CHEM 1412 Student Laboratory Manual (3)
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Texas A&M University *
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Course
1412
Subject
Chemistry
Date
Feb 20, 2024
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Experiment #1 – Safety and Laboratory Equipment. Objective:
To determine which laboratory equipment is the most accurate and precise. Background information & Theory:
Every piece of labware (whether made of glass or plastic) used to measure volume has some error associated with it. Some of the labware has a ± percent printed on it. For example a 100-mL graduated cylinder with a ±1 mL error on it means that if you measured 100 mL of a liquid it is really 100 ± 1 which is somewhere between 99 mL and 101 mL.
Equipment & Materials:
Name of labware
Quantity
beaker
1
volumetric flask
1
graduated cylinder
1
graduated pipet with bulb
1
Erlenmeyer Flask
1
Instructions: You will use each of the labware listed in the equipment & materials section to measure 100 mL
of deionized water three times. Pure water at room temperature should have a density of about 0.99796 g/mL
so use this to convert the volume into mass and weigh it on a 2 digit, 3 digit, or 4 digit analytical balance and
record the mass.
Data Section:
Labware
Mass #1
Mass #2
Mass #3
Average mass
beaker
volumetric flask
graduated cylinder
graduated pipet
with bulb
Erlenmeyer Flask
Labware
Average mass
Literature value
(in grams)
Percent error
Standard
deviation
beaker
99.7960
volumetric flask
99.7960
graduated cylinder
99.7960
graduated pipet
with bulb
99.7960
Erlenmeyer Flask
99.7960
Calculation Section: For this section calculate the average mass from all three measurements using each piece
of labware and then calculate the percent error from the literature value of 99.7960 grams and then calculate
the standard deviation.
Discussion Section: For this section you will discuss what you notice or observe in your data and
calculations. The labware with the lowest percent error should be the most accurate and the labware with the
lowest standard deviation should have the highest precision. See if you have a labware that has both the
lowest error and lowest standard deviation meaning it has the best accuracy and precision.
Conclusion Section: In this section you will make you conclusion based upon you data, calculations, and
discussions.
Experiment #2 – Effect of Solutes on the Boiling and
Freezing Points of Water Objective:
To observe the effect of adding a solute on the boiling and freezing points of a solvent. (the default solvent is water). Materials & Equipment:
Available in the Laboratory: Bunsen Burner Iron Stands and clamps DI Water (or any solvent preferred) From Prep Room Cart: Different Solutes (NaCl, sugar, vinegar) (Density of vinegar needs to be given) Thermometer Ice (stored in Ice Bucket) Salt to be added to the ice Copper Wire Loop with Handle (for Vertical Stirring) Large beaker (600 mL capacity) From Group Drawers: Beakers Test tubes
Procedure:
Part I – Effect of Solute on Freezing Point 1.
Prepare the freezing point apparatus as shown on
the illustration.
2.
Pure Water
Place exactly 20 mL of DI water into the test tube. As water freezes, continually stir the liquid for
form a “slush”. Monitor the freezing mixture such
that there is always a fresh supply of ice. Remove
the melted ice, if necessary. Note the temperature at which the solid and liquid
are in contact and the temperature stops to change. This is the freezing point. 3.
Water with Solutes.
A. Using NaCl as solute. Repeat step 2 above using another clean test tube. Before placing the tube into the freezing mixture, add a measured amount (1.0 – 2.0 grams) of NaCl and stir the mixture. Then cool the mixture as in step 2. Follow the rest of the procedure in step 2 and determine the freezing point of the mixture. B.
Using Sugar as solute. Repeat 3A, replacing NaCl with sugar C.
Using Vinegar as solute. Repeat 3A, replacing NaCl with a measured amount of vinegar.
Example Template for Data and Results:
Sample
Boiling Point
Freezing Point
Pure Water
A. Water with NaCl
Wire Stirrer
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Experimental Δt (from experiment)
The Theoretical Δt (from calculation,
using Kb and Kf of pure water from
textbook, and the concentration of the
solution)
% Error
B. Water with Sugar
Experimental Δt (from experiment)
The Theoretical Δt (from calculation,
using Kb and Kf of pure water from
textbook, and the concentration of the
solution)
% Error
C. Water with Vinegar
Δt
The Theoretical Δt (from calculation,
using Kb and Kf of pure water from
textbook, and the concentration of the
solution))
% Error
Experiment #3 – Chemical Kinetics (using Iodine Clock
Reaction) Objectives:
Part I: To determine orders of reaction and rate constant, for the reaction of Iodide with Hydrogen Peroxide to produce Iodine, and to write out the rate law. Part II: To demonstrate the applications of Arrhenius Equation and to determine the Activation Energy (Ea) using graphical analysis of results from Part I. Background/Concept
(Supplement to topics learned in Lecture)
Equation for the reaction; 2 H
+
(aq)
+ 2 I
-
(aq)
+ H
2
O
2(aq)
I
2
(aq)
+ 2 H
2
O (l)
(
#1)
The rate or speed of the reaction is dependent on the concentrations of iodide ion and hydrogen peroxide. Therefore, we can write a general rate equation for the reaction: Rate = k [I
-
]
a
[H
2
O
2
]
b
(
#2)
Starch is added to the reaction mixture in order to calculate the rate of reaction. As iodine is formed during the reaction, it will immediately react with the starch to produce a complex with a brown-
black color. This will tell you when the iodine is formed, but you have no idea as to the amount of hydrogen peroxide that has reacted in a given amount of time, i.e., the rate of reaction. So, thiosulfate ion, , is also added to the reaction mixture so that the following reaction will occur simultaneously with reaction (1). I
2
(aq)
+ 2 S
2
O
3
-2
(aq)
2 I
-
(aq)
+ S
4
O
6
-2 (aq) (
#3)
