Laboratory Manual
For
General Chemistry I
Prepared By
Manal H. Khabbas
Ebtesam Mansour
Department of Chemistry
Faculty of Science
University of Hail
2011
1
Glassware and Equipment
2
Safety Rules in the Laboratory
The following is a list of general safety rules to be followed by
everyone!
1. Safety goggles and Laboratory coat must be worn at all times in the Lab.
2. Eating, drinking, and smoking in the laboratory is strictly forbidden.
3. Laboratory chemicals are not to be tasted since many are toxic. When instructed
to smell reagents, do so with great caution, NEVER put your nose over the bottle!
Avoid looking into the mouth of any reaction vessel in which a reaction is in
progress. Never point a test tube that contains a heated liquid at anyone.
4. No one will perform any unauthorized experiments, nor will anyone work in the
lab alone, or outside of regularly scheduled hours.
5. Tie back long hair and loose clothing when performing laboratory experiments.
6. Report ALL injuries, allergies and/or medical problems to the lab instructor.
7. When pouring out of a reagent bottle, always read the label twice to be certain
that you are using the correct material; always hold the bottle by placing your
hand over the label.
8. Label every chemical container to avoid mix-ups.
9. NEVER return excess chemicals to the reagent bottle.
10. Keep all lids to chemicals closed.
11. Dispose of the excess chemicals in the proper waste container, as indicated by
the lab instructor.
12. Clean up any spills using the chemical instruction sheet.
13. Leave your work area neat and clean.
14. Wash your hands with soap and water before leaving the laboratory. This rule
applies even if you have been wearing gloves.
All chemicals in the lab are to be considered dangerous
3
Safety Equipment
Know the location of:
1.
2.
3.
4.
5.
6.
First-aid kits
Fume Hoods
Safety showers and Eyewash fountains
Fire blankets
Fire extinguishers and fire alarm
Emergency exits.
Heating and Fire Safety:
1. Never reach across an open flame.
2. Know how to light and extinguish the Bunsen burner, never leave a burner
unattended.
3. Point the test tube or bottle away from you and others when being heated;
chemicals can rapidly boil out of the container.
4. Never heat a liquid in a closed container.
5. Always use a clamp or tongs when handling hot containers.
Questions:
For each of the following questions, circle the answer that most correctly
answers the question.
1. After completing an experiment, all chemical wastes should be
a)
b)
c)
d)
left at your lab station for the next class.
disposed of according to your instructor’s directions.
dumped in the sink.
taken home.
2. If an acid is splashed on your skin, wash at once with
a)
b)
c)
d)
soap.
oil.
weak base.
plenty of water.
4
Experiment 1
Density of a Liquid and a Solid
Objective: This experiment is designed as an exercise to refine your skills in
weighing, measuring a volume of liquid using a graduated cylinder and volumetric
flask, precision in measuring, the use of significant figures in measurement and
calculations, and the determination of density.
Introduction
The density of a material is defined as its mass per unit volume. Density is an
intensive property (independent of the size of the sample) but it does depend on the
temperature.
Density =
Procedure
(A)
Density of a Liquid
This will be determined using a graduated cylinder and volumetric
flask.
(A-1) Graduated cylinder (this measures volumes to one place of decimals)
1- Weigh a 25 mL graduated cylinder on the lab balance.
2- Add approximately 10 mL of water to the cylinder (use your wash bottle) and
measure the volume of the water after you pour it in.
3- Weigh the cylinder and water.
4- Empty the cylinder. Repeat pouring the liquid, weighing and measuring, two
more times.
(A-2) Volumetric flask (this measures a volume to two places of decimals)
5
-Weigh a 50 mL volumetric flask on the lab balance. Fill the volumetric flask to
the mark on the neck with water. Weigh the flask containing the water. Record the
mass. Empty the flask. Repeat pouring the water and weighing two more times.
- Record the temperature of the water used.
(B) Density of a Solid
(Volume measured by the displacement of the water)
1- Weigh a piece of metal on the analytical balance.
2- Measure the volume of the metal:
a) Fill a graduated cylinder halfway with sink water.
b) Tap out any air bubbles.
c) Record initial volume of water.
d) Tilt the cylinder gently and slide the metal into it. It must be submerged.
e) Tap out any air bubbles.
f) Record final volume (water + metal).
