq), and concentrated ammonia, NH.a(aq).
Hold the mouths of the open bottles near enough to each other so that the gaseous solutes
from each solution will interact.
Decomposition Type of Reaction
1.
Fill the bottom curved suiface of a dry Pyrex test tube with copper(Il) sulfate pentahydrate.
Clamp the test tube at a 45° angle, (at the very top) and heat the crystals with a laboratory
burner flame. Note the color change and observe the inside wall of the test tube.
6-2
2.
Fill the bottom curved surface of a dry Pyrex test tube with copper(!!) carbonate. In a
small test tube place 10 ml of limewater, Ca(QH)iaq\ and 2 drops of phenolphthalein.
From the stockroom get a carbon dioxide test apparatus. Mount the Pyrex test tube at
about a 45° angle (See Figure 1 ). Connect the carbon dioxide test apparatus so that the end
of the delivery tube almost touches the bottcm of the limewa ter test tube as shown in Figure 1.
a)
Heat the copper(ll) carbonate gently at first, gradually increasing the temperature .
the limewater.
...L\ny gas formed will pass into
b)
Stop heating once a precipitate (cloudy) forms and a color change is observed in the
limewater test tube. (This reaction is not a decomposition reaction.)
CAL1TION - Before you stop heating the Pyrex test tube, remove the delivery tube from
the limewater to prevent sucking back the liquid into the hot Pyrex test tube.
Attach clamp here
at the top of the
test tube.
~
Figure 1: Setup for Carbon Dioxide Test
C.
Single Replacement Type of Reaction Use small test tubes.
1.
To 2 mL of copper(II) sulfate solution, CuS04(aq), add a piece of zinc.
2.
To 2 mL of zinc sulfate solution, ZnSOiaq) add a piece of copper.
3.
To 2 mL of silver nitrate solution, AgNOs(aq), add a piece of copper.
4.
To 2 mL of 6 M hydrochloric acid, HQ(aq), add a piece of copper.
5.
a)
To 2 mL of 6 M hydrochloric acid, HCl(aq), add a piece of zinc.
b)
Use a match to carefully ignite the gas produced during the reaction. (This reaction
is not a single replacement reaction.)
1
6-3
D.
Double Renlacement Tvpe of Reaction Use small test tubes.
1.
To 2 mL oflead<III nitrate, PbfNOsMaq), add 2 mL of 6 M hydrochloric acid, HCl( aq).
To 2 mL oflead(II) nitrate,
3.
E.
Pb<NOs.~(aq) 1
add 2 mL of sodium chromate, Na2Cr04(aq).
To 2 mL of zinc sulfate, ZnSO.,(aq), add 2 mL of 6 M hydrochloric acid, HCl( aq).
Other Important Reactions
1.
To 2 mL ofironfIII) chloride solution, FeCls(aq), add 2 mL of 6 M ammonia solution.
2.
To 0.1 gram of sodium carbonate, Na2C03(s), slowlv add 2 mL of 6 M hydrochloric acid,
HCl(aq.l.
6-4
Intentionally Blank
...
b
Experiment 6
Common Chemical Reactions
Station Number
Observations and Equations
A.
Combination Type of Reaction
l.
Magnesium is burned in air.
Equation:--------------------------·--------
2.
The ash from above is mixed with water and tested with indicators.
Observations: _ _ _ __
Equation:---------·------------
3.
Calcium oxide, CaO, is mixed with water and tested with indicators.
Observations: - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - Equation: - - - ·
4.
Carbon is burned in air (also in pure oxygen).
Observations:
Equation:·-------------------------------
5.
Hydrochloricacid, HCI, (aq), is placed near ammonia, :NHs, solution.
Observations:
B.
Decomposition Tvoe of Reaction
1.
CopperGD sulfate pentahydrate, CuSO.r5HzO, is heated.
i::::
-
tJ
6-6
2.
a)
Copper1Il) carbonate, CuC0 3, is heated.
Observations: _ _ _ _ __
b)
The gaseous product from (a) is passed into limewater, Ca(OHh
Observations:
Equation:_~---~---~---------------------~
(Not a decomposition reaction.)
C.
Single Replacement Type of Reaction
1.
Zinc is added to copper(Il) sulfate, CuS04 , solution.
Observations: - - - - - ·
Equation:
2.
Copperis added to a zinc sulfate, ZnS04, solution.
Observations:
Equation: - - · - - -..- - - - - · - - - - · - - - - - · - - - -
3.