Thus, the iodine that is formed in reaction (1) is immediately transformed into iodide ion and we do not see the blue-black color of the starch-iodide complex until all of the thiosulfate ion has reacted with I
2(aq)
. When this occurs, we will then know the amount of hydrogen peroxide that has reacted and the time it took to react. For example, if you start with 5.0 x 10
-5
M of Na
2
S
2
O
3
in the reaction flask and you see a color change after 20 seconds, then you can say that 2.5 x 10
-5
M of I
2 (aq)
has been formed in 20 seconds and that 2.5 x 10
-5
M of hydrogen peroxide has reacted in 20 seconds. [Remember for every 2 moles of thiosulfate ion reacting, only 1 mole of iodine and, therefore, 1 mole of hydrogen peroxide are reacting. This is the stoichiometric ratio
from the two reaction equations, (3)
and (1)
.] The rate of disappearance of hydrogen peroxide is 2.5 x 10
-5 M divided by 20 seconds. This gives a value of 1.25 x 10
-6
M/second. ***
This is the value that will be recorded under REACTION RATE
on the data sheet. Note that the concentration of the I
-
(aq)
ion doesn't change since it is regenerated in reaction (3)
. In all of the reactions the amount of thiosulfate ion used will be kept small relative to the amount of hydrogen peroxide. Therefore, the concentration of hydrogen peroxide will change very little and you will essentially be working with initial concentrations of reactants and initial rates of reaction. You will do a total of three reactions at room temperature, one reaction at cold temperature (about 5
o
C), and one reaction at hot temperature (about 40 o
C, using a hot water bath). The three reactions at room temperature are labeled # 1, #2 , and #3. These make up Part I of the experiment. For the reaction #'s 1 and 2 the concentration of iodide is the same but the concentration of the other reactant, hydrogen peroxide, does vary. For the reaction #'s 1 and 3 the concentration of hydrogen peroxide is the same but the concentration of the other reagent involved in the rate expression, the iodide ion, does vary. The amounts, and therefore the concentrations, of the starch, sulfuric acid, and
sodium thiosulfate are constant for each reaction combination. In the reaction where the concentration of hydrogen peroxide (a molecular compound) is varied, the difference in the total volume of solution is compensated by the equivalent amount of another molecular compound , water. In the reaction where the concentration of iodide ion is varied, the difference in total volume is compensated by the equivalent amount of another ion of the same charge , the chloride ion. Effectively the total volume of the reaction, and the total ionic strength of the substances present is maintained at a constant value. After you have determined the reaction rate for each of the three reactions at room temperature you will substitute the values of the reaction rate, [
I
-
], and [
H
2
O
2
] for reaction #1 into equation (2)
. Then divide this entire equation by substituting the values for the reaction rate, [
I
-
], and [
H
2
O
2
] for reaction #2 into equation (2)
. This will give you a ratio of the reaction rates for reactions #1 and #2 on the
left side; and a ratio of the products of k
, [
I -
]
a
, and [
H
2
O
2
]
b
for reactions #1 and #2 on the right side. The
values of
k
, and [
I
-
]
a
will cancel on the right side since they are the same for both reaction #'s 1 and 2. This gives you a number on the left equal to [
H
2
O
2
]
b
/[H
2
O
2
] b
on the right. The ratio on the right can be rewritten as
{[H
2
O
2
]/ [H
2
O
2
]}
b or a single number raised to the power "
b
". The value of " b
" can now be calculated. Now repeat the calculations for the two reactions where [H
2
O
2
] is kept constant and determine the value for "
a
". You can now calculate the value of "
k
" for each of the reactions since you have the values of "
a
" ,
"
b
" , [
I
-
], [
H
2
O
2
], and the reaction rate
. The average of these three values is the value of "
k
" at room temperature. Part II
The results of the three reactions run at different temperatures (cold, room temp and hot) make up Part II
of the experiment. The mathematical relation, the Arrhenius equation
, relates the rate constant , k
, and the temperature in the following equation: ln k = ( Ea/R) (1/T) + ln A (
#4)
where Ea
is the activation energy for the reaction, T
is the temperature in o
C
, k
is the rate constant at T
, and R
is the ideal gas constant in units of J/mol
}
K and having a numerical value of 8.31. The equation is related to that of a straight line: y = mx + b
. If "
ln k
" is plotted against "
1/T
", then the slope of the resulting line will equal
-Ea/R
. The activation energy can be obtained by this method. LABORATORY PROCEDURE
: Glassware and Instruments
3 - 100 or 150 mL beakers 3 - 125 or 250 mL Erlenmeyer flasks 100 mL graduated cylinder 10 mL graduated cylinder watch with a second hand or stopwatch (digital timer) Thermometer pneumatic trough with ice water bath set at~ 40
C Chemicals (Be aware of the actual concentrations used)
Try to get only enough.
0.050 M KI
0.050 M KCl
0.010 M Na
2
S
2
O
3
0.050 M H
2
O
2
keep
covered
1.OM H
2
SO
4 (Spill : B1)
1% starch
solution
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Disposal: All mixtures should be poured into a Waste Bottle. NOTHING poured into the sink.
(Pure/unused solutions should be properly returned) Procedures
Part I
(room temperatures) 1.)
Label three clean and dry Erlenmeyer flasks from 1 to 3. Do the same for three clean and dry beakers. Make up the solutions given below. Measure the solutions as accurately as possible by using graduated cylinders of the appropriate size. The graduated cylinder must be thoroughly rinsed before using with a separate solution. (Obtain and measure H
2
O
2
when ready to mix; properly covered before pouring into the mixture). FLASK 0.050 M 0.050 M #
KI
KCl
0.010 M 1 M Na
2
S
2
O
3
H
2
SO
4
1% Star
ch
1
30.0 mL
5.0 mL
10.0 mL
10.0
mL
2
30.0 mL
5.0 mL
10.0 mL
10.0
mL
3
15.0 mL
15.0mL
5.0 mL
10.0 mL
10.0
mL
BEAKER#
0.050 M H
2
O
2
Deionized water
1
30.0 mL
2
15.0 mL
15.0 mL
3
30.0 mL 2.)