6
Experiment 1 Report
Density of a Liquid and a Solid
Name: _________________
ID: __________________
Date:___________________
Instructor: _____________
Data and Data Processing
(A-1) Table 1: Density of a Liquid (graduated cylinder)
Run1
Run 2
Run 3
Mass of graduated cylinder (g)
Mass of graduated cylinder + water (g)
Mass of water (g)
Volume of water (mL)
Density of water (g/mL)
Average density of water (g/mL)
(A-2) Table 2: Density of a Liquid (volumetric flask)
Run1
Mass of volumetric flask
Mass of volumetric flask + water
Mass of water (g)
Volume of water
Density of water (g/mL)
Average density of water (g/mL)
Temperature of water (O C): -----------
7
Run 2
Run 3
B- Table 3: Density of a metal
Mass of metal (g)
initial volume of water (mL)
final volume of water and metal
(mL)
volume of metal (mL)
Density of metal (g/mL)
Questions
1- Which vessel that you used to estimate the density of water, produced the
more precise sets of measurements? Explain.
2- Which was the more accurate vessel that you used? Explain
(see table of densities on last page).
3- In the measurement of the density of solid; why won’t this technique
work to measure densities of materials that are less than 1 g/mL?
8
9
Experiment 2
Stoichiometry
Determining the Formula of a Hydrate
Objective: The purpose of this experiment is to determine the empirical formula of
a hydrate
Introduction
Hydrated salts (or Hydrates) are salts which have a definite amount of water
chemically combined. Some common hydrates are:
CuSO4. 5H2O : copper (II) sulphate pentahydrate
MgSO4.7H2O: magnesium sulphate heptahydrate
On heating, the attractive forces are overcome and the water molecules are released
leaving behind the anhydrous salt. e.g.
CuSO4. 5H2O(s)
Δ
CuSO4(s) + 5H2O(g)
(Blue)
( white)
hydrated
anhydrous
copper (II) sulphate
copper (II) sulphat
Procedure
1- Clean a porcelain crucible and cover , support the crucible and cover on a
clay triangle on a ring stand (fig. 1). Heat with a Bunsen burner, first gently,
and then to redness for about 3 minutes. Allow to cool completely.
2- Weigh the crucible and cover to nearest 0.001g
3- Transfer about one gram of the unknown hydrate to the crucible and reweigh
(Crucible + cover + hydrated salt).
10
4- Support the crucible and cover on a clay triangle on a ring stand and heat
gently for about 5 minutes. You may wish to lift the cover occasionally to
observe any changes in material. When heating is finished, place the cover
completely on crucible and allow the crucible to cool completely.
5- Weigh the crucible and anhydrous salt.
6- Return the crucible to the clay triangle and heat again for 3 minutes. Cool
and weigh again. If the weight after second heating is different from the first
one by more than 0.001 g, heat for third time, cool and re-weigh.
Figure 1: Heating the crucible
11
Experiment 2 Report
Determining the Formula of a Hydrate
Name: _________________
ID: __________________
Date:___________________
Instructor: _____________
Hydrate: ____________.XH2O
Mass of crucible
Mass of crucible + hydrate
Mass of crucible + anhydrous
Observation:
__________________________________________________________________
__________________________________________________________________
Calculation:
1- Mass of hydrate:______________________________________________
2- Mass of anhydrous:____________________________________________
3- Mass of water lost:_____________________________________________
4- Number of moles of anhydrous salt:_______________________________
________________________________________________________________
5- Number of moles of water:______________________________________
________________________________________________________________
12
6- The value of “X” in the formula:
(X = number of moles of water /number of moles of anhydrous salt)
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
7- Percent mass of water of hydration:
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
Questions
1- Calculate the percent mass of water of hydration in BaCl2.2H2O.
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
2- What is the effect (decrease, increase, no effect) of the following on the
calculated value of X? Explain
a) Incomplete dehydration of the hydrate.
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
b) Using 1.5 g of hydrate instead of 1.0 g.
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
13
Experiment 3
Stoichiometry
Empirical Formula of Magnesium Oxide
Objective: In this experiment, you will determine the empirical formula of
magnesium oxide.
Introduction
When magnesium and oxygen are heated together, they readily undergo a chemical
change (reaction):
2Mg(s) + O2(g) → 2MgO(s)
Simultaneously, magnesium reacts with the nitrogen in the air to form magnesium
nitride.
3Mg(s) + N2(g) → Mg3N2(s)
The amount of nitride that forms can be removed with the addition of water, which
converts the nitride to magnesium hydroxide and ammonia gas. The magnesium
hydroxide is heated to high temperature and decomposes to magnesium oxide and
water. Thus, at the end of the experiment, all of the magnesium has been converted
to the desired product, magnesium oxide.
Mg3N2(s) + 6H2O(l) → 3Mg(OH)2(s) + 2NH3(g)
Mg(OH)2(s)
∆
MgO(s) + H2O(l)
Materials
Safety goggles, Magnesium ribbon( Mg), Balance , Ring stand, Bunsen burner,
Ring support/ clay triangle, Crucible/ cover, Tongs
Procedure
1- Clean a porcelain crucible and cover support the crucible and cover on a clay
triangle on a ring stand. Heat with a Bunsen burner, first gently, and then to
redness for about 3 minutes. Allow to cool completely.
