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Homework Archive 20193-20211

20193.117436.Midterm1.NVA

20193.117436.Midterm1.NVA-01

Introduction

Titration is a key technique which is used as a method of quantitative analysis to identify the unknown concentration of a component (the analyte) in a solution by reacting it with a solution of another compound (the titrant) in order to find the equivalence point of this titration. [1]

Although there are many various types of titrations, acid-base titrations are used the most commonly. In this lab report, a sample of hydrochloric acid (HCl) is titrated with sodium hydroxide (NaOH) by the phenolphthalein indictor.

HCl + NaOH => NaCl + H2O

The purpose of this experiment is to determine the unknown concentration of a hydrochloric acid solution through the acid‐base titration.

Experimental procedure

1. Apparatus

Table 1: List of number of used laboratory instruments

Rubber bulb, wash bottle, ring stand

2. Chemicals

Used in the experiment are Phenolphthalein indicator, Sodium hydroxide solution (NaOH 0,1N), Hydrochloric acid solution (HCl). All the chemicals were available in the laboratory. In addition, all solutions were prepared using distillated water.

3. Produce

The steps below would be used to perform the titration described on introduce:

First, the lab wares were washed by distillated water

Second, we clamped the burette in a burette clamp and poured the titrant (NaOH 0,1N) into it carefully. Then we dropped some titrant through the stopcock into a waste beaker in order to made sure that no air bubbles remain in the stopcock.

Third,10ml of the unknown solution (HCl) was pipetted into an Erlenmeyer flask precisely and added 7-8 drops phenolphthalein indicators. At this point, the solution was colorless.

After that, we titrated Hydrochloric acid with sodium hydroxide until the endpoint of the titration was signaled when a permanent pink color change was observed (longer than 30 seconds). Next, we recorded the first volume of sodium hydroxide in the burette.

Finally, we performed three titrations at least and took average the trials (error not exceeding + -0.2%).

III) Results

Phenomena: The entire solution changed from colorless to pink after titrating. The product of the titration is shown on Fig.1

The results obtained are summarized in Table 2

Table 2: Volume of sodium hydroxide (NaOH)

The volume of the base solution after three experiments is approximately the same

Fig.1 The entire solution of the titration changed to pink color

IV) Discussion

Figure 1 shows the solution changed from colorless to pink after titrating. Phenolphthalein is an acid-base indicator which can be used to identify when the change from acidic to basic environment occurs [2]. It is colorless in acidic solutions and pink color in basic solutions. After the equivalence point, there was an excess of some sodium hydroxides and the solution had a basic environment. Therefore, the endpoint of the titration changed from colorless to pink around the equivalence point of this titration.

According to Table 2, the volume of the base solution after three experiments is approximately the same. During the experiment can be wrong. This may lead to incorrect results compared to the answer. The results obtained differ depending on the preparation of the experiment equipment and the titration operation of each person.

Then I could calculate the equivalent concentration of the acid solution (HCl).

HCl + NaOH => NaCl + H2O

Call N1, N2 and V1, V2 are equivalent concentration and volume of the solutions in the equivalence point of this titration. Applying this formula: N1.V1 = N2.V2, we found the concentration of HCl. From the table above we have:

N HCl. V HCl=N NaOH.V NaOH

Change the number into us: N HCl =0,0953 N

C(g/l) = N HCl. Đ HCl= 3,475 (g/l)

V) Conclusion

Thus, titration is a key technique of analytical chemistry which is used to identify the unknown concentration of a component in a solution. In this experiment, we titrated hydrochloric acid with sodium hydroxide through the phenolphthalein indictor. We produced three titrations at least and took average the trials. Then we calculated the equivalent concentration of the acid solution (HCl) According to this, we determined the unknown concentration of the acid solution (HCl)

20193.117436.Midterm1.NVA-02

I. Introduction

The purpose of this report is to synthesize the dye Para Red from p-nitroaniline that basing on diazotization reaction.