Add the contents of beaker #1 to flask #1. Immediately start your stopwatch as the two are added together. Continuously swirl the flask to mix the contents
. When the solution turns (dark) blue\black record the time and the temperature of the solution on the Data Sheet
. Do exactly the same for the remaining 2 pairs of solutions. If a beaker or a flask is to be reused it must be thoroughly rinsed with deionized water. The temperature of the reactions should remain within 1 o
C of each other. The KCl solution is used so that the ionic strength of the various solutions may be kept at a fairly constant level. On the
Data Sheet
, record the initial concentrations of hydrogen peroxide and iodide ions in the reaction solution. Remember that the original H
2
O
2
and I
-
solutions were diluted when they were mixed together. 3) Use the following equations to determine the new concentrations and the reaction rates: [( Molarity individual substance] = __{( Original Molarity ) ( Volume use in mL
)}
__ Total Volume of Mixture
* For: example:
[H
2
O
2
] =_ (Molarity) x (mL used)
Total mL of mixture Rate of Reaction
= Rate of formation of Iodide = ½ the Rate of Usage of Na
2
S
2
O
3 =1/2( concentration of Na
2
S
2
O
3 in reaction divided by time) *Note that the final (dilute) solution will contain 85 mLs in all cases
.
Part II
(cold and hot temperatures) 1. Prepare two flasks, each containing the same solutions as in flask #1
of part I
. Then prepare two beakers, each containing the same solution as in beaker #1 of
part I
.
2. Place one flask and one beaker in a pneumatic trough containing water at
40 C.
3.
Place the other flask and beaker in another pneumatic trough containing water at approximately 0
C.
4.
When the solutions have reached the temperature of the surrounding water, mix the appropriate solutions. Again, record the time that the solutions were mixed,
the time that the brown-black color appears, and the temperature of the reaction solutions.
5.
Prepare a graph for with x = 1/T and y = ln K. Draw a line of best fit and solve for the slope. From Arrhenius equation, the slope = - Ea/R. Determine Ea for the experimental reaction, Sugested format for Data & Results Part I
Reaction
Number
Initial
[H
2
O
2
]
(Refer to
Page 3 for
Calcs)
Initial
[ I
-
]
(Refer to
Page 3 for
Calcs)
[Na
2
SO
3
] /2
Reaction
Time
(seconds)
Rate
(Refer to
Page 3
for Calcs)
Exact
order
value
Rate
Constant
(K)
1
2
3
Orders of Reaction: ________ for H
2
O
2 and _______ for I
- Rate Law: Part II
Reaction Number
Description of
Condition
Actual
Temp (˚ C)
Actual
Temp (˚ K)
Rxt Time (sec) Reaction Rate 1 Cold
2 (use #1 results
from part I)
Room Temp
3 Hot
Combined Data
(Needed for Graphical Analysis) Data for Graphing:
Reactio
n
Numbe
r
Rate
Constant
(k)
ln k
(y-
axis)
T (K)
1/T (K) (x-
axis)
Comments
1
Cold Temperature
2
Room Temperature
3
Hot Temperature
Activation Energy from graphical analysis: __________________ : Prepare a Report following all guidelines
.
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Experiment #4 – Chemical Equilibrium DETERMINATION OF K
c
VALUES USING UV ABSORPTION I.
Objectives:
To determine the equilibrium constant for a given reaction.
To understand the concept of Le Chatelier’s Principle
To gain experience in the use of a UV Spectrophotometer and the Beer’s Law
Graph II.
Background/Concepts (Topics be researched by students and include literature in
Report)
Equilibrium and equilibrium constant
Le Chatelier’s Principle
The principles behind the use of a “spectrophotometer” (including the Beer’s
Law) III.
Procedure
A.
Calibrate of the spectrophotometer using a 2M solution of HNO
3
. (Be sure to use the same kind of cuvette or vial as those used for the remainder of the procedure) (Note: Use 450 nm as wavelength setting) Refer to page 3 of this procedure.
B.
Preparation of Solutions for the the Beer’s Law Graph 1.
To a clean 250-mL calibrated beaker, add the following: 10 ml of 0.0020 M KSCN 25 ml of 2.0 M HNO
3
65 ml of DI Water (Total volume at this point is 100 ml) 2.
Add to the solution in the previous step, 1.0 ml of 0.10 M Fe(NO
3
)
3 in HNO
3
solution. Mix with a clean stirring rod. Total volume at this point is 101 ml 3.
Fill a clean cuvette halfway from the rim with the solution prepared
above. (Be sure to handle the cuvette at the top portion so as not to leave any
finger prints on the lower portion of the tube.) 4.
Place this cuvette (from step 3) in the sample holder of the spectrophotometer and measure the % transmittance (and/or absorbance). 5.
Return the solution in the cuvette to the original reaction solution in the 250 ml beaker. 6.
Add another 1 ml quantity of the 0.10 M 0.10 M Fe(NO
3
)
3 / HNO
3 to the 250 ml beaker with the original solution. Again stir the solution. The total volume at this point is 102 ml.
7.
Repeat Steps 3-6 until ten (10) 1-ml additions and ten (10) transmittance readings have been recorded. 8.
Prepare a graph by plotting the X and Y values as determined from the succeeding calculations. The slope of the plotted line is -K
c
. 9.
Perform needed calculations and draw your conclusions. C.
Determination of K
c
for the reaction: FeSCN
+2
Fe
+3
+ SCN
-1
Use the results from the plotted values
. (Kc = - slope of best fit line) IV.
Suggested Chart for Data and Results To have values fit the cells, convert all values
to/proper scientific notation so that every value in each row or line will have
consistent scientific notation.)
