1
Experimental Procedures
General Chemistry I
KI-1101
Translated from
Penuntun Praktikum Kimia Dasar I KI-1101
by
Theodorus Felix Darpieto Abik, S.Si
Muhammad Arizki, M.Si
Dr. Rizqiya Astri Hapsari
GENERAL CHEMISTRY LABORATORIUM
FIRST YEAR PROGRAM
INSTITUT TEKNOLOGI BANDUNG
2 0 1 7
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Table of Content
Experiment I Chemical Reactions ............................................................................................................... 3
Experiment II Stoichiometry of Chemical Reactions ............................................................................. 7
Experiment III Energy Changes in Chemical Reactions .................................................................. 11
Experiment IV Chemical Bonding and Molecular Polarity .............................................................. 15
Experiment V Physical and Chemical Properties Of Gases ........................................................... 19
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Experiment I
Chemical Reactions
Introduction
A chemical reaction is a process involving two or more reactants to produce a product that has
physical/chemical properties that are different from its reactants. In general, chemical
reactions are divided into two groups, acid-base reactions and oxidation-reduction reactions.
Acid-base reactions are chemical reactions involving the neutralization of H+ and OH− ions
(Arrhenius theory), proton ions (H+) acceptor-donor (Bronsted-Lowry theory), electron pair
acceptor-donor (Lewis theory), or oxide ion (O2-) acceptor-donor. Oxidation-reduction
reactions are chemical reactions involving electrons transfer between reductor and oxidator,
followed with a change in the oxidation state. Some phenomenon that can be observed in a
chemical reaction, including: (i) the presence of gas as a reaction product; (ii) the presence of
precipitate; (iii) pH changes; (iv) color changes; or (v) temperature changes. Below are some
examples of chemical reactions:
(i) Oxidation-reduction reaction:
Gas formation: 2 Al(s) + 6 HCl(aq) → 2 AlCl3(aq) + 3 H2(g)
Purification of oxide ore: Fe2O3(s) + 3 CO → 2 Fe(s) + 3 CO2(g)
Analysis of qualitative/quantitative ethanol: 2 K2Cr2O7(aq) + 3 C2H5OH(aq) + 8
H2SO4(aq) → 2 Cr2(SO4)3(aq) + 3 HC2H3O2(aq) + 2 K2SO4(aq) + 11 H2O(l)
(ii) Acid-base reaction:
Neutralization: NH3(aq) + HCl(aq) → NH4Cl(aq)
Precipitate formation: AgNO3(aq) + Na2CrO4(aq) → Ag2CrO4(s) + 2 NaNO3(aq)
Thermal decomposition: CaCO3(s) → CaO(s) + CO2(g) (Occurring at 900 °C, the
oxide acceptor-donor, Ca2+ ion receives O2- ion from CO32- ion)
The main focus of this experiment is to study the chemical reactions that use water as the
solvent via observing any changes that occur in each reaction. After completing the
experiment, the student is expected to be able to: (i) recognize the used chemicals, (ii) write
down the chemical formula, (iii) write the chemical equation correctly, and (iv) recognize the
various kinds of chemical reactions.
Chemicals and Equipment
The chemicals required in this experiment are listed below:
0.1 M CuSO4, 0.1 M HCl, 0.1 M AgNO3, 0.1 M Pb(NO3)2, 0.1 M NaC2H3O2, 0.1 M KI, 0.1
M KOH, 0.1 M Na2CO3, 0.1 M NH3, 0.1 M HC2H3O2, 0.1 M K2CrO4, 0.1 M K2Cr2O7, 1 M
HCl, 1 M NaOH, 0.05 M KMnO4, 0.1 M H2C2O4, 0.1 M Fe(II), 2 M H2SO4, 3% H2O2,
CuSO4.5H2O solid and KI solid, Mg, Cu, and Zn metals.
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The equipment needed in this experiment are listed below:
Test tube, tube rack, pipette, and spatula.
Procedures
Section I: Metal Oxidation Reaction
a. Into a test tube, put 1 mL of CuSO4, and add a piece of Mg metal into the solution.
Observe the changes that occur at the beginning of the reaction and 5 minutes after the
reaction started.
b. Put 1 mL of HCl into a test tube, and add a piece of Zn metal into the solution. Observe
the changes that occur at the beginning of the reaction and 5 minutes after the reaction
started.
c. Into a test tube, put 1 mL of AgNO3, subsequently add a piece of Cu metal into the
solution. Observe the changes that occur at the beginning of the reaction and 5 minutes
after the reaction started.
d. Based on the observation results, for all of the above reactions, answer these following
questions:
Do they react spontaneously?