Para Red (paranitraniline red, Pigment Red 1, C.I. 12070) is a chemical dye which has characteristic similar to Sudan I, and it was explored as the first azo dye by two scientists Von Gallois and Ullrich in 1880. Para Red stains cellulose fabrics a bright red, so it has been applied commonly as colors or dye [2]. Diazotization is a method that primary aromatic amine is combined with the sodium nitrite in medium alkaline and acidic environments at low temperature to create a diazonium salt. This method is applied popularly for determining dyes [1].

Para Red is generated by diazotizing p-nitroaniline in a range of 5^oC to 10^oC, followed by a diazonium coupling reaction of diazonium salts and β-naphtholnate salt in a fusion of weakly acidic and weakly alkaline environments [3]:

Figure 1: The reactions of Para Red preparation [3]

The reason we must execute this experiment in ice-cold temperatures and a flexible environment was the decomposition of intermediate compounds at high temperatures, strongly alkaline or strongly acidic environments.

II. Methodology

1. Chemicals

P-nitroaniline (NO₂C₆H₄NH₂), sodium nitrite (NaNO₂), β-naphthol (C₁₀H₈O), a 1M sulfuric acid solution (1 M H₂SO₄), a 2.5 M sodium hydroxide solution (2.5 M NaOH), ice distilled water, and salt. All chemicals were available in the laboratory.

2. Synthesis of Para Red

Para Red was synthesized from p-nitroaniline via a diazotization reaction.

Firstly, we diluted 0.5 mL of hydrochloric acid with 10 mL of water in a 50 mL Erlenmeyer flask. Then, we stirred the solution while adding 1 g of p-nitroaniline. Therefore, this solution was decreased temperature to 0 ^oC by a freezing mixture which was prepared from some crushing ice mixed with a small amount of salt. Secondly, after 0.5 g of sodium nitrite was inserted into another 50 mL Erlenmeyer flask, 2 mL of water was poured into it, and the temperature of the solution was decreased to below 10 ^oC by the freezing mixture simultaneously. Thirdly, a medicine dropper was used to suck the solution in the second Erlenmeyer flask, and then, we dripped gradually it to the primary Erlenmeyer flask. Simultaneously, the mixture was made cool below 5^oC. At this time, we got diazonium salt. Finally, 10 mL of a 2.5 M sodium hydroxide solution was added to a 100 mL beaker containing available 0.52 g of β-naphthol. Next, we cooled the solution to 0 ^oC by the freezing mixture. After that, we stirred and continued to cool the solution in the beaker under 10 ^oC while we poured it gradually into the diazonium salt solution. At the end of the procedure, the solution was acidized by a 1M sulfuric acid solution

III. Results

Our results demonstrated that the Para Red product was brilliant red and indicated in Figure 2. The results presented in this report match the factual theory. [2]

Figure 2: The color of Para Red

IV. Discussion

The temperature was one of the most critical factors that had a direct impact on the success of the preparation for Para Red. The temperature of diazotization reaction of p-nitroaniline should be maintained within the range of 5^oC to 10^oC by the freezing mixture because diazonium salt would be decomposed in high temperatures that followed by a reaction [4]:

Figure 3: The decomposition of diazonium salt [4]

Additionally, all the reactions were exothermic reactions which were more likely to local overheating. Thus, we maintained the temperature of the solution under 5^oC.

It was noticeable that an available HNO₂ was not combined with p-nitroaniline to create diazonium salt in the diazotization reaction because the HNO₂ is not a durable compound and easy to be decomposed under the influence of light. Instead, we used a fusion of H₂SO₄ and NaNO₂ to create a newborn HNO₂, then coupling with p-nitroaniline to prepare the diazonium salt. This leads to optimize react performance.

