Chemistry as a science that has several branches, one of them is thermochemistry. Thermochemistry is concerned with the study of heat changes during a chemical process, for example, in some experiment heat is absorbed or emitted during a chemical reaction. The formula for change in enthalpy and Hess law can be used to determine heat change in a chemical reaction. In this laboratory work, the heat change of different reactions was studied using Hess law. The law predicts that the total enthalpy change during a whole course of a chemical reaction is similar whether the reaction occurs subsequently or in a single step.
Mathematically change in enthalpy is expressed as;
………………………………………………………………………………………..(1)
Where q is the heat change during the reaction and can be expressed as, m is mass of the substance in question and is temperature change, given as final temperature (Tf) minus initial temperature (Fi), and c the specific heat of the substance.
Objective
The primary purpose of this lab work is to employ Hess law in finding the heat of neutralization of a chosen reaction by using the observed heat of neutralization from three different reactions. The found value will then be compared with a theoretical value to ascertain its accuracy.
Materials
Two stacked styrofoam cups, cover, thermometer, ring stand, burette clamp, sodium hydroxide hydrochloric acid, acetic acid, and ammonia.
Procedure
The two stacked Styrofoam cups and the cover were used as a calorimeter to determine the temperature change during the various reactions. The thermometer supported by a ring stand and a burette clamp was used to determine the temperature during and before the reaction. The reactions that were considered are the reaction between sodium hydroxide and hydrochloric acid, sodium hydroxide and acetic acid and ammonia, and hydrochloric acid.
First, 25 ml each of 1M sodium hydroxide and 1M hydrochloric acid were measured and the initial temperature of each measured. The solutions were mixed together and transferred to the calorimeter, and a thermometer was inserted through the cover into the calorimeter. Second, 25 mL each of 1M sodium hydroxide and 1M hydrochloric acid were measured and their temperatures noted. The solutions were again mixed and transferred into the Styrofoam cups, and the temperature of the mixture was measured. For both reactions, the temperature change was observed and recorded in Table 1, after every 30 seconds. The final temperature of the reaction was determined after the temperature of each reaction stabilized. Equation 1 was used to calculate the heat change during the reactions; the found value was then multiplied by 40 and divided by 1000 to find the heat change per mole and in kilojoules. The procedures were repeated for the reactions between 1M sodium hydroxide with 1M acetic acid and 1M ammonia with 1M hydrochloric acid. ΔH for the reaction between ammonia and acetic acid was determined using Hess law.
Data and Calculations
Table 1
Base
Acid
TI
+30 sec
+30 sec
+30 sec
Tf
ΔT
ΔH
NaOH (1M) 25mL
HCl (1M) 25mL
19° C
24.5° C
25° C
25° C
25° C
6° C
-1255.2 J
NaOH (1M) 25mL
CH3COOH (1M) 25mL
19° C
23° C
24.5° C
24.5° C
24.5° C
5.5° C
-1150.6 J
NH3
(1M) 25mL
HCl (1M) 25mL
19° C
21.5° C
23.5° C
24° C
24° C
5° C
-1046 J
NaOH(aq)
+ HCl(aq) → NaCl(aq)
+ H2O(aq)
NaOH(aq)
+ CH3COOH(aq) → NaCH3COO(aq)
+ H2O(aq)
NH3(aq)
+ HCl(aq) → NH4Cl(aq)
NH3(aq)
+ CH3COOH(aq) → CH3COONH4(aq) ΔHf° = ???
Applying Hess Law
NaCl(aq)
+ H2O ← NaOH(aq)
+HCl(aq) ΔHf° = +50.208 kJ
NaOH(aq)
+ CH3COOH(aq) → NaCH3COO(aq)
+ H2O(aq) ΔHf° = -46.024 kJ
NH3(aq)
+ HCl(aq) → NH4Cl(aq) ΔHf° = -41.840 kJ
NH3(aq)
+
CH3COOH(aq) →
CH3COONH4(aq) ΔHf° = -37.656 kJ
Theoretical Values
Literature Values
NaOH(aq)
+ HCl(aq) → NaCl(aq)
+ H2O ΔHf° = -57.1 kJ
NaOH(aq)
+ CH3COOH(aq) → NaCH3COO(aq)
+ H2O ΔHf° = -55.2 kJ
NH3(aq)
+ HCl(aq) → NH4Cl(aq) ΔHf° = -52.2 kJ
NH3(aq)
+ CH3COOH(aq) → CH3COONH4(aq) ΔHf° = ???