Write down the balanced chemical equation for each of the above reaction. Use standard
reduction potential data, E0, for each of the above reagent.
Section II: Pb2+ Ion Acid-Base Reaction
e. Into a test tube, put 1 mL of 0.1 M Pb(NO3)2, then into that solution, add 1 mL of 0.1 M
NaC2H3O2. Observe the changes that occur.
f. Put 1 mL of 0.1 M Pb(NO3)2 into a test tube, then add 2 mL of 0.1 M KI into the solution.
Observe the changes that occur.
g. Based on the observations of the two reactions above, write down the balanced chemical
equation for each of the above reaction.
h. Do those above reactions produce any precipitate? If yes, give an explanation why the
precipitate was formed. Ksp for PbI2 (25 °C) = 7.9 × 10-9 and the solubility of
Pb(C2H3O2)2 (20 °C) = 44.31 g/100 mL.
Section III: Cu2+ Ion Reduction Reaction in Solid- & Aqueous-Phase
a. Prepare four test tubes.
Tubes 1 & 2: put a little amount of CuSO4.5H2O solids. Then label them as A and B
Tubes 3 & 4: put a little amount of KI solids. Label them C and D.
b. Into tube C, pour down the solids in the tube A. Observe the changes that occur.
c. Into each tube B and D add 2 mL of water and then stir it until the solids is dissolved.
Pour the solution in tube B into tube D and then observe the changes that occur.
d. Based on the observations of step b and c, what is the difference between solid phase
(step b) and aqueous phase reaction (step c)?
e. Write a balanced chemical equation for each of the reaction.
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Section IV: Indicator Color Changes in Acid-Base Reactions
a. Into a test tube, put 1 mL of 0.1 M NaOH, subsequently add 2 drops of indicator solution
into the same test tube. Into the mixture, add 1 mL of 0.1 M H2C2O4. Observe whether
there is any color changes on NaOH solution after the addition of indicator and H2C2O4.
Explain your observation results.
b. Put 1 mL of 0.1 M NH3 (note: NH3 solution, not NH4OH) into a test tube, then add 2
drops of indicator solution into the solution. Into the mixture, add 1 mL of 0.1 M
H2C2O4.. Observe whether there is any color changes on NH3 solution after the addition
of indicator and CH3COOH. Explain your observation results.
c. Write a balanced chemical equation for each reaction.
d. Based on the acid/base strength, discuss the difference of reaction (a) and reaction (b).
Section V: Chromate (CrO42-) and Dichromate (Cr2O7
2-) Ions Equilibrium
K2CrO4 and K2Cr2O7 are oxoanions of Cr(VI), which dissolve well in the water. The
existence of each oxoanion, CrO42- and Cr2O7
2-, in solution is strongly affected by the pH of
the solution. The dissolved Cr2O72- ions in the water will produce an orange-colored solution,
while CrO42- ions will produce a yellowish solution. Note: Cr(VI) compounds are toxic, be
careful, don’t expose it to your skin. When exposed, immediately wash or rinse your skin.
a. Prepare two test tubes then fill each of the test tubes with 1 mL of K2CrO4. Into tube
1, add 5 drops of HCl and then gently shake the mixture. Observe whether the color of
the solution changed or not. Into tube 2, add 5 drops of 1 M NaOH and then shake the
mixture slowly. Observe whether the color of the solution changed or not. Keep both
of the reaction mixtures.
b. Do the same as above, but replace the K2CrO4 solution with K2Cr2O7 solution.
c. Compare the experimental results from subsection (a) with subsection (b). Determine
the pH (acidic or alkaline solution) of each reaction mixture that you made.
d. Write the equilibrium reaction of Cr2O72- and CrO4
2- ions under acidic and alkaline
environments.
Section VI: The Reduction Reaction of Hydrogen Peroxide
It has been known that the reaction of H2O2 with KI happens in two stages:
H2O2(aq) + I-(aq) → 2 H2O(l) + IO-
(aq) (i)
H2O2(aq) + IO-(aq) → H2O(l) + O2(g) + I-
(aq) (ii)
As can be seen from the above reactions, I- is presence at the beginning and end of the
reaction, which mean KI act as a catalyst for the reduction reaction of H2O2.
Do the following experiment in a fume hood.
Into a test tube, put 2 mL of 3% H2O2 solution then add a little amount of KI solid (about a tip
of a small spoon) into the solution. Observe the changes that occur. Are there any changes in
temperature and color of the solution?