Besides, the environment of the diazotization reaction was also a crucial segment that contributed to the success of this experiment. We implemented the procedures in a fusion of weakly acidic environments and weakly alkaline environments. The reason for this action was in a strongly acidic environment, β-naphtholnate salt will be demolished and transformed into β-naphthol which is difficult to dissolve and not be able to participate in a diazotization. On the other hand, the diazonium salt will be decomposed in a strongly basic environment [4]

V. Conclusion

We synthesized the dye Para Red from p-nitroaniline successfully by diazotizing p-nitroaniline in a range of 5^oC to 10^oC, and then, the resulting solution was reacted with β-naphtholnate salt in weakly acidic and weakly basic environments at a temperature below 10^oC. The received result was the brilliant red dye Para Red. [2]

20193.117436.Midterm1.NVA-03

Introduction

The aim of this experiment was refining acetanilide by recrystallization method.

I.1. Properties and applications of acetanilide

Acetanilide was an odourless solid [1] that doesn’t have color. In pure state, Acetanilide did not have savour. It melted at 114-115°C and boiled at 303.8°C [2]. Acetanilide dissolved well in etylic, ete, clorofrom, axeton and aniline. However, it dissolved badly in xylin and benzen. In addition, it didn’t dissolve in acid and lye. Acetanilide participated in nitrification reaction, halogenated reaction, hydrolysis reaction and reactions with HNO2.

Acetanilide was used in organic synthesis and medicine. In a small dose, Acetanilide had analgesic, antipyretic and anti-inflammatory effects. However, due to its toxicity, it was rarely used [1]

I.2. Principle of recrystallization method

Recrystallization method was used to separate and purify solids. The recrystallization method was based on marked differences in the solubility of substances in a solvent (or mixture of solvents) at different temperatures, or there was a difference in solubility between the main substance and the impurity substance at the same temperature.

I choosed a solvent which was comfortable with following conditions:

Solubility of the crystallizer was necessary.

Solvent dissolved the crystallizer.

The solvent did not react with the crystallizer.

The boiling temperature of the solvent must be greater than its substance.

Experimental Procedure

II.1. Making a saturated solution

Firstly, I created a solution which was mixture of distilled water and ethanol followed 7 : 3 respectively. Then, I placed 3 gram of acetanilide and 20 mL of the solution into an one necked round- bottom flask . Next, I assembled apparatuses. The condenser was inserted in the flask in vertical position, the bottom of the flask was far away 1cm from the heat source. Secondly, I took water flow through the condenser and pugged power. Therefore, a saturated solution was made. During the heating process, I observed the dissolution of acetanilide. When the solution in the flask was boiling for about 5 minutes, if the solids were large, a small amount of solvent was added through the condenser. I continued to cook until the solids and impurities were dissolved completely. Finally, I turned off the heat and let the mixture be cool.

II.2. Clarifying the saturated solution

The solution was clarified by Bucne funnel. A new solution was created. It was hot, clear and colorless. During the clarification, I must operate quickly to avoid premature crystallization in the flask.

The new solution was divided into two 100 mL beakers. One beaker was covered at room temperature. Other beaker was refrigerated with ice and stirred. During crystals appear in two cases, I compared and made conclusions. I combined the crystals in two beakers, cooled them with ice and clarified by Bucne funnel. The crystals were washed on the funnel with a small amount of cold water.

II.3. Drying product

Pure acetanilide was dried at 80°C. The product that was completely dry was determined its melting temperature and weight. Therefore, I could calculate efficiency of acetanilide refinement process.

Results

My result was indicated in figure 1:

Figure 1: Pure acetanilide

Figure 1 showed that pure acetanilide which I get in this experiment was white and lustrous.

The mass of pure acetanilide crystals was 2.3 gram.

The efficiency of the refining process was 76 %

Discussion

The result was obtained due to many causes:

The hot filtration process was not fast, leading to premature crystallization.

The process of cooling to crystallize was not complete, the time for crystallization was still short, leading to a part of acetanilide in solution.

Process of cooling filter was defective.

Practitioner made some mistakes.