Line 1 2 3 4 5 6 7 8 9 Solution No.
1
2
3
4
5
6
7
8
9
1
0
Measured
Quantities
%
Transmittanc
e
Absorbance
(2 – Log T)
[Fe
+3
]
M
[SCN
-
] M
([Fe
+3
]
i +
[SCN
-
]
i
) ([Fe
+3
]
i x
[SCN
-
]
i
) A{([Fe
+3
]
i +
[SCN
-
]
i
) A([Fe
+3
]
i +
[SCN
-
]
i
)
X =
-------------------
------
([Fe
+3
]
i x
[SCN
-
]
i
)
A
Y =
-------------------
-----
([Fe
+3
]
i X
[SCN
-
]
i
)
Slope
A
A{([Fe
+3
]
+ [SCN
-
])
----------------------- = -
K -------------------------- + b (
Algebraic equation for
the reaction)
([Fe
+3
]
[SCN
-
]) ([Fe
+3
]
[SCN
-
])
Y = m X + b (
Slope-intercept form of the line)
V.
Calculations and
Discussion of Results (Include actual graphical representation)
VI.
Conclusion (Address all issues of the objectives and discuss your results, in terms
of concepts and principles—not the procedure) -- Show and explain how the K
c
was obtained. -- Provide a short discussion why was it possible to use UV Absorption method for this experiment. Additional Information:
CALIBRATING THE SPECTROPHOTOMETER:
1)
Turn the power “ON” and allow the spectrophotometer to warm-up for at least 20 minutes. 2)
Set to required wavelength (be sure the wavelength range capability is also set) 3)
Be sure the setting is on “Transmittance Mode”
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4)
With the sample compartment empty and the cover closed, adjust Zero Control so that the meter reads “zero”. 5)
Fill a blank cuvette (up to 2/3 or ¾ mark) with reference liquid (the required solvent for the experiment. In this case it is HNO
3
solution). Place the cuvette with the solvent in the sample compartment and close compartment. 6)
Set the Transmittance to 100% using the transmittance control knob. 7)
Carefully open the sample compartment and remove the calibrating cuvette. 8)
The spectrophotometer is now ready for the experimental sample ------------------------------------------------------------------------------------------------------ Line # 3
: Molarity of Fe(NO
3
) 3
( Vol. of Fe
+3
soln used) [Fe
+3
] = Total Volume of Reaction Solution Line #4: [Molarity of SCN
-
] (Vol. Of SCN
-
soln. Used) [SCN
-
] = Total Volume of Reaction Solution
Experiment #5: Chemical Equilibrium – Part 2
Le Châtelier’s Principle
Required reading:
Ebbing, 11th Edition Chapter14.
Le Chatelier’s Principle
Effect of adding or removing products or reactants from an equilibrium.
Effect of the common ion.
Effect of the change the pH of an equilibrium.
Effect of changing the temperature to an equilibrium.
Learning Goals:
To understand how to manipulate equilibrium reactions.
To evaluate the validity of the Le Chatelier’s principle.
To determine whether a reaction is endo or exothermic based on qualitative experimental results.
Background information and theory
:
The equilibrium constant
, K,
is determined as a ratio of the product of the equilibrium concentrations
of the products raised to their stoichiometric coefficients divided by product
of the equilibrium concentrations
of the reactants raised to their stoichiometric coefficients. Equation 2 shows the expression to calculate K
for the generic equilibrium shown here:
a A
(
aq
)+
b B
(
aq
)
⇄
cC
(
aq
)+
d D
(
aq
)
(1)
K
=
(
[
C
]
eq
)
c
∙
(
[
D
]
eq
)
d
(
[
A
]
eq
)
a
∙
(
[
B
]
eq
)
b
(2)
Where lower case letters represent the stoichiometric coefficients and capitalized letters represent hypothetical chemicals involved in the reaction. A K
value which is > 1 indicates a
product dominate system, and one < 1 indicates a reactant dominant system. It is also possible to calculate a similar ration of products to reactants using values of concentrations under non-equilibrium conditions. This calculation is known as the reaction quotient
,
Q
, and its value can help identify whether the reaction is still on its way
to an equilibrium condition before or after a potential perturbation to such system. The expression to calculate this quotient for the generic equilibrium above would look very similar to that used to calculate K
.
Q
=
(
[
C
]
)
c
∙
(
[
D
]
)
d
(
[
A
]
)
a
∙
(
[
B
]
)
b
(3)
•
If Q
=
K
, then the system is at equilibrium. •
If Q
<
K
, then the system is still favoring the formation of products on its way to reach the equilibrium.
•
If Q
>
K
, then the system is favoring the formation of reactants on its way to reach the equilibrium.
Perturbing an Equilibrium System
Once a system achieves equilibrium, it will remain at equilibrium until an outside stress alters its relative concentration of products and reactants. There are several ways by which
an equilibrium system can be disturbed. This can be by directly changing the concentration of products and/or reactants, or, for a system which is in a gas phase, by altering the volume or pressure of the reaction system, or, changing a change in the temperature of the system. The Le Chatelier’s principle
states that any system at equilibrium will always find its way to re-establish equilibrium conditions after a perturbation has been caused to the initial equilibrium state. This principle helps scientists predict how any equilibrium system will behave in order to re-establish an equilibrium condition.