Copper is added to a silver nitrate, AgNOs, solution.
Equation:
4.
Copper is added to hydrochloric acid, HCl(aqJ.
Observations: - - - - · - - - - - - - - - - - - - - - - - - - - - - - - - - - - -
5.
a)
Zinc is added to hydrochloric acid, HCI (aq).
b)
The gas produced by the reaction in (a) is ignited with a match.
Equation: _____________________________
(Not a single replacement reaction.)
~
6-7
D.
Double Replacement Tvpe of Reaction
1.
A solution oflead(U) nitrate, Pb(NOs~, is mixed with hydrochloric acid. HCl.
Observations: - - - - - - - - - - - - - - - - - - - - - - - - - - - - - Equation:-----------------
2.
Solutions oflead(II) nitrate, Pb(NOah and sodium chromate, N azCr().h are mixed.
Observations:_~~~----------
3.
Solutions of zinc sulfate, ZnS0 4, and hydrochloric acid, HCl (aq), are mixed.
Observations:
E.
Other Important Reactjon Tvoes
1.
Solutions of iron(lli) chloride~ FeCla, and ammonia, NHs, are mixed.
Observations: - - - - - - - - ·
Equation: __
2.
Hydrochloricacid, HCl(aq), is added to sodium carbonate, Na'.;COs(s).
Observations:
Equation: _ _ _ __
Exp. 6 “Common Chemical Reactions”
Properties of Substances
Name of the Substance
Chemical
Formula
copper(II) sulfate pentahydrate
CuCO3
copper(II) sulfate
PbCl2
hydrogen chloride
Fe(OH)3
lead(II) chromate
NaNO3
ammonia
CaCO3
carbon dioxide
zinc sulfate
ZnCl2
hydrogen
Na2CO3
copper
FeCl3
Color
Physical
State
(g, l, or s)
Solubility in
Water (good
or poor)
Experimentll
Conductivity
Objectives
l. To study the electrical conductivity (electronic or ionic) of various substa..'llces and solutions.
2. To classify a substance according to the conductivity of its solutions.
3. To relate the conductivity of solutions to the nature of the dissolved solute (i.e., ions andJor
molecules).
Discussion
Electrical conductivity is dependent on the mobility of charged particles (electrons or ions!. Certain
types of pure substances and certain types of solutions are capable (to varying extents) of conducting
electricity. Metals are good electronic conductors due to the mobility of their valence electrons.
Molten ionic substances are good ionic conductors due to the liquid state mobility of the ions which
compose them. Certain metalloids called semiconductors are poor conductors. .Most other pure
substances do not conduct and are referred to as nonconductors. Carbon in the form of graphite is an
exception to this; it is a good electronic conductor.
Materials that are ionic conductors are called electrolytes. Many solutions are ionic conductors (or
electrolytes) due to the presence of mobile cations and anions. Solutions that are good conductors are
called strnng electro1vtes; those that are poor conductors are called weak electralvtes; and those that
are nonconductors are called none1ectrolytes.
The primary focus of this experiment is to test the conductivity of various aqueous solutions and from
the results infer whether the solute produces charged species (ions) and/or neutral species (molecules 1
in solution. The aqueous ions will conduct electricity; the molecules will not. The experiment will
be carried out by L.-nmersing two electrodes in each solution and noting the affect on an attached light
bulb. The bulb may glow brightly, dimly, or not at all depending on whether the solution is a strong,
weak, or nonelectrol:vte, respectively. Other factors that affect how brightly the bulb lights up include
solution concentration aI1d the depth the electrode.s are inserted into the solution.
The concepts and results of this experiment lay the foundation for writing net ionic equations.
Procedure
1. The above discussion states the terminology and key concepts involved in this experiment. The
theory and explanation of these concepts will be thoroughly discussed in class.
2. This experiment vvill be conducted as a classroom demonstration during its scheduled lab period
or during convenient lecture periods. The instructor will demonstrate all conductivities of the
materials listed on the report form.
INTENTIONALLY BLANK
Experiment 11 “Conductivity”
Name: ____________________________
Questions
1.
Referring to the experimental results, classify each type of solid and liquid as generally being good (G),
poor (P), or nonconductors (N) of electricity.
Elements
Compounds
Metals
Nonmetals
Ionic
Covalent
Solid State:
______
______
______
______
Liquid State:
______
______
______
______
2.
List those types (categories, not examples) of compounds whose aqueous solutions are generally classified
as strong, weak and nonelectrolytes.