Hess’ Law Equation
NaCl(aq)
+ H2O ← NaOH(aq)
+HCl(aq) ΔHf° = +57.1 kJ
NaOH(aq)
+ CH3COOH(aq) → NaCH3COO(aq)
+ H2O ΔHf° = -55.2 kJ
NH3(aq)
+ HCl(aq) → NH4Cl(aq) ΔHf° = -52.2 kJ
NH3(aq)
+
CH3COOH(aq) →
CH3COONH4(aq) ΔHf° = -50.3 kJ
Percentage Error for the Experimental Values
NaOH(aq)
+ HCl(aq) → NaCl(aq)
+ H2O
NaOH(aq)
+ CH3COOH(aq) → NaCH3COO(aq)
+ H2O
NH3(aq)
+ HCl(aq) → NH4Cl(aq)
NH3(aq)
+ CH3COOH(aq) → CH3COONH4(aq)
Discussion
All the experimental enthalpy values were negative, showing that all the reactions were exothermic. Exothermic reactions release heat into the environment and that is why the values of enthalpies in the reactions were negative. The rise in temperature measured is an indication that heat was being released during the reaction.
The experimental values did not match the theoretical values; the experiment values had a large percentage error as compared to the theoretical values. The large percentage error could have resulted from errors during the experiment. One possible source of error is human errors; the errors could have occurred during measurement, such as parallax error. Also, during the mixing of the solutions, some little solution could have remained in the original beakers where they were prepared. Another source of error is the chemical reaction, maybe all the solutions did not react completely or the reaction was to slow for any change to be noticed. The calorimeter used in the experiment was not a perfect one; it could have affected the temperature measurement.
The Hess’ law was however confirmed, the energy released in the first reaction was greater than the second and the second was greater than the third, this result was similar to the theory.
Thermochemistry is important since it enables people to know whether energy is released or absorbed in a reaction. For example, in this lab work, the energy released can be converted to usable forms and be used to drive machine engines or to do other work. Therefore, by studying the heat change of a reaction, an individual can determine the best use of a reaction.
Conclusion
In this lab, Hess law was successfully applied to find the heat of neutralization of ammonia and acetic acid. The enthalpy change was a little far from the theoretical value, but the experiment objective was achieved.
Titration of Creek Water
Introduction
Titration is known as the slow addition of one solution of a known concentration, a titrant to a solution of a predetermined volume, but unknown concentration, analyte until the reaction reaches neutralization, which is often indicated by a color change.
Objective
Titration can be used to determine the concentration of elements in a solution, in this lab the concentration of chlorine in Newton creek water was to be determined through titration.
Materials
Creek water, Silver Nitrate (AgNO3)
Methodology
Creek water was titrated with a 2M solution of Silver Nitrate (AgNO3). At initial stages, a precipitate was observed, which an indication of the formation of AgCl. After some time, the precipitate disappeared. The disappearance was due to neutralization of all the chloride ions in the water. The silver displaced potassium in chromate (K2CrO4) to form silver chromate (Ag2CrO4). The chemical reaction equations, before and after the neutralization process is as shown below.
Before neutralization: Ag+(aq)
+ Cl-(aq)→ AgCl(s)
After Neutralization: 2Ag+(aq) + CrO42-(aq) → Ag2CrO4(s)
The titration was done in three trials, using 10.0 mL of creek water. On the third trial, 15.10 mL of the .2M AgNO3 solution was used to titrate the sample. Since the third trial gave better results, it was used without averaging other trials. The 15.10 mL of .2M AgNO3 was converted to .00302 moles of AgNO3 as follows.