Changes in concentration of either reactants or products, a simple analysis of equation 3 would help predict what the system will do. For the generic equilibrium from equation 1:
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a A
(
aq
)+
b B
(
aq
)
⇄
cC
(
aq
)+
d D
(
aq
)
Adding moles of C or D: Q > K, reaction shifts left
Adding moles of A or B: Q < K, reaction shifts right
Removing moles of C or D: Q < K, reaction shifts right
Removing moles of A or B: Q > K, reaction shifts left
Changes in volume and pressure
will depend on the relative values of the stoichiometric ratios of the gaseous
reactants and products. If the sum of the stoichiometric ratios of reactants equals the sum of the stoichiometric ratio of products, P and V changes have no effect. If the two are not equal, P and V changes will favor the side of the reaction with the smallest value of the sum of the gaseous stoichiometric ratios. Changes in temperature
of a system depend on whether the reaction is exothermic or endothermic. We can predict the effect temperature change has on the position of an equilibrium for exothermic and endothermic reactions by considering energy to be either a
product (exothermic) or a reactant (endothermic). This is done by looking at the thermochemical equation
of the reaction in consideration, as follows.
Endothermic (
Δ
H > 0):
heat
+
a A
(
aq
)+
bB
(
aq
)
⇄
c C
(
aq
)+
d D
(
aq
)
(4)
Exothermic (
Δ
H < 0):
a A
(
aq
)
+
bB
(
aq
)
⇄
c C
(
aq
)
+
d D
(
aq
)
+
heat
(5)
Once we identify which side of the reaction arrows the heat belongs in, then the analysis is identical to the one we would do when changing amounts of reactants. The Reaction
In this experiment, you will be exploring the behavior of one reaction system under different perturbations caused to its equilibrium condition. You will do this in order to determine the validity of Le Chatelier’s principle. All the reactants and products in question
in this experiment are aqueous, so volume and pressure changes will not be evaluated. Equation 6 shows the equilibrium system we will work with.
[Co(H
2
O)
6
]
2+
+ 4 Cl
-
⇄
[CoCl
4
]
-2
+ 6 H
2
O
(6)
Cobalt(II) salts form complex ions readily. Aqueous solutions of cobalt(II) form the complex
ion [Co(H
2
O)
6
]
2
, which is pink (as shown). If Cl
-
ions are present in sufficient concentration,
they will form a deep blue (as shown) solution of [CoCl
4
]
-2
. Because of this, we can expect that if the solution becomes bluer in color, the concentration of [CoCl
4
]
-2
increases; while if it becomes pinker, the concentration of [Co(H
2
O)
6
]
2+
.
Aside from evaluating the behavior of this Co
2+
equilibrium upon changes to the amounts of
reactants, we will also analyze the temperature effects of the reaction. We should expect the same color changes to occur in solution by temperature effect as we would by changing concentrations. At the end of the experiment, students should be able to identify if the reaction is endothermic or exothermic by exploring the color changes in solution occurring when heating and cooling samples of the equilibrium mix.
Name: ____________________________________
Procedure:
Part 0: Predictions.
Got to the Data Sheet page below and fill up tables 1 and 2 with your predictions in regard to the experimental procedure for parts 1 and 2 in this experiment.
Part 1: Cobalt (II) Complex Equilibrium: Concentration Effects
For this part, make sure to take notes (on table 3 in the Data Sheet section) of your experimental results and observations. 1)
For the equilibrium reaction shown above, make your predictions regarding what you should expect to observe when adding: 1) concentrated HCl to the equilibrium mix 2) pure NaCl to the equilibrium mix, and 3) H
2
O to the equilibrium mix. 2)
Obtain three clean 13x100 mm test tubes. Fill each test tube with approximately 3mL of the violet colored cobalt equilibrium mixture. Caution: strongly acidic solution. Take proper precautions. 3)
To test tube 1: add 15 drops (dropwise) of concentrated hydrochloric acid to the mixture.
4)
Carefully shake the test tube until the acid is fully mixed in solution and observe the
color change.
5)
To test tube 2: add a small amount of NaCl (s) to the mixture (roughly 0.30 g – 0.50 g approx.).
6)
Carefully shake the test tube until the salt is fully dissolved in solution and observe the color change.
7)
To test tube 3: add 1 mL of water to the mixture and shake.
8)
Carefully shake the test tube and observe the color change. 9)
Record your observations and establish your conclusions. Think on whether they match your initial predictions and why/why not? (no need to answer this anywhere).
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Part 2: Cobalt (II) Complex Equilibrium: Temperature Effects
For this part, make sure to take notes (on table 4 in the Data Sheet section) of your experimental results and observations. 1)
Obtain one clean 15x150 mm pyrex test tube. Fill it about ½ full with the cobalt (II) equilibrium solution from part 1. 2)
Obtain two medium sized beakers. Fill each about ½ full with tap water. Heat one beaker on a hot plate, and add ice to the other, preparing a hot water and ice water bath (any temperature between 50°C and 80°C should be more than enough for the warm bath). 3)
Immerse the test tube in the hot water bath and hold until the color no longer changes
. 4)
Pull the test tube out of the warm bath and immerse it in the ice water bath until the color no longer changes
. 5)
Record your observations and reflect on what these results mean. 6)
You can repeat steps 3 and 4 as needed in order to further analyze the phenomenon.
7)
Is the reaction endothermic or exothermic? Justify your answer.
Name: ____________________________________
Le Châtelier’s Principle Data Sheet
:
Part 0: Predictions
1.
Please complete this table before watching the video demonstration. Your job is to predict what you think you will observe and then briefly explain your prediction in reference to the equilibrium mix between the hexaaquocobalt (II)
and the trtrachlorocobalt (II)
complexes, as shown again below:
[Co(H
2
O)
6
]
2+
+ 4 Cl
-
⇄
[CoCl
4
]
-2
+ 6 H
2
O
Table 1. Predictions to changes in concentrations.
Test
Tube
Perturbation
caused by
What do you think
you will observe?
Why?
#1
Conc. HCl added
to equilibrium
mix.
#2
NaCl(s) added
to equilibrium
mix.
Chem 1412 Page 20
of 36
Test
Tube
Perturbation
caused by
What do you think
you will observe?
Why?
#3
H
2
0(l) added to
equilibrium mix
2.