Strong Electrolytes:
Weak Electrolytes:
Nonelectrolytes:
3.
Classify each material (a substance or a mixture) as a good conductor (G), poor conductor (P), or
nonconductor (N) of electricity.
4.
5.
(a)
Ca(s)
_______
(e)
MgCl2(l)
_______
(i)
P4O10(s)
_______
(b)
HCl(g)
_______
(f)
KCl(s)
_______
(j)
HC2H3O2(aq)
_______
(c)
HCl(aq)
_______
(g)
H2S(g)
_______
(k)
HC2H3O2(l)
_______
(d)
NH3(aq)
_______
(h)
H2S(aq)
_______
(l)
CuSO4(aq)
_______
Classify each of the following substances as a strong (S), weak (W), or nonelectrolyte (N).
(a)
HI
_______
(e)
Na2CO3
_______
(i)
CH3CH2OH
_______
(b)
Rb2SO4
_______
(f)
CH3NH2
_______
(j)
Sr(OH)2
_______
(c)
HNO2
_______
(g)
KMnO4
_______
(k)
NH4NO3
_______
(d)
HCN
_______
(h)
FeCl3
_______
(l)
H2SO4
_______
Write the notation appropriate for each substance in full-formula and ionic equations. Examples are given
for the first two substances.
Name
potassium
cyanide
lead (II)
chloride
hydrobromic
acid
cesium
hydroxide
Notation in FullFormula Equaion
Notation in Ionic
Equaion
KCN(aq)
K+(aq) + CN−(aq)
PbCl2(s)
PbCl2(s)
Name
carbonic acid
aluminum
sulfate
zinc
hydroxide
sulfuric acid
ammonia
ethyl alcohol
(CH3CH2OH)
oxygen
water
Notation in FullFormula Equaion
Notation in Ionic
Equaion
INTENTIONALLY BLANK
N/A
N/A
N/A
N/A
Graphite
Silicon
Pure Silicon
Mercury
Sugar Dissolved in
Water
Sugar
Sodium Chloride
Dissolved in Water
Sodium Chloride
Tap Water
Deionized Water
N/A
Copper
Type of
electrolyte (for
pure substance or
solute): Strong,
Weak, or
Nonelectrolyte?
N/A
Observed
Type of
Conductivity:
Electronic or
Ionic?
Aluminum
Predicted
Conductivity:
Good, Poor, None?
N/A
Type of Material:
Element (metal,
nonmetal, metalloid),
Compound (covalent or
ionic) or an Aqueous
Solution?
Iron
Material
Formula (if
applicable)
and (s), (l),
(g), or (aq)
N/A
N/A
N/A
N/A
N/A
N/A
N/A
N/A
N/A
Predominant
Species in
Solution
(besides water):
Molecules or
Ions?
N/A
N/A
N/A
Predominant
Species in
Solution
(besides water):
Molecules or
Ions?
Molten Potassium
Chlorate (Demo!)
Type of
Type of
electrolyte (for
Conductivity: pure substance or
Electronic or solute): Strong,
Ionic?
Weak, or
Nonelectrolyte?
N/A
Predicted
Conductivity:
Good, Poor, None?
Observed
Potassium Chlorate
(Demo!)
Suspension of Lead(II)
Iodide in Water
Suspension of Barium
Sulfate in Water
Ammonia Dissolved in
Water
Phosphoric Acid
Sulfuric Acid
Hydrochloric Acid
Acetic Acid Dissolved
in Water
Pure Acetic Acid
Sodium Hydroxide
Dissolved in Water
Ethanol Dissolved in
Water
Ethanol
Material
Formula (if
applicable)
and (s), (l),
(g), or (aq)
Type of Material:
Element (metal,
nonmetal, metalloid),
Compound (covalent or
ionic) or an Aqueous
Solution?
c. -1
Exercise C
Selected Equations from Experiments 3 and
~
~bjective
To practice writing three forms of chemical equations - full formula, total ionic, and net ionic.
Directions
In this exercise you are to balance full formula, total ionic and net ionic equations for selected reactions from
experiments 3 and 6. The sets of equations are divided into two sections, one for experiment 3 and one for
experiment 6. Each reaction is identified by its original reaction number and by the names of the reactants.
This assignment has two parts, and may be treated as two separate assignments. If completed in this manner,
each part will be assigned at different times and graded separately. \Vb.ether completed as two separate
assignments or as one assignment, follow the directions below.