2- Weigh the crucible and cover.
3- Clean Mg ribbon weighing approximately 0.3 grams with sandpaper to
14
remove any oxide coating.
4- Coil the ribbon very loosely and place on the bottom of the crucible. Then,
weigh the crucible with the Mg ribbon inside. Record the weight (crucible +
Mg)
5- Place the cover on the crucible. Heat the crucible gently for 5 mins. while
using the tongs to lift the cover slightly every 30 sec. to admit air. Should the
Mg start glowing brightly when the cover is lifted, quickly cover the
crucible, remove the bunsen burner, and wait one min before continuing to
heat.
6- Heat the covered crucible strongly for 15 min. (lifting the cover
occasionally).
7- Lift the cover to determine whether the ribbon has become a whitish ash. If
the ribbon still has its original color, reheat for 10 min. Repeat step 9 until
the ribbon has become a whitish ash then, allow the crucible to cool.
8- To a cooled crucible, add 10 drops of deionized water.
9- Partially cover the crucible (leave a slight crack) and heat gently for 2 mins.,
then strongly for 10 min. Allow the crucible to cool to room temperature.
10- Weigh the crucible and the product.
Experiment 3 Report
Empirical Formula of Magnesium Oxide
Name: _________________
Date:___________________
ID: __________________
15
Instructor: _____________
Data :
Molar mass of Mg = 24.31 g/mol,
Molar mass of O
=16.00 g/mol
Mass of crucible + cover (g)
Mass of crucible + cover + Magnesium (g)
Mass of crucible + cover + Magnesium oxide (g)
Observation:
__________________________________________________________________
__________________________________________________________________
Calculation:
1- Mass of magnesium: ___________________________________________
2- Mass of magnesium oxide:_______________________________________
_____________________________________________________________
3- Mass of oxygen: _______________________________________________
________________________________________________________________
4- Moles of Mg:__________________________________________________
_____________________________________________________________
5- Moles of O: __________________________________________________
_____________________________________________________________
6- Calculate the empirical formula of the magnesium oxide
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
16
Experiment 4
Precipitation Reactions
Objectives
1) To observe precipitation reactions involving various ionic solutions.
2) To practice writing complete ionic and net ionic equations and to identify
spectator ions.
Introduction
Most ionic solids are soluble in water because the polar water molecules surround
the individual ions of the salt. Those that do not dissolve and go into solution form
solid products called precipitates. These precipitates have many colors and often
help scientists identify what the precipitate is present. A precipitate can be
identified by the cloudy, milky, gelatinous, or grainy appearance it gives to the
mixture.
A barium sulfate precipitate can be produced by the reaction of barium chloride
and sodium sulfate. A chemical equation to describe the reaction is written and
balanced like this:
BaCl2 (aq) + Na2SO4 (aq) Æ 2 NaCl (aq) + BaSO4(s)
This form of the equation is generally referred to as the molecular equation.
To write a total ionic equation, rewrite all aqueous substances as their component
ions and keep all solid substances unchanged.
Ba+2(aq) + 2Cl- (aq) + 2Na+(aq) + SO4-2(aq) Æ 2Na+(aq) + 2Cl-(aq) + BaSO4(s)
This form is known as the complete ionic equation. This reaction occurs because
the insoluble substance, BaSO4, precipitates out of solution. The other product,
barium nitrate is soluble in water and remains in solution. We see that Na+ and Clions appear on both sides of the equation and thus do not enter into the reaction.
Such ions are called spectator ions. If we eliminate them from both sides, we
obtain the net ionic equation.
Ba+2(aq) + SO4-2(aq) Æ BaSO4(s)
17
Procedure
1.
The report lists 10 pairs of chemicals to be mixed. Use about 1 mL
(20 drops) of each of the reagents to the test tube, as indicated on the
report sheet. Record your observations on the report sheet. If it is not
apparent that a reaction has occurred, write “nothing observed.”
2.
Dispose of the solutions containing barium, copper, and chromate in
the appropriate waste container provided. All other solutions can be
rinsed down the sink
18
Experiment 4 Report
Precipitation Reactions
Name: _________________
ID: __________________
Date:___________________
Instructor: _____________
Results
For each reaction listed below, record your observations.
REACTION
OBSERVATIONS
1. Copper (II) sulfate and sodium
hydroxide
2. Silver nitrate and dilute hydrochloric
acid.