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Section VII: Potassium Permanganate Reduction Reaction
Potassium permanganate, KMnO4, is a strong oxidizing agent that is widely used in chemical
reactions. Manganese (Mn) can form compounds, which vary widely in the oxidation
numbers, e.g., +2, +3, +4, +5, +6 and +7. Under acidic conditions, MnO4− ion can be reduced
to MnO42- ion (green solution), MnO2 (blackish brown solid), or Mn2+ (magenta solution)
depends on the reducing agent used in the reaction. Reducing agents, which can be used to
reduce MnO4- ion, including: Zn, H2C2O4, and Fe. This is related with the potential reduction,
E0, of KMnO4 with each reducing agent.
a. Into a test tube, mix 1 mL of 0.1 M H2C2O4 with 2 mL of 2 M H2SO4. Subsequently,
while shaken, drop wisely adds 0.05 M KMnO4 solution until the solution color is
changed. Note the time required for the KMnO4 solution color to change as well as the
amount of KMnO4 required.
b. Into a test tube, mix 1 mL of 0.1 M Fe(II) and 2 mL of 2 M H2SO4 and then, while
shaken, drop wisely adds 0.05 M KMnO4 until the solution color is changed. Note the
time required for the KMnO4 solution color to change as well as the amount of KMnO4
required.
c. Which reactions require a shorter time to change the KMnO4 color, reaction (a) or (b)?
Explain your observation results as well as your answer.
d. Write a balanced chemical equation for each reaction.
e. If one drop of KMnO4 solution is assumed to be equivalent to 0.05 mL, then count the
number of moles of KMnO4 required on each of the above reaction. Is the number of
moles of KMnO4 required in both reactions different? Give an explanation for your
results.
Please bring them together: • Lab Journal • Long sleeve lab coat • Goggle • A rag/ tissue
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Experiment II
Stoichiometry of Chemical Reactions
Introduction
Stoichiometric ratio of reactants is important to monitor a reaction quantitatively. A reaction’s
progress can be observed through changes in temperature, pH, color, or the amount of the
product formed, such as its mass (when the reaction includes precipitation) or gas volume
(when the reaction evolves gases). One of the most common methods to determine
stoichiometric ratio of a reaction is through Job’s Method, also known as the Method of
Continuous Variation. In this method, several reaction runs are made using different ratios of
the reactants while keeping the total molar amount of the reactants constant. For each runs, a
certain physical property (it may be changes in pH, temperature, mass, volume, or
absorptivity) is monitored, then correlated to the ratios of the reactants using a plot (Job’s
plot) to determine the reaction’s stoichiometric ratio.
For example, let’s observe the following table, showing data obtained to determine the
stoichiometric ratio for the reaction between silver nitrate (AgNO3) and potassium chromate
(K2CrO4). This is a precipitation reaction, therefore it is possible to monitor the weight of
precipitates formed for different ratios of silver nitrate and potassium chromate. A Job’s plot
is then constructed with respect to silver nitrate’s molar amount.
𝑅𝑢𝑛 𝑉𝐴𝑔𝑁𝑂3
(𝑚𝐿)
0,24 𝑀
𝑉𝐾2𝐶𝑟𝑂4 (𝑚𝐿)
0,24 𝑀
𝑇𝑜𝑡𝑎𝑙 𝑉𝑜𝑙𝑢𝑚𝑒 (𝑚𝐿)
𝑀𝑜𝑙𝑎𝑟 𝑎𝑚𝑜𝑢𝑛𝑡 (× 10−3 𝑚𝑜𝑙)
𝐴𝑔𝑁𝑂3 𝐾2𝐶𝑟𝑂4 𝑇𝑜𝑡𝑎𝑙 1 5 45 50 1.20 10.80 12.0
2 10 40 50 2.40 9.60 12.0
3 15 35 50 3.60 8.40 12.0
4 20 30 50 4.80 7.20 12.0
5 25 25 50 6.00 6.00 12.0
6 30 20 50 7.20 4.80 12.0
7 35 15 50 8.40 3.60 12.0
8 40 10 50 9.60 2.40 12.0
9 45 5 50 10.80 1.20 12.0
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𝑅𝑢𝑛 𝑀𝑜𝑙𝑎𝑟 𝑎𝑚𝑜𝑢𝑛𝑡 (× 10−3 𝑚𝑜𝑙)
𝑀𝑎𝑠𝑠 𝑜𝑓 𝑃𝑟𝑜𝑑𝑢𝑐𝑡 (𝑔𝑟𝑎𝑚) 𝐴𝑔𝑁𝑂3 𝐾2𝐶𝑟𝑂4
1 1.20 10.80 0.225
2 2.40 9.60 0.396
3 3.60 8.40 0.564
4 4.80 7.20 0.696
5 6.00 6.00 0.885
6 7.20 4.80 1.030
7 8.40 3.60 1.194
8 9.60 2.40 0.892
9 10.80 1.20 0.598
According to the Job’s plot that has been constructed, the stoichiometric ratio between
AgNO3 and K2CrO4 is 2:1, and the molecular formula of the resulted product is Ag2CrO4
(AgNO3 mole : K2CrO4 mole = 8.4 : 3.6 = 2.33 : 1 ≈ 2 : 1).