Part I: Full Formula Equations
Complete the full formula equation only. This would include formulas for reactants and products as well as their
states of matter.
Use your observations from the appropriate experiment to help determine if a reaction occurred.
1. If there was no reaction, write the formulas of the reactants on t11e left side of the full formula
equation box. Next draw an arrow and write N.R. (no reaction) to the right of the arrow.
2. For those reactions that did occur, write the full formula equation. Be sure to balance the equation and
give the states for all formulas as (s }, ill, (g), or (aq).
Part II: Total Ionic and Net Ionic Equations
1. For those reactions that did not occur, write the formulas as appropriate for a total ionic equ~tion.
These formulas should be written on the left in the total ionic equation box. Next draw an arrow and
write N.R. to the right of the arrow . .Also, write the word ,none!! on the far right of the net ionic equation
box.
1
2. For those reactions that did occur, complete the total ionic and net ionic equations. Be sure to
balance the equation and give the states for all formulas as (s), (t), (g), or (aq).
Ex.am pies
Experiment 3
12. Reaction 7 solid product is mb:ed with hydrochloric acid.
G
8
....
Full
Formula
Equation
BaSO is) + HCl(aq) _____.. N.R.
Total
Ionic
BaSOJs) + *H.,.o+ + *Cl-(aq) ____.. N.R.
'!<Hydrochloric acid, HCl{aq), may be represented in ionic form as [ Hp··(aql + c1-1aqJ
Use your instruct.or's preference.
Equation
Net
none
J
QLM.
[
Ionic
Equation
H·\aq) + Cl-(aq)
J.
C-2
20. Solutions of ironmD chloride and sodium hydroxide are mixed.
D
Full
FeCl3 (aq) + 3 NaOH(aq)
~
Fe(OH\/s) + 3 NaCl(aq)
Fe +(aq) + 3 c1-(aq) + 3 N a+(aq) + 3 OH-(aq) ___. Fe{OH)~(s) + 3 N a"'"(aq) + 3 c1-(aq)
G
3
Formuia
Equation
'!'otal
Ionic
Equation
Net
3
Fe +(aq) + 3 OH-(aq) __.. Fe(OH)'.ls)
Io me
Equation
Exercise C
Name
Selected Equations from Experiments 3 and 6
Experiment. 3
Hydrochloric acid is mixed with sodium carbonate solution.
L
[]
Full
Formula
Equation
Total
ioruc
Equation
G
Net
Ionic
Equation
4. Sulfuric acid and potassium carbonate solutions are mixed.
ID
I
I
~
Formula
F===============Equai:km
IGI
II
Ionic
I ~~tion
Ii Total
.
!.!:::::~===============================================================================::!
Equation
5. Sulfuric acid and barium chloride solutions are mixed.
1r--i
Fnll
E'orm.ttla
Equation
I I
I~
l
IGI
l
II
I
J
Totai
Ionic
Equation
Net
Ionic
Equation
6. Solutions of barium chloride and sodium sulfate are mixed.
~
IDr1·~=======~formula
~
Equation
Total
Ionic
G
IIil====================;1Equatlot1
l.!::::::=::==============================================================================::::!i
Net
Ionic
Equatiou
7. Solutions of barium chloride and ammonium. sulfate are mixed.
G
1L::====================================================================~Formcla
11
M
;;;;;;
Equation
:
.I
·l==============;'I
'11LJ
!!:::::db:====================================================================~
Total
lor.i.c
Equation
Net
lomc
Equation
C-4
8. Solutions of barium chloride and sodium carbonate are mixed.
D
I
~
Formula
~==================================~Equation
Total
Iomc
l:=====================t'lEqua~ion
G
II
Net
lonic
Equation
9. Solutions of barium chloride and potassium carbonate are mixed.
D
G
Full
Formula
Equatlon
Total
Ionic
Equation
Net
Ionic
Equation
10. Solutions of barium chloride and ammonium carbonate are mixed.
IOI
~I
Full
For.nula
Equat10n
I
================;1
§
G
, I
1
TI
tion
lortic
1
i!:=:::::!J========================================================::====================================..l~quation
11. Reaction 6 solid product is mbrnd with hydrochloric acid.
D
G
Full
Formula
Equation
l::===============================================================================::;::::==========~Total