The ratio of bonding of Ag+ to Cl- in 1:1, from this ratio it was easy to conclude that 0.00302 M of Cl- took part in the reaction process, and that was the concentration of the Cl- in the creek water used. The concentration was converted to molarity to determine the concentration of Cl- per liter of creek water, 0.302 M.
The concentration was then converted to mg/L; this was done by multiplying the determined concentration by the molar mass of Chlorine which is 35.453 AMU and converting the resulting answer to milligrams. Since the test sample was 10.0 mL, the final answer was multiplied by 100.
The milligrams per liter was finally converted to parts per million (ppm) as follows
Results
The molar concentration of Cl- in the water 0.302 M
Mass per liter of Cl-
I in the water, 10, 706.8 mg/L
Discussion and Conclusion
From this lab, it was discovered that Cl- ions are present in Newton creek water, but the level was high as compared to Environmental Protection Agency standards of 250 mg/L. There could have been an error in the experiment and the possible sources are human errors and experimental errors. The experiment involved color change and it is not easy to precisely determine when the color started changing. Another human error could have occurred during measurement of the solutions used during the experiment. The creek water could have also contributed to other errors as there was no surety about the existing elements in the water, maybe the other existing elements reacted with the titrant to give false results. The sources of chlorine in creek water include road salts, industrial processes, fertilizers, and sewage.
Determining the Empirical Formula of a Hydrate
Objective
The objective of this lab was to determine the empirical formula of unknown hydrate, by evaporating its water content and to understand what empirical formula stands for in the chemical formula of a compound.
Materials
Equipment: Crucible, crucible cover, crucible tongs, bunsen burner, clay triangle, wire screen, bunsen burner, ring stand, clamp, scale.
Chemicals: Ionic hydrate sample,
Methodology
The crucible and the lid were washed with soap and water and then dried with a paper towel. The cleaned crucible was placed opened on a clay triangle using crucible tongs and supported by an iron ring attached to a ring stand. A Bunsen burner was lit and its flame adjusted to a blue color with a lighter blue in the inner cone. The burner was then placed under the crucible to dry it. The heating was initially done slowly for 5 minutes; the heat was then increased for 3 minutes until the bottom of the crucible glowed red. The crucible was finally removed from the triangle using the tongs and placed covered on a wire screen on a lab bench to cool to room temperature. The cooled crucible without the cover was weighed, and the result recorded in Table 1. The steps were repeated until the crucible was heated to a constant mass.
Ionic hydrate was obtained from the lab and its identification code noted. At least 1 g of the hydrate was weighed directly into the crucible, and the mass of the sample and the crucible were measured and recorded in Table 2. The crucible plus the sample was placed on the clay triangle using the tongs, the lid was placed on top of the crucible (slightly ajar), and the crucible and its contents heated slowly for 5 minutes, the flame was increased to medium for 10 minutes. The hot crucible was transferred to the wire screen covered and allowed to cool to room temperature. The crucible was weighed without the cover and the result recorded in Table 2. The crucible plus the contents as further heated in a medium flame for 10 minutes with the cover opened, this was repeated until a constant mass of the crucible and its content was achieved. The results obtained were recorded in Table 3 and used to determine the formula of the anhydrous compound.
Results
Table 1
Mass of empty crucible
13.356
Mass after first heating and cooling
13.224
Mass after second heating and cooling
13.172
Table 2
Mass of the Crucible + Sample
14.189
Mass after the first heating and cooling
14.083
Mass after the second heating and cooling
14.089
Table 3
Mass of original ionic hydrate sample
1.0213g
Mass of water given off by sample
0.1249g
Mass percent of water in the ionic hydrate*
12.23%
Mass of anhydrous compound
0.901g
The chemical formula of the anhydrous compound
BaCl2
Molar mass of the anhydrous compound
Ba(137.33)+2Cl(35.45)= 208.236 AMU
Number of moles of anhydrous compound*
0.0043 moles
Molar mass of water
18.013 AMU
A number of moles given water given off by the sample.