Fill up the following table based on the hypothetical change in enthalpy shown in the first column. This still refers to the cobalt (II) equilibrium mix.
Table 2. Predictions to changes in temperature.
Enthalpy
Change
Perturbation caused by
What should you observe?
Δ
H > 0
A rise in temperature
of the equilibrium mix
Δ
H > 0
Cooling of the
equilibrium mix
Δ
H < 0
A rise in temperature
of the equilibrium mix
Δ
H < 0
Cooling of the
equilibrium mix
Chem 1412 Page 21
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Part 1: Changes in Concentrations.
3.
Complete this table with the observations you make from “Part 1” of the experiment’s video and give a brief explanation to each behavior.
Table 3. Observations to changes in concentrations.
Test
Tube
Perturbation
caused by
Observations
Explain
#1
Conc. HCl added
to equilibrium
mix.
#2
NaCl(s) added
to equilibrium
mix.
#3
H
2
0(l) added to
equilibrium mix
Chem 1412 Page 22
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Part 2: Changes in Temperature.
4.
Complete this table with the observations you make from “Part 2” of the experiment’s video and give a brief explanation to each behavior.
Table 4. Observations and analysis to changes in temperature.
Perturbation caused by
Observation
A rise in temperature of
the equilibrium mix
Cooling of the
equilibrium mix
5.
Is this reaction exothermic or endothermic? Explain.
Chem 1412 Page 23
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Experiment #6
PREPARATION AND STANDARDIZATION OF A SODIUM HYDROXIDE SOLUTION Objectives
To prepare a solution (of required concentration) of sodium hydroxide
To review the principles
of titration
To standardized solution of sodium hydroxide reagent.
Part 1 – using a solid acid as primary standard Part 2 – using a standardized acid as primary standard
PRE-LAB ASSIGNMENT: Perform a Background Research on: (1)
Titration, (2) Standard Reagent, (3) Primary Standard, (4) Preparation of solutions of required concentration (Chapter 12) Procedures
Preparation of Solution 1)
Obtain the needed sample of sodium hydroxide. (Use the analytical balance to get the exact mass of the sample, as determined in the pre-lab assignment.) 2)
Pour the sample into a clean beaker and add enough DI water (~100 mL) to make 100 mL solution. 3)
Use a Stirring Plate for mixing. Make sure that the stir bar is clean. 4) Use this solution for titration and standardization. Part 1: Standardization of NaOH using a solid Acid as primary standard. 1)
Obtain a clean buret
and set-up for proper titration (be sure to use proper buret stand and buret clamp). Run DI water thru the buret. 2)
Run a small amount (less than 5 miL) of NaOH solution (that your group prepared)
thru the buret. Have a container ready for the washings and pour this out into a waste bottle. 3)
Fill the buret with the NaOH sample solution, and calibrate for proper measurement of volume. Make certain that the buret tip is filled with the solution. Keep the stopcock closed when buret is not in use. Note down your initial buret reading. Chem 1412 Page 24
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4)
Preparation of the primary standard: 5a) Obtain a stock sample of approximately 3.0 grams
of KHP. (use an evaporating dish or a weigh boat). 5b) Using an analytical balance, obtain the mass of a clean & dry 125-mL Erlenmeyer flask to the nearest 0.0001 gram. 5c) Remove flask from balance and carefully transfer your stock sample of KHP (
close to 3.000 grams
) to the flask. 5d) Re-weigh the flask with the KHP sample (using the same balance) Record all mass measurements. 5e) Add about 40 mL of DI water to the flask and stir until the KHP is dissolved (use
a stirring plate, if necessary) 5f) Add 2-3 drops of phenolphthalein indicator. 5)
Position the flask with KHP on top of a stirring plate and under the tip of the buret. Titrate the KHP solution with the NaOH in the buret by adding the titrant slowly (1 to 2 mL portions at a time). When the pink color becomes more observable, add the titrant one drop at a time until the pale pink color becomes permanent. 6)
DO NOT OVER TITRATE. The titration is complete when a pale pink color
persists for a minimum of 30 seconds. 7)
Record the final buret volume reading . 8)
Repeat the experiment (Steps 1-8)for a second trial using a similar amount of KHP. 9)
At the end of the second trial, return any remaining NaOH “stock solution” into the original container. Make sure it is not contaminated. Part 2: Standardization of NaOH using a standard HCl solution as Primary Standard. Follow the same procedure as Part I, except changing Step (4) by replacing KHP with 40.0 ml of standard 1.0 M HCl provided by the Instructructor. 10)
Clean the buret and return to the Laboratory Cart. 11)
Perform all required calculations (outside of the lab) and determine the average concentration of the NaOH sample solution. Chem 1412 Page 25
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Suggested format for Data and Results
Description
Trial 1
Trial 2 Mass of flask with KHP
Mass of flask only
Mass of KHP
Moles of KHP
Initial buret volume reading
Final buret volume reading
Volume of NaOH used
Molarity of NaOH
Average Molarity of NaOH
(Submit a technical report showing all pertinent calculations). Chem 1412 Page 26
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Experiment # 7
ACID-BASE TITRATIONS & TITRATION CURVES Pre-Lab
From the lecture series, a titration curve
is prepared by plotting the pH of the titration mixture versus the volume of the titrant
added. When the points are traced properly, the resulting curve will look like the letter “S” or the mirror image. The equivalence point
corresponds to the midpoint of the steep portion of the curve. This can be graphically located by drawing a vertical
line dividing the curve into equal parts. The pH that corresponds to this point (where the vertical
line and the curve intersect) is the pH at equivalence point. Objectives:
To demonstrate the titration
of acid with a base in different combinations (strong and weak)
To prepare titration curves for four different combinations of acid-base titration. Part 1 (Strong Base) Titrant 1- A: Titrating a Strong Acid (HCl) with a Strong Base (NaOH) 1.