In this experiment, you will determine the stoichiometric ratio for three reactions: (i) between
CuSO4 and NaOH; (ii) between HCl and NaOH; and (iii) between H2SO4 and NaOH,
monitoring temperature change as the observable property. You will also determine yield
percentage for the precipitation reaction between lead(II) acetate (Pb(C2H3O2)2) and
potassium iodide (KI). Yield percentage is the ratio between the actual mass obtained to the
theoretical mass yield (made under the assumption that the reaction would go to completion).
%𝑦𝑖𝑒𝑙𝑑 =𝑎𝑐𝑡𝑢𝑎𝑙 𝑚𝑎𝑠𝑠
𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑚𝑎𝑠𝑠× 100%
0.000
0.200
0.400
0.600
0.800
1.000
1.200
1.400
0 2 4 6 8 10 12
Mas
s o
f P
rod
uct
s (g
ram
)
Moles of AgNO3 (10-3 mol)
Job's Plot, Mass of Products to Moles of AgNO3
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Chemicals and Equipment
The following chemicals are required for this experiment:
Solutions of Pb(C2H3O2)2 (0.1 M), KI (0.1 M), NaOH (1 M and 2 M), HCl (1 M), H2SO4 (1
M), CuSO4 (1 M), filter paper
The following equipment are required for this experiment:
Analytical scale, 50 mL measuring cylinder, 50 mL and 100 mL beaker glass, 100 mL
Erlenmeyer flask, funnel, thermometer, watch glass, oven
Procedures
I. Reaction between Pb(C2H3O2)2 and KI solutions
1. Put 2 mL of Pb(C2H3O2)2 0,1 M solution to 50 mL beaker glass, then add 2 mL of KI
0,1 M solution.
2. Put a filter paper on a watch glass, then measure the combined dry mass; write down
the measured combined mass on your note.
3. Create a filtering set up using the funnel, Erlenmeyer flask, and filter paper.
4. Pour the resulted solution and precipitate from step 1 to the filtering set up, aided by a
stirring stick. This step is to separate the precipitate from the remaining solution. Take
care not to spill anything.
5. Rinse the stirring stick and beaker glass using distilled water 3 times, pouring the
resulted solution into the filter. Make sure all precipitate has been filtered.
6. Once there are no more drips of filtrate from the funnel, put the filter paper containing
yellow solids on the watch glass, then put them inside 100oC oven for 30-45 minutes
or until it has dried completely.
7. Cool the dried solids, filter paper, and watch glass to room temperature.
8. Weigh the filter paper and watch glass once again, this time containing the yellow
solids; write down the resulting mass.
9. Calculate the theoretical mass for the resulting product when Pb(C2H3O2)2 and KI are
mixed, assuming the reaction goes to 100% completion. Then, calculate the yield
percentage of the reaction using the measured mass of the yellow precipitate.
II. Reaction between CuSO4 and NaOH solutions
1. Pour 50 mL NaOH 1 M solution into 100 mL beaker glass, then measure its
temperature.
2. Pour 10 mL CuSO4 1 M solution into 50 mL beaker glass, then measure its
temperature.
3. Pour the CuSO4 into the NaOH solution. Shake gently while measuring the mixture’s
temperature. Write down the highest temperature measured.
4. Repeat step 1-3 under the following compositions:
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Run Volume of NaOH
(mL)
Volume of CuSO4
(mL)
1 50 10
2 40 20
3 30 30
4 20 40
5 10 50
5. Plot a graph of T against volume of NaOH, with T as the change in temperature (T =
Tf – Ti, Tf = the mixture’s final temperature, Ti = the initial temperature of each
solution)
III. Acid-base reaction (between HCl and NaOH; H2SO4 and NaOH)
1. Pour 5 mL NaOH 1 M solution into 50 mL beaker glass, then measure its temperature.
2. Pour 25 mL HCl 1 M solution into 50 mL beaker glass, then measure its temperature.
3. Pour the NaOH into the HCl solution. Shake gently while measuring the mixture’s
temperature. Write down the highest temperature measured.