3. Barium nitrate and copper (II)
sulfate
4. Silver nitrate and potassium iodide.
5. Silver nitrate and potassium
chromate.
6.Ferric Chloride and dilute ammonia
Solution
7.Chromium chloride and dilute
ammonia solution
8.Calcium chloride and copper (II)
sulfate
9.Sodium chloride and Silver nitrate
10. Sodium chloride and copper (II)
sulfate
11. Calcium chloride and potassium
carbonate
12. potassium carbonate and Silver
nitrate
19
Write a molecular equation, total ionic equation, and net ionic equation for all 12
of the reactions that were carried out. Physical states, aqueous (aq), solid (s),
liquid (l), and gas (g), must be shown for the molecular and net ionic equations.
These equations should be hand-written and handed in before leaving the
laboratory.
1) Copper (II) sulfate and sodium hydroxide
Overall Equation:
__________________________________________________________________
Complete Ionic Equation:
__________________________________________________________________
Net Ionic Equation:
__________________________________________________________________
2) Silver nitrate and dilute hydrochloric acid. . Don’t forget to continue and
write the decomposition reaction that leads to the formation of the gas. Show this
at the net ionic equation.
Overall Equation:
__________________________________________________________________
Complete Ionic Equation:
__________________________________________________________________
Net Ionic Equation:
__________________________________________________________________
3) Barium nitrate and copper (II) sulfate.
Overall Equation:
__________________________________________________________________
Complete Ionic Equation:
__________________________________________________________________
Net Ionic Equation:
__________________________________________________________________
20
4)Silver nitrate and potassium iodide
Overall Equation:
__________________________________________________________________
Complete Ionic Equation:
__________________________________________________________________
Net Ionic Equation:
__________________________________________________________________
5)Silver nitrate and potassium chromate
Overall Equation:
__________________________________________________________________
Complete Ionic Equation:
__________________________________________________________________
Net Ionic Equation:
__________________________________________________________________
6)Ferric Chloride and dilute ammonia Solution
Overall Equation:
__________________________________________________________________
Complete Ionic Equation:
__________________________________________________________________
Net Ionic Equation:
__________________________________________________________________
21
7) Chromium chloride and dilute ammonia solution
Overall Equation:
__________________________________________________________________
Complete Ionic Equation:
__________________________________________________________________
Net Ionic Equation:
__________________________________________________________________
8) Calcium Chloride and copper (II) sulfate
Overall Equation:
__________________________________________________________________
Complete Ionic Equation:
__________________________________________________________________
Net Ionic Equation:
__________________________________________________________________
9)Sodium chloride and silver nitrate
Overall Equation:
__________________________________________________________________
Complete Ionic Equation:
__________________________________________________________________
Net Ionic Equation:
__________________________________________________________________
22
10) Sodium chloride and copper (II) sulfate
Overall Equation:
__________________________________________________________________
Complete Ionic Equation:
__________________________________________________________________
Net Ionic Equation:
__________________________________________________________________
11) Calcium chloride and potassium carbonate
Overall Equation:
__________________________________________________________________
Complete Ionic Equation:
__________________________________________________________________
Net Ionic Equation:
__________________________________________________________________
12) potassium carbonate and Silver nitrate
Overall Equation:
__________________________________________________________________
Complete Ionic Equation:
__________________________________________________________________
Net Ionic Equation:
__________________________________________________________________
23
Solubility Rules:
1. Most nitrate (NO3−) salts are soluble.
2. Most alkali (group 1A) salts and NH4+ are soluble.
3. Most Cl−, Br−, and I− salts are soluble (except Ag+, Pb2+, Hg22+).
4. Most sulfate salts are soluble (except BaSO4, PbSO4, Hg2SO4,
CaSO4).
5. Most OH− salts are only slightly soluble (NaOH, KOH are soluble,
-----Ba(OH)2, Ca(OH)2 are marginally soluble).
6. Most S2−, CO32−, CrO42−, PO43− salts are only slightly soluble.
24
Experiment 5
Acid Base Titration
Objective:
Determine the concentration of acetic acid in a vinegar sample
Introduction:
Volumetric analysis is a quantitative analytical process based on measuring
volumes. The most common form of volumetric analysis is the titration, a process
whereby a standard solution of known concentration is chemically reacted with a
solution of unknown concentration (analyte) in order to determine the
concentration of the unknown.