Calibrate the pH meter by placing the electrode into a buffer solution (of known pH) and adjusting the knob to the pH of the given buffer. Then lift the electrodes and rinse with DI water before using in the succeeding steps. 2.
Obtain a clean buret and rinse with DI water. Then clamp the buret to a ring stand. 3.
Obtain a stock solution of standard NaOH solution (approximately 100 mL placed in an Erlenmeyer flask). Using a beaker, fill the buret with the NaOH solution, just above the calibration mark. Drain a small amount into a beaker until the meniscus is at a point within the zero mark (or any chosen calibration mark for initial reading. 4.
Obtain 35 mL of approximately 0.10 M HCl solution in a clean, dry200-mL beaker. 5.
Position the beaker with the acid on top of a magnetic agitator so that the pH electrode is immersed in the solution. Note the initial pH of the acid. 6.
Then, place the tip of the buret with the base titrant just inside the beaker, with the tip below the rim. 7.
Titrate the acid by adding the base from the buret in 2-4-mL increments, stirring continuously and, noting the pH after each addition. (Note: When a weak base is involved, Chem 1412 Page 27
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use your discretion on the volume of titrant to be added… sometimes 6-8 mL increments may be necessary) 8.
When successive pH readings begin to increase rapidly, add the titrant in smaller increments (1 mL or in drops) until successive pH readings are increasing only slightly. Gradually increase the amount of base added (not to exceed 2 mL at a time), until the pH reaches about pH = 10 (or over), or when the letter “S” is formed by the points. 9.
Prepare a titration curve by plotting the volume of base added as the abscissa and the corresponding pH as the ordinate. Determine the equivalence point for each titration. Part 1-B: Titration of a Weak Acid (acetic acid) with a strong base (NaOH) – Repeat Steps 1
– 10 in Part A, replacing HCl with acetic acid. Part 2 – Using Weak Base as Titrant
Part C: Titration of Strong Acid with Weak Base •
Use Ammonium hydroxide to titrate HCl. Follow the same procedure as Part 1-A Part D: Titration of Weak Acid with Weak Base •
Use Ammonium hydroxide to titrate Acetic acid) Follow the same procedure as Part 1-B . Data and Results
(Suggested Chart on next page) (Add more spaces when needed
) Note
: when the pH reading does not change for more than 5 additions of the base, there is no need to record all readings. Do not forget to use proper Waste Disposal. Part A (HCl w/ NaOH) (Identify specific samples
used) Part B (CH
3
COOH w/ NaOH) (Identify specific samples
used) Part C (NH4OH w/ HCl) (Identify specific samples
used) Part D (NaOH w/ HCl) (Identify specific samples
used) mL Base (total) pH mL Base (total) pH mL Base (total) pH mL Base (total) pH Chem 1412 Page 28
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Chem 1412 Page 29
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Experiment #8 The Solubility Product-Constant (Ksp) I.
Objectives:
To calculate the K
sp
values of some salts II.
Background:
Research and provide definition/principles for:
K
sp
and how to obtain its value I
Procedure
.
Before performing the actual procedure, prepare a perfectly tarred evaporating dish for consistent “dry” mass by heating it, cooling and weighing. Heat again, cool and weigh until the difference in masses is = or < 0.0010. (Note: Use the same analytical balance and DO NOT touch the dish with bare hands.) 1)
(Preparing a saturated solution) 1-a. In a large test tube, mix a few grains of the first salt with 20 mL of water
and stir or agitate. 1-b Continue adding the salt a few grains at a time (with stirring or agitation) until the added salt no longer dissolves. Let the solution stand for a few minutes 1)
Withdraw 5 mL of the supernatant liquid
using a pipette and place this volume in a pre-weighed
clean and very dry evaporating dish. 2)
Cover with a watch glass and evaporate gently using an air bath over a Bunsen Burner or a hot plate on high setting. 3)
Cool and re-weigh the evaporating dish to obtain the mass of salt
in the 5 mL saturated solution. 4)
Determine the concentration of the salt used in the saturated solution. 5)
Calculate the K
sp
.
Repeat Procedures 1 – 6 using a second salt. II
Suggested format for Data and Results (Tabulated)
Mass of Salt +
Evaporating Dish Avg. Mass of Evap. Dish (empty) Mass of
Salt Molar Solubility
Of Salt K
sp for
Salt I.
Calculations (All values that were not “measured” during the procedure should have pertinent calculations) Chem 1412 Page 30
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II.
Conclusions Page 1
of 1
Experiment # 9 SPONTANEITY Pre-Lab Assignment: Determine the mass of Sodium Nitrate that would be required to prepare 100 mL of 1.0 M solution. Chem 1412 Page 31
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Review the heat calculations from Thermochemistry (Chem. 1412), including the “heat balance” I.
Objectives:
To measure the enthalpy change that occurs when a salt is dissolved in water.
To estimate the minimum value for the entropy change.
To predict the sign of the “free-energy
” change for the process. Background Research and understand the definitions/principles/concepts for:
H (enthalpy) ,
S (entropy) and
G (free-energy)
Spontaneous and non-spontaneous processes & the Coffee-Cup Calorimeter. Be
sure to include the material that you have researched in your Lab Report.
Procedure 1)
Perform the Pre-Lab Assignment 2)
Set-up a “coffee-cup calorimeter.” (composed of 3 cups and prepare the set-up according to sample shown by instructor from lecture slides) 3)
Using a weighing container, obtain the needed amount of Sodium Nitrate Sample. 4)
Place 100 mL of distilled water in the calorimeter cup. 5)
Measure and record the temperature of the water (to the nearest fraction in
C.) This is the initial temperature (
t
i
). 6)
Add the solid sodium nitrate to the cup in such a way that none adheres to the side of the cup. 7)
Place the top (or lid) to the cup immediately and begin mixing or stirring. 8)
Measure the temperature of the solution to the nearest fraction of
C after 30 seconds and every 30 seconds thereafter, until the temperature attains either a maximum or a minimum value. This temperature is the final temperature (
t
f
).