4. Repeat step 1-3 under the following compositions:
Run Volume of NaOH
(mL)
Volume of HCl (mL)
1 10 20
2 15 15
3 20 10
4 25 5
5. Plot a graph of T against volume of NaOH, with T as the change in temperature (T =
Tf – Ti, Tf = the mixture’s final temperature, Ti = the initial temperature of each
solution)
6. Repeat step 1-5, substituting HCl with H2SO4.
Please bring them together: • Lab Journal • Long sleeve lab coat • Graphical paper • Goggle • Calculator • Ruler • A rag/ tissue
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Experiment III
Energy Changes in Chemical Reactions
Introduction
Thermochemistry is a branch of chemistry that studies heat changes in a chemical reaction.
Heat changes involved in a reaction can be measured by simplify some of the reaction system
and reaction surrounding parameters. In this experiment, the heat changes are studied in a
constant pressures and only involves solid- and liquid-phase substances (volume changes can
be neglected). Therefore, the work relating to the reaction system ( 𝑤 = 𝑷 ∆𝑽), can be
ignored. Based on the first law of thermodynamics, the internal energy changes, E, which
accompanies the reaction in this kind of experimental conditions is same as the reaction
enthalpy changes, Hrx. Besides the law of Conservation of Energy, the Principles of Black –
the released heat equals to the absorbed heat – can also be used to solve problems in this
experiment. Through simplification of various reaction parameters, heat changes resulted
from a chemical reaction in a calorimeter can easily determined by measuring the temperature
changes within a reaction mixture.
In this experiment, we will determine the released of heat that comes from acid-base
neutralization reaction between hydrochloric acid (HCl) and sodium hydroxide (NaOH) in
two different conditions:
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) ∆Hrx1 = ?
HCl(aq) + NaOH(s) → NaCl(aq) + H2O(l) ∆Hrx2 = ?
The enthalpy of solution reaction of sodium hydroxide, NaOH(s), in the water will also be
determined in this experiment,
NaOH(s) → NaOH(aq) ∆Hrx3 = ?
Determination of heat changes in those three reactions can be calculated using the Hess’s
Law. Hess's Law states that the enthalpy changes of a whole process is the sum of the
enthalpy changes of each reaction stage or in other words, the enthalpy of a chemical reaction
is independent of the path taken from the initial to the final state. Note that when we sum up
the third reaction with the first reaction, we will get the second reaction. In consequence,
when the concentration of NaOH and HCl solution is controlled to be equally large in all of
the three reactions, it can be stated that:
∆Hrx1 + ∆Hrx3 = ∆Hrx2
To determine the value of ΔHrx1, ΔHrx2, and ΔHrx3, in this experiment we will use a simple
calorimeter made of styrofoam cup. The styrofoam cup is closed using a punched-stryrofoam-
cap with a thermometer and stirrer rod into the styrofoam cup. Styrofoam is a good insulator,
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although some heat would be absorbed by the styrofoam and some would be released to the
surrounding, its value is quite small compared to the amount of heat absorbed by the solution
inside the calorimeter. Consequently, in this experiment, we assume that there is no heat
absorbed by styrofoam cup, styrofoam cap, thermometer, stirrer rod, and the surrounding of
the styrofoam cup. The other types of calorimeter can also be used to determine the amount of
the reaction heat involved in acid-base neutralization reaction.
(EXPLANATION OF THE CALCULATION WILL BE GIVEN BY THE ASSISTANT
DURING THE EXPERIMENT)
Chemicals and Equipment
The chemicals required in this experiment are listed below:
HCl 2M, NaOH 2M, solid NaOH, Demineralized water. Watch Out! Solid NaOH are
hygroscopic and can cause skin irritation
The equipment needed in this experiment are listed below:
Simple calorimeter, stopwatch, rod stirrer, analytical balance, test tubes, Erlenmeyer flask, 50
ml measuring cup, 50 mL/100 mL beaker glass, thermometer.