Iorue
II~~~~~~~~~~~tioo
IL!::=:::::!.!:::====================================================================::::::!J
Net
lomc
Equation
13. Reaction 8 solid product is mixed with hydrochloric acid.
D
G
Full
For::nula
Equation
Total
ioruc
Equation
Net
Ionic
Equation
c -3
16. Sodium hydroxide solution is mixed with hydrochloric acid.
Full
Formula
\:---1
u
1.---.~============================================================1
II : I
~~
Total
Ionic
~~~
~========================================================================================:::;'!Net
ll
lo
l!::::================================================================================.l
Ionic
Equation
17. Sulfuric acid and potassium hydroxide solutions are mixed.
[~
Full
Formula
Equation
B
:~========================================~ Total
~
I
I nl!l~=================-'iEquation
Net
I ==============================================================================:=l
Ionic
Equation
18. Ammonia solution is mL'i:ed with hydrochloric aci~:~
l,l~I
LJ
1~
Equation
i======================================================================~ Total
~I
I
II
I
Ionic
t==================;'iEquation
:I
~et
:oruc
II
==~IL~========================================================================================:::::!.!Bquation
19. Sulfuric acid and ammonia solutions are mi.""{ed.
D
Full
Formula
Equaton
Total
Ior.ic
.Equation
.Bl
I
Net
Iorue
Equation
21. Solutions of iron{III) chloride and potassium hydroxide are mixed.
[]
Full
Formula
Equation
Total
Iomc
Equation
18
Net
Ionic
Equation
C-6
23. Solutions of copper(Il) sulfate and sodium hydroxide are mixed.
[]
Full
Formula
Equation
Total
Ionic
G
Equation
Net
Ionic
.Squat1on
24. Solutions of copper(ll) sulfate and potassium hydroxide are mixed.
Full
?or:nula
[ I]
Equation
~==========~Total
G
IIi=============:;'J
Ionic
Equation
}iet
Ionic
t===.t::======================================================================~
Equation
Experiment 6
C. Chemical Equations for Single Replacement Reactions
1. Zinc is added to copper (II) sulfate solution.
[j
Full
Formui::i.
Eqliat1.:m
l================lTotal
Gi
Ionic
Equation
~==========================================================================================;iNet
l!:::::=2.!:::======================================================================~
~~-
Ionic
Equation
Copper is added to zinc sulfate solution.
[]
Full
F1Jrmula
'Equation
.l================lTotal
lomc
Equation
GI1===========================================================================;1Net
Ionic
Equation
3. Copper is added to silver nitrate solution.
DI
I
Full
Formula
·===============~Equation
D
Total
Ionic
Equation
l========:===:=:::;'lNet
1
!!=:::!b:========================================================================~
:Orne
Equation
....
(_,
4. Copper is added to hydrochloric acid.
IOI
I
I
I F=========================================================================================::::::jTotal
l
?ull
F~~
T
1
j
Equation
I
Ionic
II
Equation
Net
Ionic
Equation
iJ
lJ
i:==::.;==========================================================================================~
fia. Zinc is added to hydrochloric acid.
ID
G
Full
Formula
Equation
r======-=======~Total
Ionic
Equation
!=================================================================================;'?Net
Ionic
i====.::::=====================================::=:====================================================.t
Equation
D. Chemical Equations for Double Replacement Reactions
l. Lead(Il) nitrate solution and hydrochloric aeid are mixed.
L
lr-1
I
1,\rl
l __j
~~~~~~~~~~~·~~~~~~~~~~~~~~~~~~~~~~~~~~~~~,
I Equat..ion
~'~muta
1
1
1
=::====================================================================================~
Total
~~~~~tion
I
J ~et
I .I
I !::==============================================111 ~~tlon
1'1
2. Lead
<ID nitrate solution is mixed with sodium chromate soiution.
.
D
G
Full
formula
Equation
Total
Iomc
Equation
'.\et
Ionic
E·qua~on
3. Zinc sulfate solution and hydrochloric acid are mixed.
l[J
Full
Formula
Zquation
Total
H
-
Ionic
Equation
Net
Ionic
.
Equation
- ,...'
c- s
E. Equations for other Important Reactions
1. Solutions of iron(III) chloride and ammonia are
mi.~ed.
D
iIG
Full
Formula
Equation
Totai
Ionic
Equation
-
Net
I
Ionic
Equation
2. Hydrochloric acid is added to solid sodium carbonate.
I
ID\
I
I!GI
Formula
Equat1cn
~==============================================================================:::=======:
!
II
·1·_
I
~
.1
! ~:U~tior.
!Total
]J'~~
:!.!::=============================================i Equatton