Vinegar is a dilute solution of acetic acid (CH3COOH), and in this experiment the
acetic acid (CH3COOH) is the analyte and sodium hydroxide (NaOH) is the
standard solution. The reaction is:
CH3COOH(aq) + NaOH(aq) → CH3COONa(aq) + H2O(l)
Procedure:
1- Make sure that your buret is clean. If necessary wash it with soap solution
using a buret brush. Rinse it several times with deionized water. Then rinse
it twice with approximately 10 mL of the 0.10 M NaOH solution to be used
in the titration. Drain the solution through the buret tip. Fill the buret with
the 0.10 M NaOH solution; make sure there are no air bubbles in the tip of
the buret or just above the stopcock. Run the base out of the buret until the
level is at 0.00 or below. Record the level of the base, estimating the
reading to two decimal places.
2- Rinse a clean 10.0 mL pipet with several small samples of the vinegar
solution. Use the pipet to transfer 10.0 mL of vinegar to a clean 250 ml
25
erlenmeyer flask. Your instructor will demonstrate the use of the pipet. Add
3 drops of phenolphthalein indicator to the vinegar sample.
3- Slowly run the base out of the buret into the vinegar solution, swirling the
flask and contents (figure 1). As you approach the equivalence point, the
area in the vinegar where the drop of base falls will turn pink; then the pink
color will disappear as the solution becomes mixed. From this point on, add
the base dropwise with constant swirling. Occasionally wash down the sides
of the flask with water from your wash bottle. The equivalence point is
where 1 drop (or less) of base causes the solution to become very pale pink
throughout. Record the final buret reading, estimating it to the nearest 0.05
mL.
4- Repeat the titration using a clean flask. After the first titration, the second
should go more quickly since you now have some idea of how much base is
required per 10.0 mL sample of vinegar. The base may be added quickly
until you are within 2 or 3 mL of the equivalence point; then change to
dropwise addition.
26
Figure 1: Acid base titration 27
Experiment 5 Report
Acid Base Titration
Name: _________________
ID: __________________
Date:___________________
Instructor: _____________
Data Collected:
Volume of vinegar solution (ml)
Run 1
Run 2
10.0 ±
10.0 ±
Initial buret reading (ml)
final buret reading (ml)
Volume of sodium hydroxide
needed to reach end point (ml)
Average volume of sodium
hydroxide (ml)
Data Processing
1- From the above data calculate the molarity of acetic acid in vinegar
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
2- Calculate the mass of acetic acid in the vinegar per liter (g/L).
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
28
Questions:
1- Why is it a good idea to rinse the buret with the NaOH solution, instead of
with water, before filling it at the start of the titration?
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
2- If there is an air bubble in the buret, how this will affect the calculated
concentration of acetic acid in vinegar?
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
29
Experiment 6
Redox Titration
Analysis of Commercial Bleach
Objectives:
To determine the amount of sodium hypochlorite present in commercial bleach.
Introduction:
Many commercial products, such as bleaches and hair coloring agents,
contain oxidizing agents. The most common oxidizing agent in bleaches is sodium
hypochlorite, NaClO. Commercial bleaches are made by bubbling chlorine gas
into a sodium hydroxide solution. Some of the chlorine is oxidized to the
hypochlorite ion, ClO-, and some is reduced to the chloride ion, Cl-. The solution
remains strongly basic. The chemical equation for the process is:
Cl2(g) + 2OH-(aq) → ClO-(aq) + Cl-(aq) + H2O(l)
The amount of hypochlorite ion present in a solution of bleach can be determined
by an oxidation-reduction titration. One of the best methods is the iodinethiosulfate titration procedure. The iodide ion, I-, is easily oxidized by almost any
oxidizing agent. In acid solution, hypochlorite ions oxidize iodide ions to form
iodine, I2. The iodine that forms is then titrated with a standard solution of sodium
thiosulfate. The analysis takes place in a series of steps:
1. Acidified iodide ion is added to hypochlorite ion solution and the iodide
is oxidized to iodine.
2 H+(aq) + ClO-(aq) + 2 I-(aq)
Cl-(aq) + I2(aq) + H2O(l)
2. Iodine (I2) is only slightly soluble in water, but it dissolves very well in
an aqueous solution of iodide ion, in which it forms a complex ion called
the triiodide ion (I3-). Triiodide ion is a combination of a neutral I2
molecule with an I- ion. The triiodide ion is yellow in dilute solution and
dark red-brown when concentrated.
I2(aq) + I-(aq)
→
30
I3-(aq)
3. The triiodide ion is titrated with a standard solution of thiosulfate ions,
which reduces the iodine back to iodide ions.
I3-(aq) + 2 S2O32-(aq)
3 I-(aq) + S4O62-(aq)
During this last reaction the red-brown color of the triiodide ion fades to
yellow and then to the clear color of the iodide ion. It is possible to use the
disappearance of the color of the triiodide ion as the method of determining the end
point, but this is not a very sensitive procedure. Addition of starch to a solution
that contains iodine or triiodide ion forms a reversible blue complex. The
disappearance of this blue colored complex is a much more sensitive method of
determining the end point. However, if the starch is added to a solution which
contains a great deal of iodine, the complex which forms may not be reversible.