0.0069 moles
The ratio of moles of water divided by moles of anhydrous compound*
1.6125:1 which is approximately (2:1 ratio)
The chemical formula of ionic hydrate*
BaCl2. 2H2O
The hydrate used was barium chloride salt. As the hydrate was heated it gave out water, and the reaction can be written as;
Amount of water in 3 g of the ionic hydrate is;
Discussion
A compound is a chemical substance that contains more than one element.
Hydrate is a salt bonded with water (H2O) and has a formula salt·x H2O where x is the number of water molecules attached to each salt molecule.
The water of hydration is water that is chemically combined with a substance to form a hydrate; the water can be removed from the substance though heating without changing the composition of the substance.
The formula of the compound used was BaCl2. 2H2O. The experimental formula is similar since I also achieved a ratio of 1:2 (ratio of the hydrous salt to water in the hydrated salt). My ratio was not exact, but it was rounded to the nearest whole number. The lack of accuracy could have been caused by errors during heating, may be part of the hydrous salt was lost in the air. Also, measurement errors could have occurred.
Since the crucible was washed the heating was important to help in getting rid of all the water and that is why it had to be heated to a constant mass. The water in the crucible could have altered the values of the experiment if it was not fully evaporated.
Conclusion
The objective of the experiment was met, the empirical formula of the unknown salt was determined.
Classification of Chemical Reactions
Introduction
The evidence observed in this lab included effervescence, color changes, heat absorption, production of sound, the formation of new odors, and the formation of a precipitate.
Objectives
We performed and observed the results of a variety of chemical reactions in this lab. The lab also entailed familiarizing ourselves with different observable signs of a chemical reaction. Also, we were to identify the products of the chemical reactions and to write a chemically balanced equation.
Materials
Clean test tubes, Bunsen burner, beakers, stirring rods, soap, and water (for washing the glassware) distilled water, 10mL graduated cylinder and tongs.
Chemicals: 3M HCL, solid Na2CO3, 0.1M CaCl2, 0.1M Na3PO4, solid Cu(OH)2, 1M CuSO4, Zn metal, Mg metal, 1M ZnSO4, Cu wire, 6M HCL, 3M NaOH, 0.1M FeCl3, 0.1M KSCN, 3M HCL, 3% H2O2, solid MnO2, CuSO4, 1M NaCl, 1M KNO3.
Methods
1st
experiment
5 mL of HCl was added into a test tube standing upright in a beaker. A small piece of zinc metal was dropped into the acid in the test tube and observation made. The mouth of the test tube was closed using a thumb for 30 seconds. A burning wooden splint was carefully brought next to the thumb, closing the test tube. The thumb was removed and the burning splint placed at the mouth of the beaker and observation recorded under results.
2nd experiment
Using crucible tongs, a copper wire was held in the hottest part of a burner flame for 2 minutes. The wire was examined and it was noted that the color changed from brown to black.
3rd experiment
5 ml of HCl was mixed with 5 ml of o.5 M sodium carbonate in a test tube; a burning wooden splint was brought at the mouth of the test tube after 2 seconds. The wooden splint went off after 5 seconds.
4th experiment
Using the crucible tongs, a small piece of magnesium metal was placed in the flame of a Bunsen burner, after which it was removed and allowed to continue burning. After, it cooled to room temperature it was noted that the metal had changed to a grey powder.
5th experiment
25 mL of hydrogen peroxide was poured in a beaker; a catalyst was added to speed up the reaction. When a burning splint was brought to the mouth of the beaker, the flame glowed brightly, a sign that the reaction in the beaker was producing oxygen.
Data and Calculations
1st
experiment
When the burning splint was brought to the mouth of the test tube, a pop sound was heard an indication of hydrogen gas. Theoretically, metals react with acids to produce hydrogen gas and a respective metal salt. This is an example of a single replacement reaction where an element and a compound come together, but the element takes the compound’s ion.