Procedures
PART A
Section A.1: Determination of the Heat of Neutralization Reaction: HCl(aq) + NaOH(aq)
a. Prepare a styrofoam cup that will be used as calorimeter.
b. Pour 25 mL of 2M HCl into the calorimeter and close the calorimeter using the
calorimeter cap that has been mounted with thermometer.
c. Take 25 mL of 2M NaOH solution and put the NaOH solution into a 50 mL beaker glass.
d. Measure the temperature of each of the solution.
e. Turn on the stopwatch. At t = 0 seconds, transfer the NaOH solution into the calorimeter
containing 25 ml of HCl 2M immediately, then close the calorimeter immediately
(thermometer has been mounted through the cap).
f. Stir the HCl and NaOH mixture, until it is well-mixed.
g. Measure the temperature of the solution at t = 10 seconds.
h. Stir the mixture and measure the temperature of the solution in the calorimeter every 10
seconds, until maximum temperature is obtained and the temperature is relatively
constant or decreases slowly and then relatively constant.
i. Calculate the moles of each reagent (HCl and NaOH) and the products.
j. Calculate the heat of the neutralization reaction per mole for the occured reaction.
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Section A.2: Determination of the Heat of Neutralization Reaction: HCl(aq) + NaOH(s)
a. Mix 30 mL of 2M HCl with 20 mL of demineralized water in calorimeter. Measure and
record the temperature of the solution.
b. Weigh 6.00 g of solid NaOH.
c. Turn on the stopwatch. At t = 0 seconds, immediately pour the solid NaOH into the
calorimeter using a spatula. Watch Out! Solid NaOH are hygroscopic and can cause
skin irritation.
d. Stir the mixture until it is well-mixed.
e. Measure the temperature of the solution at t = 10 seconds.
f. Stir the mixture and measure the temperature of the solution in the calorimeter every 10
seconds, until maximum temperature is obtained and the temperature is relatively
constant or decreases slowly and then relatively constant.
g. Calculate the moles of each reagent (HCl and NaOH) and identify which reagent that acts
as the limiting reagent.
h. Calculate the moles of the obtained product.
i. Calculate the heat of the neutralization reaction per mole for the occured reaction.
Section A.3: Determination of the Heat of Solution: NaOH(s) → NaOH(aq)
a. Prepare 50 mL of demineralized water inside a calorimeter. Measure and record the
temperature.
b. Weigh 6.00 g of solid NaOH.
c. Turn on the stopwatch. At t = 0 seconds, immediately transfer the solid NaOH into the
calorimeter.
d. Stir the mixture until it is well-mixed.
e. Measure the temperature of the solution at t = 10 seconds.
f. Stir the mixture and measure the temperature of the solution in the calorimeter every 10
seconds, until maximum temperature is obtained and the temperature is relatively
constant or decreases slowly and then relatively constant.
g. Calculate the heat of solution per mole of solid NaOH in water. Use the calculation in
section (A.1) and (A.2) to calculate the heat of solution (use Hess's Law).
PART B. Energy Changes in Chemical Reactions: exothermic and endothermic
reactions
a. Coat the bottom (outer part) of a test tube using silicon grease and stick a few grains of
solid iodine (I2) on it.
b. Put some solid CuSO4 into the test tube.
c. Place the test tube in an erlenmeyer flask.
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d. Add a few drops of water into the test tube until all of the CuSO4 become soaked.
Immediately close the test tube with a cork/rubber cap. Observe and record the
phenomenon!
DATA PROCESSING
Part A.
Follow these steps to process the acquired data for section A.1-A.3:
a. Plot the T (temperature, ° C) changes to t (time, seconds) for each reaction (A.1-A.3).
b. Indicate the initial temperature and the final temperature for each of the above reactions,
as shown in Figure 4.1.
c. Calculate the difference in temperature (T) for each of the above reaction.
d. Calculate the heat absorbed by the calorimeter, q1.
e. Calculate the heat absorbed by the solution, q2.
f. Calculate the heat produced in the reaction, q3 (q3 = q1 + q2).
g. Calculate the reaction enthalpy per mole, H (H = q3 / mol substances involved in the
reaction).
Please bring them together: • Lab Journal • Long sleeve lab coat • Graphical paper • Goggle • Calculator • Ruler • A rag/ tissue
15
Experiment IV
Chemical Bonding and Molecular Polarity
Introduction
Free atoms are not commonly found in nature because most of them are too reactive, thus
they tend to interact with each other to form chemical bond. How atoms form chemical bonds
depend on their electronic structure and types of bonds that can influence the chemical
properties of that compound. In binary ionic compound, two atoms that involved in the
formation of ionic bond are metallic and non-metallic ions. Those atoms are relatively
different but they complement each other; metal atoms tend to release electrons while the
non-metallic atoms are more likely to receive electrons. As the result, electron transfer occurs
from the metal to non-metal. Meanwhile, covalent bond is formed from the mutual usage of
electrons from both of atoms that produces attraction forces in between two nuclear atoms
with relatively strong interaction. The two types of bonds may look completely different, but
further differences and similarities may be identified by looking at continuous spectral data of
the two bonds. The relationship between both types of bonds can be understood through the
concept of electronegativity. The electronegativity is a parameter to describe the relative
attraction forces of an atom by dividing their electrons in a chemical bond. The more
electronegative an atom is, the bigger is its tendency to attract electrons. The difference in
electronegativity between 2 atoms in a chemical bond can be used to predict the bond’s
polarity. Polarity is a parameter to describe the difference of electron distribution in a
chemical bond. When two identical atoms (which would have the same electronegativity)
share one or more its electron pairs, each atom will have equal attraction force to the
electrons. This type of bond is called non-polar covalent bond. When two different atoms
(which has difference in electronegativity) are involved in covalent bond, the more
electronegative atom will attract electrons stronger than the other atom, and this will create an
imbalance in electron density distribution between the two atoms. This type of bond is known
as polar covalent bond.