Therefore, the starch is not added until shortly before the end point is reached. The
quantity of thiosulfate used in step 3 is directly related to the amount of
hypochlorite initially present.
ClO-(aq) + 2 S2O32-(aq) + 2H+(aq) →
Cl-(aq) + S4O62-(aq) + H2O(l)
SAFETY ALERT : Concentrated bleach is damaging to skin, eyes, and
clothing. If you spill bleach solution on yourself, wash off with plenty of water.
Materials:
250 Erlenmeyer flask, pipets; 5 and 25 mL, 100 mL Volumetric flask,
buret, graduated cylinders, 2% starch solution, 3.0 M HCl, 0.10 M Na2S2O3, solid
KI.
Procedure:
1. Dilute the concentrated bleach.
Use a 5-mL volumetric pipet to measure 5.00 mL of a commercial bleach
solution into a 100-mL volumetric flask. Dilute to the mark with distilled
water, stopper and mix well.
2. Weigh out approximately 1.5 g solid KI. This a large excess over that which
is needed.
31
3. Pipet 25.00 mL of the dilute bleach into an Erlenmeyer flask. Add the solid
KI and about 25 mL distilled water. Swirl to dissolve the KI. Slowly, with
swirling, add approximately 5 mL of 3 M HCl. The solution should be dark
yellow to red-brown from the presence of the I3- complex ions.
4. Rinse the buret with about 5 mL of 0.100 M sodium thiosulfate solution and
then fill the buret. Record the initial volume.
5.
Titrate the solution in the flask with a standard 0.100 M sodium thiosulfate
solution until the iodine color becomes light yellow. Add 2 mL of starch
indicator. The blue color of the starch-iodine complex should appear.
Continue to titrate until one drop of Na2S2O3 solution causes the blue color to
disappear. Record the final volume.
6. Repeat the titration beginning with step 2.
32
Experiment 6 Report
Analysis of Commercial Bleach
Name: _________________
ID: __________________
Date:___________________
Instructor: _____________
Results and Calculations:
Trial 1
Volume of bleach solution
Initial buret reading
Final buret reading
Volume of Na2S2O3 solution
Number of moles of Na2S2O3
Number of moles of ClOMolarity of NaClO in diluted
bleach solution
Molarity of NaClO in original
bleach solution
Average molarity of NaClO in
original bleach solution
Trial 2
mL
mL
mL
mL
mol
mol
M
mL
mL
mL
mL
mol
mol
M
M
M
M
- Calculate the mass percent of NaClO in the original bleach solution.
(The density of bleach solution is 1.084 g/mL)
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
33
Questions:
1. Write balanced oxidation and reduction half-reactions for the equations in
step1 and step 3. For each half-reaction, identify which substance is oxidized
or reduced.
________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
2. How would each of the following laboratory mistakes affect the calculated
value of the percent NaClO in the commercial bleach (decrease, increase, no
change)? Explain.
a) The pipet was rinsed with distilled water immediately before being used to
measure the commercial bleach solution.
_____________________________________________________________
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________
b) If 3 g of KI was used instead of 2.
_____________________________________________________________
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________
34
Experiment 7
Determination of Molar Mass for a Volatile Liquid
Objective: The molar mass of a volatile liquid will be determined by measuring
what mass of vapor of the liquid is needed to fill a flask of known volume at a
particular temperature and pressure.
Introduction
This experiment involves measuring the molar mass of a volatile liquid by
applications of the Ideal Gas Law:
PV = nRT
Where
•
•
•
•
•
P = Pressure
V = Volume
n = number of moles
R = Universal gas constant (0.08206 L. atm /K . mol)
T = Temperature in Kelvin
In this experiment, a small amount of easily volatilized liquid will be placed in a
flask of known volume. The flask will be heated in a boiling-water bath and will
be equilibrated with atmospheric pressure. From the volume of the flask used, the
temperature of the boiling-water bath, and the atmospheric pressure, the number of
moles of gas contained in the flask may be calculated. From the mass of the
volatile liquid, m, required to fill the flask with vapor when it is in the boiling
water bath, the molar mass of the volatile liquid, M, could be calculated.
35
Equipment:
250 mL Erlenmeyer flask, large beaker, Needle or pin for poking hole, aluminum
foil, thermometer, hot plate, balance, graduated cylinder.
Procedure:
1. Put on your safety goggles.
2. Record the barometric pressure.
3. Set up a boiling water bath using a large beaker 3/4ths full of water.
4. Obtain a sample of the unknown volatile liquid.
5. Cut a small piece of aluminum foil (big enough to fit over the mouth of
the Erlenmeyer flask) and poke a small hole in the middle of it with a
pin.