Zn(s) + HCl(aq)
→ ZnCl2(aq) + H2(g)
2nd
experiment
The second experiment is an example of a combination reaction, where two elements come together to form a single compound. The black solid observed is a copper oxide, formed as a result of burning copper in the air. The reaction is an example of a combination or synthesis reaction.
2Cu(s) + O2(g)
→ 2CuO(s)
3rd
experiment
The wooden splint went off as a result of the presence of CO2 gas. From theory, metal carbonates react with acids to produce a new metal salt, carbon (iv) oxide gas and water. The reaction was an example of a double displacement reaction.
2HCI(aq) + Na2CO3(aq)
→ 2NaCl(aq) + H2O(aq)
+ CO2(g)
4th
experiment
The new material formed after burning of magnesium metal is known as magnesium oxide, and this was an example of a combustion reaction. Combustion reacts occurs if an element burns completely in the air to produce a new compound.
2Mg(s) + O2(g)
→ 2MgO(s)
5th
Experiment
The oxygen produced was as a result of decomposition of hydrogen peroxide and this type of reaction is known as decomposition reaction.
2H2O2(aq)
→ 2H2O(l) + O2(g)
Discussion
Various chemical compounds and elements were used in this lab to observe the five different chemical reactions. The observed reactions were single displacement, double displacement, combination or synthesis, combustion, and decomposition. Different chemicals were used to observe the reactions, but only chosen reactions were recorded in this report.
Most of the reactions in this lab could not give physical observable results, but from a different test like a burning splint, we were able to identify some products of the reactions. Other tests like the use of litmus paper were also useful, in knowing the progress of a reaction during the experiment.
Errors like improper chemical ratios could have occurred in some cases since we were not able to witness changes in some reactions. Another error could have resulted from chemical contamination, we had several chemicals and chances are high some were mixed accidentally.
Conclusions
The experiment was a success since we were able to identify all the five chemical reactions. The chemical formulas were also written and balanced properly.
This lab was good but a lot of precautions need to be put in place. For example, hydrogen peroxide is dangerous and should be handled carefully. The acids also need proper handling.
Quantitative Analysis of Ions
Objective
A number of methods were used to test for the presence of ions; the purpose was to determine the ion in an unknown solution. Flame tests were used to test the presence of cations, such as; sodium, potassium, calcium, aluminum, and copper. For anions, we did a range of tests where we reacted the ions with compounds, looking for precipitate or by using a color indicator. These indicated the presence of carbonate, sulfite, phosphate, chloride, and thiocyanate.
Materials
Clean glassware, Distilled water (to clean glassware between experiments), Hot water bath (200mL distilled water in the 250mL beaker), Hotplate, Bunsen burner, Flame test wire, 6M-HCl (for cleaning), Cobalt blue glass, Litmus red.
Methods
Sodium, Na+
a. A flame test wire was dipped in HCl and placed in a flame until the flame became colorless.
b. The text wire was then dipped in 0.2M-NaCl and placed in a flame to test for Na+. A bright yellow flame is an indication of sodium ion, and that what was expected.
Potassium, K+
Step 1(a) was repeated.
b. The text wire was then dipped in 0.2M-KCl and placed in a flame to test for Na+. A fleeting lavender color flame is an indication of potassium ion, and that what was expected.
Calcium Ca
10 drops of 0.1 M Ca(NO)2 was combined with 3 drops of 1 M NaCO3 in a test tube, a white precipitate of calcium oxalate was expected as an indication of the presence of calcium.
In another tube 10 drops of 0.1 M Ca(NO)2 was mixed with 3 drops of 6 M HCl and a flame test was performed. Fleeting orange was to be a sign of calcium ion presence.
Aluminium, Al
NH (may have a miasmic odor) was combined dropwise to 10 drops of 0.5M AlCl3 until the solution became basic, (testable with litmus red).
White gelatinous Al(OH)3 formed, to this 3M of HCHO was added drop by drop and mixed until it dissolved.
2 drops of catechol violet reagent were added, which indicated the presence of aluminium ion.
Copper (II), Cu
Drops of concentrated NH was added to 10 drops of 0.5M CuSO4 in a test tube. Dark blue color confirmed the presence of copper (II) ion.