Experimental data have shown that most chemical bonds are neither 100% covalent nor 100%
ionic, but instead have characters of both bonds at a certain proportion. To simplify, there are
rules to follow:
1. If there is no difference in electronegativity, the chemical bonds is classified as non-
polar covalent bond.
2. If the difference in electronegativity ranges between zero and 1.7, the chemical bond
is classified as polar covalent bond.
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3. If the difference in electronegativity is bigger than 1,7 it is classified as ionic bond.
For compounds that consist of 3 atoms or more, the polarity is determined by its molecular
geometry because it affects the total electron density in that molecule. Thus, the general
directions to predict molecular polarity are:
1. Molecules that consist of identical atoms are classified as non-polar molecules.
2. For molecules that consist of non-identical atoms:
a. If the 3-dimensional arrangement of the molecule is symmetrical, the molecules
are classified as non-polar.
b. If the 3-dimensional arrangement is asymmetrical, the molecules are classified as
polar.
This experiment aims to give a comprehensive knowledge for students on how to discriminate
types of chemical bonds, how to predict the polarity of molecules based on the polarity of its
chemical bonds and its geometry, and to also be capable to operate Avogadro software.
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Experimental Procedures
1. Arrange 7 molecule models of the first set of molecules on the worksheet.
a. Use the following colors to represent the atoms:
H = white; C = black; O and S = red; N = blue; F, Cl, Br, I = green;
b. Use a flexible connector (bond) to connect two atoms that form double bonds.
c. Evaluate what types of bonds are formed, determine the molecular geometry, and
write down whether the molecule is polar or non-polar.
2. Arrange 7 molecule models of the second set on the worksheet.
a. Evaluate what types of bonds are formed, determine the molecular geometry, and
write down whether the molecule is polar or non-polar.
b. If the molecule consists of three or more atoms, evaluate the type of bond on each
atom, determine the molecular geometry, and determine whether the whole
molecule is polar or non-polar.
3. When you have done the whole task, disassemble the model that you have created.
4. Download (http://avogadro.cc/wiki/Get_Avogadro) and install Avogadro software in
your group computer. Each group consists of 2-3 peoples. Every student must learn
the software before doing this experiment (mandatory!). Draw 14 molecules (7
molecules of the first set and 7 molecules of the second set), that you have been
choose, using the Avogadro software. Optimize the molecule geometry and analyze
the bond lengths and bond angles of each molecule. Discuss those molecules geometry
whether they are in accordance with the data in the textbook or not.
Please bring them together: • Lab Journal • Long sleeve lab coat • Laptop • Installed Avogadro software
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Experiment V
Physical and Chemical Properties Of Gases
Introduction
Gas is one of the states in which a matter can exist. Most properties of gases are independent
from its content. Several physical properties of gases include:
1. Compressibility: gases can be compressed and expanded according to the size and
shape of its container.
2. Gases are 1000 times less dense than solids and liquids.
3. Gas expands when heated.
4. Gas dissolves in other gas in various proportions.
The properties of gases are measured by four main parameters: pressure (P), volume (V), the
amount of gas particles in mole (n) and temperature (T). During the 18th century, these
properties are formulated into ideal gas laws. These laws are:
1. Boyle’s Law – the volume of a gas is inversely proportional to its pressure,
mathematically written as:
𝑃 ~ 1
𝑉 ; 𝑃1𝑉1 = 𝑃2𝑉2 (𝑎𝑡 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 𝑛 𝑎𝑛𝑑 𝑇)
2. Charles’ Law – the volume of gas is directly proportional to its temperature,
mathematically written as:
𝑉 ~ 𝑇 ; 𝑉1
𝑇1=
𝑉2
𝑇2 (𝑎𝑡 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 𝑛 𝑎𝑛𝑑 𝑃)
When V is plotted against T, a linear graph will be produced. When the line is
extrapolated until the volume reaches 0, it will intersect the axis at T= -273,15 oC,
setting the value as absolute zero temperature. Based on this calculation, the scale for
absolute temperature is defined as Kelvin temperature scale, with 0 K = -273,15 oC
and 273K = 0 oC. For all calculations using gas laws, the temperature must be stated in
kelvins.