6. Determine the mass of the flask, aluminum foil.
7. Remove the foil cap. Place a 5.0-mL of the liquid into the flask and
replace the foil.
8. Clamp the flask with a single buret clamp. Transfer the flask to the
boiling water bath, immerse, and heat.
9. Monitor the temperature of the water in the water path with a
thermometer.
10. When the volatile liquid has completely evaporated, remove the flask
from the hot water bath and record the temperature of the water (this
will be the temperature of the vapor in the flask also).
11. Remove the clamp and carefully place the flask into a beaker of ice
water.
12. Once the vapor has condensed (this should only take a few minutes), dry
the flask and determine the mass of the flask, aluminum foil, and
volatile liquid.
13. Dispose of the contents of the flask according to instructions. Fill the
flask with water. Pour the water into a 250-mL graduated cylinder,
measure the volume, and record.
36
Experiment 7 Report
Determination of Molar Mass for a Volatile Liquid
Name: _________________
ID: __________________
Date:___________________
Instructor: _____________
Data Collected:
Mass of flask and aluminum foil
Mass of flask, foil, and condensed liquid
Mass of condensed liquid
Atmospheric pressure
Temperature of hot water bath
Volume of flask
Atmospheric pressure
Temperature of hot water bath
Volume of flask
g
g
g
mmHg
°C
mL
atm
K
L
Calculations:
Calculate the molar mass of the volatile liquid g/mol
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
37
Questions:
- Would your calculated value for the molar mass of an unknown liquid
INCREASE, DECREASE, or REMAIN UNCHANGED with the following
procedural errors. Explain your reasoning!
a) All of the liquid was not vaporized before the flask was taken away from the
heat.
__________________________________________________________________
__________________________________________________________________
b) The flask was not allowed to remain in the water bath as long enough to reach
thermal equilibrium with the water.
__________________________________________________________________
__________________________________________________________________
c) The inside of the flask contained drops of nonvolatile impurity for the initial
massing.
__________________________________________________________________
__________________________________________________________________
d) The flask was not thoroughly dried for the final massing.
__________________________________________________________________
__________________________________________________________________
e) If the atmospheric pressure is recorded 1 atm instead of real pressure inside
the lab.
________________________________________________________________
________________________________________________________________
38
Experiment 8
Heat of Neutralization
Objective:
To measure, using a calorimeter, the energy change accompanying neutralization reactions.
Introduction:
Every chemical change is accompanied by a change in energy, usually in the
form of heat. The energy change of a reaction that occurs at constant pressure is
termed the heat of reaction or the enthalpy change. The symbol ΔH is used to
denote the enthalpy change. In this experiment, you will measure the heat of
neutralization when an acid and base react to form 1 mole of water.
HCl(aq) + NaOH(aq)
→ H2O(l) + NaCl(aq) + heat
This quantity of heat is measured experimentally by allowing the reaction to take
place in a thermally insulated Styrofoam cup calorimeter. The heat liberated in the
neutralization reaction will cause an increase in the temperature of the solution and
of the calorimeter. The heat lost by the neutralization reaction will equal the heat
gained by the water and calorimeter. Because we are concerned with the heat of the
reaction and because some heat is absorbed by the calorimeter itself. The heat
capacity of the calorimeter (in this case, two Styrofoam cups) usually would be
calculated first; however, the heat capacity of the cups is so small that it can be
neglected.
The First Law of Thermodynamics applies and:
Heat (q) lost by the reaction + heat (q) gained by the salt water solution = 0
39
q = (s) (m) (ΔT)
where: s = specific heat capacity of the solution
m = mass of the solution
ΔT= change in temperature of the solution (Tfinal – Tinitial)
The initial temperatures of the acid and base solutions before mixing will be
averaged to get Tinitial. The maximum temperature that is reached after the time of
mixing will be recorded to be the solution’s final temperature. Start monitoring the
temperature as soon as the solutions have been mixed.
Equipment
2 Styrofoam cups, Thermometers ,50 or 100 mL graduated cylinder, stirring rod
150 mL beaker
Chemicals
2.00 M HCl, 2.00 M NaOH
Procedure
1. Rinse and dry the Styrofoam cups (calorimeter). Place one cup inside the
other.
2. Measure out 25.0 mL of 2.00 M NaOH and pour it into the calorimeter. In a
clean dry 150 mL beaker, take 25.0 mL of 2.00 M HCl.
3. Determine the temperature of both the acid and the base to the nearest 0.1 ͦ
C. Average the temperatures. Record this average as the initial temperature:
Tinitial.