Conclusion

The lab’s purpose was to refine acetanilide by recrystallization method.The experiment made a good result with quite high efficiency ( 76 %). I refined acetanilide by recrystallization method successful by making and clarifying a saturated solution, and then, crystals were made from the resulting solution.

20193.117436.Midterm1.NVA-04

Introduction

At a certain and constant pressure, the pure solvent begins to crystallize at its maximum temperature. This temperature and temperature remains constant throughout the crystallization process, ie when the liquid and solid phases are weighed equally. When adding a solute that does not evaporate and does not interact chemically with the solvent The starting temperature of the solution (crystallizing solvent, non-crystallizing solute) will be lower than the crystallization temperature of the pure solvent. . During crystallization, the concentration of the solution increases gradually (since only solvents can crystallize), the crystallization temperature of the solution decreases gradually until the solute begins to crystallize (at the same time with the solvent). , the crystallization temperature remains constant. Starting temperature The first crystallization (first appearance of solvent crystals) is called the crystallization temperature of the solution. The To - Tdd difference is called the decrease in crystallinity (ATL) of the solution. If the solution is dilute enough, the ATkt is proportional to the solute concentration.

Determine the molecular mass (M) of a substance (for example urea) by measuring the decrease in crystallinity of a dilute solution containing the solute in a suitable solvent (eg water).

Experimental procedure

1. Apparatus

Average Thermostats

Test tube bearing volume 100mL

Stirring rod

Beckmann thermometer

Copper air bag

2. Chemicals

Distilled water

Urea crystal

Table salt, ice used as a cooling mixture

3. Produce

- Assemble tools as shown in the picture

-Measure 25 mL of distilled water with a pipette into a dry test tube. Place the test tube in the ice cup.

- Weigh about 0.5 g of urea in the weighing bowl (use the analytical balance with an accuracy of 0.0001).

-Prepare the cooling mixture: mix crushed ice and salt.

Measure the crystallization temperature of the solvent

· Test tubes containing 25 mL of distilled water were placed in an air bag placed in the cooling mixture.Stir gently with your hand and monitor for lowering of temperature (use a magnifying glass). When the exact temperature can be read, continuously observe the solvent temperature (read every 15 seconds). Crystallization can be delayed by up to 1 - 2 ° C. The large crystalline point (large symbol because the measured temperature is the relative temperature) of the solvent is the temperature at its highest point and stops for a long time, (line (1), Figure 1). After determining the aloud value (read to accuracy to 0.002 - 0.003 ° C) remove the test tube from the bag, dissolve the water crystals with hand heat and stir).

· Repeat the measurement process as above until the values ​​of to do not differ by more than 0.003 ° C.

Measure the crystallization temperature of the solution

§ Pour the weighed urea into a test tube containing the solvent above through a dry funnel.

§ Weigh the weighing bowl again, thereby deducing the amount of urea dissolved in the solvent. Determining the crystallization temperature of the resulting solution is the same as determining the crystallization temperature of the solvent. The crystallization point of a solution is the temperature at which it peaks and then falls. Repeat the determination at (loud symbol because the measured temperature is relative) until the vessel values do not differ by more than 0.003 ° C.

Results

Mass of solvent (g) G= 25

Mass of solute (g) g = 0,4962

Solvent crystallization point Tdm = 3,33

Crystallization point of solution Tdd = 2,71

The cold solution constant of H2O K = 1,86

Results calculated:

M=K××=1.86××=59.5

Survey error:

Experimental error :

M=K××=>= ++

=++= 1.01%

Instrument error :

(∆g)dc=10^-4×0.4962=4.96×10^-5(g)

(∆G)dc=0.005×24=0.125(g)

(∆T)dc=0.01×0.62=6.2×10^-3(°C)

Discussion

In this case the crystallization point of the solvent is higher than that of the solution because when adding solute to the solvent, during crystallization the concentration of the solution will increase gradually (due to the crystallization solvent only), so the crystallization temperature of the solution gradually decreases until the solute begins to crystallize, then the crystallization temperature will start to stay the same

Not any two-component system is the same. If there is a chemical interaction with the solvents in the additive constituent, the crystallization temperature may be lowered or increased depending on the nature of the chemical interaction with the solvent.