Carbonate, CO3
10 drops of 1M NaCl was mixed with 10 drops of 6M HCl. CO2 bubbles and color/odorless gas indicated carbonate ion.
Sulfate, SO4
10 drops of 1M HCl was carefully added to a test tube with 10 drops of 0.5M Na2SO4
and mixed to acidify solution.
4 drops of 1M BaCL2
was added. A white powdery precipitate was observed.
Phosphate, PO
10 drops of 0.1M NaPO was mixed with 10 drops of 6M HNO.
10 drops of 0.5M (NH ) was added to MoO and mixed.
The test tube was heated in a warm water bath for 10 minutes.
Yellow precipitate (not a solution) indicated phosphate ion
Chloride, Cl-
10 drops of 0.2M NaCl was mixed with five drops of 1M HNO.
3 drops of 0.1M AgNO was added to the solution from step 9a, a white precipitate of AgCl indicated the presence of chloride.
Precaution: Should SCN be present, it will also form the white precipitate invalidating the previous test if your unknown contains this compound follow this procedure.
Mix 10 drops of the unknown with 10 drops of HNO.
Boil the solution in a water bath until only half the volume remains.
Add 2-3 drops of 0.1M AgNO and look for the formation of white AgCl precipitate.
Thiocyanate, SCN
Mix 10 drops of 0.5M KSCN with 10 drops of 3M HCHO(CHCOOH).
Add 0.1M Fe(NO) dropwise until the solution becomes blood-red in color indicating for the formation of [Fe(HO)SCN] +, meaning her is a presence of thiocyanate.
Unknown
For unknown, replace the substance being tested for and continue experiment as instructed above.
Results
ION TESTED
Known
Unknown
Na+
Bright yellow flame
Bright yellow flame
K+
Fleeting lavender flame
No observation
Ca2+
A white precipitate, fleeting orange flame.
White precipitate
Al3+
Red on litmus paper, white gelatinous Al(OH) forms, catechol reagent turns violet.
No observation
Cu2+
Dark blue color
No observation
CO32-
Bubbles and color/odorless gas
No observation
SO42-
White powdery precipitate
White precipitate
PO43-
Yellow precipitate
No observation
Cl-
White precipitate
No observation
SCN
Blood red color
No observation
Discussion
Qualitative analysis is the process by which components of mixtures are singled out through chemical process identification. The primary purpose of the qualitative analysis is confirming the existence or absence of certain materials from an unknown compound.
When the flame test was conducted on the unknown solution, a bright yellow flame was observed. From theory a bright yellow flame is an indication of sodium ion, hence it can be concluded that there were Na+
ions in the solution. When 10 drops of the unknown solution were combined with 3 drops of 1 M NaCO3 solution, a white precipitate was observed, the white precipitate is a sign of formation of calcium oxalate, which an indication of the presence of Ca+. The SO42-
ions test done on the sample was positive. Therefore, it can be concluded that the unknown solution contained Ca+, Na+, and SO42- ions.
The results obtained in this lab are important to chemists and the general public; it shows that chemical reactions can be used to test the presence and absence of various materials in the daily compounds we encounter. For example, such a test can be used to analyze the components in a water body, especially water bodies near agricultural lands, and to ascertain their safety.
The results achieved in this study are in accordance with theoretical expectations. The consistency of the results is an indication that the experiment was successfully done and the objectives achieved. Also, the consistency shows that qualitative analysis can be confidently used to determine the presence of different elements in a compound.
Conclusion
Based on this data, I believe that sodium, calcium, and sulfate are present in the unknown solution. The experiment should be done in a fume chamber because some of the gases released during the test are poisonous. Also, since the tests involve the use of acids and other corrosive chemicals, students should put on protective clothing such as gloves to prevent unnecessary accidents. Errors due to human errors such as human and experimental could have occurred during the experiment, the errors could have prevented us from identifying more elements in the unknown solution. Therefore, students should be given more accurate measurement methods other than the drops we used. The drops may have not been accurate as expected.