4. Pour the acid into the base quickly and carefully with gentle stirring. Start
monitoring the temperature as soon as the two are mixed. Continue to stir
40
and monitor the temperature. Record the maximum temperature that the
solution reaches. This is Tfinal .
5. Repeat Steps 1 - 3 (above) for Run 2.
6. Determine the temperature change for the reaction.
7. Calculate the heat absorbed by the solution. The heat lost by the
neutralization reaction will be equal to the heat absorbed by the solution but
opposite in sign. Using the volume and molarity the number of moles of
reactants can be determine and the ΔH, enthalpy of neutralization per mole,
can be calculated.
41
Experiment 8 Report
Determination of Molar Mass for a Volatile Liquid
Name: _________________
ID: __________________
Date:___________________
Instructor: _____________
Data Collected
Volume of 2.00 M HCl (ml)
Volume of 2.00 M NaOH (ml)
o
T initial
T final
C
o
C
Δ T = T final - T initial
o
C
Volume of resulting NaCl solution (ml)
Mass of resulting NaCl solution (density = 1.04 g/mL)
Specific heat capacity of the solution (3.87 J/g o C)
Calculations
1. Calculate the heat absorbed by the solution in each run. (q gained)
q = Heat gained by the solution = (s) (m) (ΔT)
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
2. Calculate q lost from the reaction (q lost)
q (lost) = ____________________________
42
3- Calculate the number of moles of water formed (same as moles of HCl used):
Moles H2O = Moles of HCl = M HCl x VHCl(L)
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
4- Calculate ΔH neutralization (ΔH°neut) per mole of water (kJ/mole) formed. Make
sure the sign is correct.
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
ΔH°neut = ___________________________
43
Experiment 9
Predicting shapes of molecules using VSEPR theory
Lewis Structures
A Lewis structure (or an electron dot formula) shows how the valence electrons are
arranged and indicates the bonding between atoms in a molecule. We represent the
elements by their symbols. The shared electron pair is shown as a line/bond
between the two atoms. All the other valence electrons are shown as dots or lines
around the symbol of the element.
For Example: The Lewis structure for F2 is
How to draw a Lewis Structure/Electron dot formula
a) Sum up the valence electrons of all the atoms
I. For anions, add one electron for each negative charge.
II. For cations, subtract one electron for each positive charge.
b) Choose the least electronegative element (other than H) as the central atom
For a molecule of type XYn, X is always the central atom.
c) Use single bonds to connect the central atom to the surrounding atoms.
d) Remember to subtract 2 electrons for each bond from the original total
number of electrons
e) Complete the octet for all outer atoms (other than H)
f) Complete the octet for the central atom
g) If you run out of electrons before the octet is complete for the central atom,
start forming multiple bonds till that atom has a complete octet
CO2
Examples:
H 2O
Molecular Geometry
The electron pairs around the central atom tend to be as far apart from each other
as possible. This is the main idea of the VSPER theory. We can apply the VSEPR
theory to predict the molecular shape/geometry of a molecule.
44
1. Draw the Lewis structure for the molecule in question.
2. Count the number of electron groups on the central atom. While counting
electron groups, each multiple bond is counted as a single electron group.
3. The arrangement of the electron pairs is determined by minimizing the
repulsions between them.
The attached table gives the relation between the number of electron pairs and the
molecular geometry.
No. Of
Electron
Pairs
2
3
4
Arrangements of Electron Pairs
Molecular Shape
(Geometry)
Examples
Linear
180°
Linear
BeCl2, CO2
Trigonal
planar
120°
Trigonal planar
BCl3
V-shape
NO2-
Tetrahedral
109.5°
(1 lone pair )
Tetrahedral
CH4
Trigonal
pyramidal
V-shape
(1 lone pair )
NH3
H2O
(2 lone pair )
5
Trigonal
Bipyramidal
120° & 90°
Trigonal
bipyramid
"see saw"
PF5
SF4
(1 lone pair )
T-shaped
ClF3
(2 lone pair )
Linear
XeF2
(3 lone pair )
6
Octahedral
90°
Octahedral
square pyramid
SF6
BrF5
(1 lone pair )
square planar
IF4(2 lone pair )
45
Experiment 9 Report
Predicting shapes of molecules using VSEPR theory
Name: _________________
ID: __________________
Date:___________________
Instructor: _____________
Make molecular models of the compounds listed in the table below and
complete the table.
Lewis
Molecular Polarity
No. of
No. of
Molecular
Structure
Shape
groups on lone
Formula
(Geometry)
central
pairs on
atom
Central
Atom
BeH2
CS2
H2S
SF6
BH3
NH4+
PCl5
CNSiCl4
46
47