Note: q of the system
= -
q of the environment
The environment consists of water and the coffee/Styrofoam cup. 9)
Calculate q for the environment, using 4.184 J/g
C as specific heat of water, and 1.0 g/mL for density of water. Assume the heat capacity of the coffee-cup as
10.0 J/
C. 10)
Calculate the enthalpy change (
H) using “q” and the number of moles of NaNO
3
. 11)
Repeat the same procedure for a second trial. 12)
Calculate the mean enthalpy change for the process. 13)
Figure out the approximate minimum value of
S and predict the sign for
G .
(Note: this can be done by using only the signs of
H and
S, and the temperature at which the reaction was performed. (Refer to Lecture Notes) 13)
Complete a lab report following all guidelines Chem 1412 Page 32
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Experiment #10 ELECTROCHEMISTRY Pre-Lab Assignment: Know the two general types of Electrochemical Cells and learn how each one operates. Chem 1412 Page 33
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I.
Objectives:
To prepare a set-up containing both types of electrochemical cells.
To predict formula mass of the metal ,Lead, from the solid deposited on one electrode of the electrolytic cell. II.
Background (Know and understand the following topics. Research and include in Lab
Report))
Electrochemistry
Redox reactions
Voltaic Cell
Electrolytic Cell
Faraday Law The Double-Cell Set-Up for Electrochemistry: Chem 1412 Page 34
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. III.
Procedure
Obtain two lead strips
from the laboratory cart (or from your instructor). Sandpaper each strip to remove impurities from the surfaces. (Place a paper towel beneath the strip in order to collect the dust. Dispose of the dust particles) 1)
Using the analytical balance, get the mass of each strip to the nearest 0.001 g. Mark each strip and designate which one is the anode and which one is the cathode (for the electrolytic cell) 2)
Fill the porous cup
containing a zinc strip
with 1.0 M zinc sulfate solution
. 3)
Place this cup in an electrochemical beaker
(provided) containing about
120 mL of 1.0 M copper (II) sulfate solution
. Place a copper strip
into the beaker as shown in the illustration. Chem 1412 Page 35
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Zrive:\locruz\Chem 1412\lab procedures Expt #9, Page 1 of 2 4)
Place the lead strips into another electrochemical beaker containing about 150 mL of lead (II) nitrate solution. 5)
Switch on the ammeter (set to read to 250 mAmp, if not reading set to 50 mAmp) 6)
Connect the zinc, copper, and lead strips using copper wire connectors provided.
Refer to the illustration for proper connections.
Connect the wire from the zinc strip in the porous cup to one of the lead strips in the lead (II) nitrate solution. This strip will be the cathode
of the electrolytic cell.
Connect the other lead strip to the negative terminal of the ammeter
; this will be the anode
.
Connect the copper strip
(in copper (II) sulfate solution) to the positive terminal of the ammeter. 1)
Immediately begin recording the time on your Data Sheet. 2)
Read and record the current reading on the ammeter at intervals of 5 minutes. Allow the cell to operate for a minimum of 30 minutes. 3)
After the time of operation, disconnect one wire from the ammeter and record the time on the Data Sheet. 4)
Disconnect the two lead electrodes carefully and let them stand on a beaker. BE
SURE NOT TO SHAKE OFF ANY POSSIBLE DEPOSITS). Gently rinse them with acetone (using the squirt bottles provided). If any solid crystals fall off, filter the solution to recover the crystals. Allow the strips to dry at room temperature. 5)
Weigh each strip to the nearest 0.001 of a gram, and record the mass of both strips. 6)
DO NOT POUR OUT ANYTHING INTO THE SINK. Return all solutions to their respective containers except the Lead Nitrate
which should be pouredinto
the waste bottle. Pour any acetone washings into the waste bottle. 7)
Using the Faraday Law and stoichiometric equations, calculate the formula mass
of the lead deposits. (Tabulate your results and show ALL pertinent calculations in your Lab Report). Chem 1412 Page 36
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- 2. A student is told to determine the density of a solid unknown cylinder. The mass of the unknown cylinder is 56.38 grams. The unknown cylinder is 4.95 cm high and has a radius of 1.28 cm. The student measures the correct mass but measures the height as 4.96 cm and the radius as 1.30 cm. What will be the percent error in the student's density value? True- Experimental| x100 Percent Error = Truearrow_forwardgraduated cylinder weighs 35.825 g. When 10 mL of water is measured in it, the weight of the graduate and water is 45.835 g. Calculate the weight of the water, and express any deviation from 10g as percentage of error.arrow_forwardIt was still incorrect please solve it again.arrow_forward
- A student’s percent error for the density of gold was 7% under the accepted value of 19.3 g ml-1. What was the student’s experimental value?arrow_forwardplease make sure each answer has the correct number of significant digitsarrow_forwardWhat is the smallest mass that we can measure on an analytical balance that has a tolerance of ±0.1 mg, if the relative error must be less than 0.1%?arrow_forward
- A volumetric flask that is to be used to measure volume is supposed to be marked at 10.00 ml but is actually marked at 9.95 ml. This error affects its precision. True or false? please explainarrow_forward3. A student is using a 10 mL volumetric pipet to dispense water and wants to test how accurate the volume is by weighing the water dispensed from the pipet. First, they measure the mass of a beaker to be 30.045 ± 0.005 g. They then collect 10 mL of water into the volumetric pipet and dispense the water into the beaker. The mass of the beaker plus the water is 40.081 ± 0.005 g. Calculate the mass of the water added to the beaker and use propagation of error to calculate the uncertainty of the mass. Report the mass of the water with the uncertainty with appropriate significant figures.arrow_forwardCan you help me, please?arrow_forward
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