3. Avogadro’s Law – equal volumes of gases, at the same temperature and pressure,
have the same number of molecules. It can be written as the equation:
𝑉1
𝑛1=
𝑉2
𝑛2
4. Gay-Lussac’s Law – the pressure of a gas is directly proportional to its temperature,
mathematically written as:
𝑃 ~ 𝑇 ; 𝑃1
𝑇1=
𝑃2
𝑇2 (𝑎𝑡 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 𝑛 𝑎𝑛𝑑 𝑉)
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The combination of all the gas laws produces the Ideal Gas Law, which is formulated as the
following equation:
𝑃𝑉 = 𝑛𝑅𝑇
with R being the universal gas constant; R=8.314 JK-1mol-1 or R=0.08206 L.atm.mol-1.K-1
The Ideal Gas Law is commonly used to calculate properties of gases. Most real gases behave
like ideal gas in low pressure and concentration. It is also necessary to define molar gas
volume, 𝑉𝑚 , which is the volume of 1 mole of gas at a certain temperature. It may be
calculated using the Ideal Gas Law at various pressure and temperature.
Chemicals and Equipment
Chemicals needed for this experiment is listed below:
Solid Zn, Mg, Al, HCl 6M, methylene blue, ice, CH3COOH, NH3, universal pH indicator and
demineralized water.
Equipment needed for this experiment is listed below:
Standard glassware, volume determination equipment set, microscale chemistry set, plastic
glass, etc.
Procedures
A. Determination of An Unknown Metal’s Relative Atomic Mass and Molar Volume of
Hydrogen
1. Weigh approximately 0,2 gram of a powdered unknown metal (write down the exact
measurement), then put the powder into flask 3 (the reactor).
2. Fill flask 2 with tap water up to the bottom part of the flask’s neck.
3. Empty both tubing mounted on flask-2’s cork by using a hand-pump (tubing B and
one half of tubing A), then assemble the set according to the diagram. Make sure the
corks are tightly fitted.
4. Measure 25 mL of 6 M HCl solution, pour it through the funnel on flask 3. The funnel
is bounded to a squeeze-tap, therefore in order to let the acid solution flow down, the
tap must be pressed. Let all the solution drop down.
5. Rinse the measuring cylinder and the funnel using a known amount of demineralized
water, let it flow into the reactor. The exact volume of water used must be known.
6. During the experiment and reaction takes place, water will flow from flask 2 to glass
1. Wait until the reaction between HCl and the metal completely finished and no more
water flows into glass 1.
7. Measure the temperature of water in glass 1, then measure its volume using the
cylinders.
8. Measure room temperature and pressure using a digital barometer and thermometer
(provided by the lab staff’s bench)
9. Calculate the relative atomic mass of the metal used based on acquired data.
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B. Temperature Impact on Gas Volume
1. Prepare plastic pipette and microplate.
2. Fill one of the well on the microplate with water (choose the larger well). Add 2-3
drops of methylene blue.
3. Place the bulb of the plastic pipette into a plastic measuring cup (as indicated at the
figure below).
4. Insert the open end of the plastic pipette to the water-filled well.
5. Hold the open end under the water surface.
6. Fill the plastic measuring glass with ice cube, then let the bulb touch the ice cube.
Observe the water level changes inside the tube.
7. Remove all ice cubes from the cup. Let the air temperature inside the bulb reach room
temperature.
8. Fill the plastic measuring cup with some hot water.
9. Dip the bulb into the hot water. Observe the water level changes inside the tube.
C. Gas Diffusion
1. Choose one line of small wells in the microplate in your microscale chemistry set.
2. Add one drop of universal indicator into every well in the line except the first and the
last wells.
3. Put 10 drops of water into each well containing universal indicator, or until the well is
completely full.
4. Put 10 drops of vinegar into one of the empty well within the row (either the first or
the last hole).
5. Put 10 drops of ammonia into the other empty well within the row.
6. Immediately cover the microplate using the plastic tray provided.
7. After 5 minutes, lift the tray and observe the color changes of the universal indicator
solution .
Please bring them together: • Lab Journal • Long sleeve lab coat • Goggle • Calculator • A rag/ tissue