Conclusion

Through the experiment we can determine the mass molecule of urea. The obtained results have identified errors and during data calculation and processing. If we increase the solute concentration, the error of T decreases, but the thicker the solution, the less accurate the application of the formula ∆T = K×m.

20193.117436.Midterm1.NVA-05

I . Introduction

Measure the saturated vapor pressure P of a volatile liquid at different temperatures, then make the Clausius - Clapeyron equation, determine the constants of this equation, then determine the vaporization heat and boiling temperature of liquid

The saturated vapor pressure of a liquid or solid is the pressure of that substance in a slightly equilibrium with liquid or solid.

* Clausius - clapeyron equation :

The vapor pressure depends on the temperature according to the clausius - clapeyron equation :

Inside :

P : is the saturated vapor pressure

T : is the temperature

R : is the ideal gas constant

Λ : is the heat of vaporization

II. Experimental procedure

Add the liquid to be measured into 2 branches A and B of the peach jar (push about 2/3) and connect with tube N. Place the flask in a cup of water placed in the thermostat (the flask should be submerged in water).

1. First of all, remove the air from liquid A in the following way :

Keep the thermostat at the desired temperature. Let the suction pump work and then open the J key, the gas pressure in the stabilizer C gradually decreases, the liquid level H₂ and M rise, the level H₁ and I₂ drop, to one At some point, the air bubbles from A escaped and pulled the air out.

2. When you are sure that the air above A has been expelled (the thermostat remains fixed), close lock J. Then open the key K and add the incoming air until the B₁ and B₂ levels are equal, then close the key K. Read temperature t, barometric pressure H, and the difference in pressure h.

III. The result

Test temperature t℃ = 29 ℃

H = 760 (mmHg)

IV. Discussion

- The graph is a straight line.

- From the above graph, we deduce a and B from which we can construct the Clausius-Clapeyron equation.

- Determine the constants of this equation and then determine the vaporization temperature and boiling temperature at normal pressure (760mmHg) of the liquid.

V. Conclusion

- From the experiment, we measure the saturated vapor pressure of a volatile liquid at many different temperatures, from which we make the Clausius -Clapeyron equation, determine the vaporization temperature and boiling temperature of that liquid.

20193.117436.Midterm1.NVA-06

Introduction:

Benzoic acid, C₆H₅COOH, is a white crystalline solid which is used chiefly as a food preservative, in the synthesis of dyes and in medicine. This aromatic acid usually appears as a needle shape and It is dissolved in alcohol, ete and hot water [1]. Because It is so easy to get some mixed impurities, this acid needs to be refined. The recrystallization method is commonly used to separate Benzoic acid from the mixture containing it. This method is mainly based on the solubility of the acid in a solvent at high temperature.

The aim of this experiment is to recrystallize Benzoic acid and determine the purity of the obtained acid . The expected acid melts 122,35 ° C like the manual temperature [2]

Experimental Procedure:

Experimental process is as follows [3]:

Firstly, I made the saturated solution. I placed in a 250 ml 1-necked round- bottom flask about 210 ml of distilled water and 5 grams of Acid benzoic to be refined. Then I installed the allihn condenser on the flask in a vertical position, the bottom of the flask is about 1cm from the electric stove . Next, I filled that condenser with water, turned on the electric stove and observed the dissolution phenomenon. When the solution in the flask boiled for about 5 minutes, I added a small amount of solvent through the condenser to dissolve the remaining solids. I continued to boil until the amount of solids did not change or they dissolved completely.

Secondly, the saturated solution must remove colors. Turning off the electric stove to cool the flask for a while, then I removed the condenser and put in the flask about 0.1 gram of the finely ground activated carbon. Then I installed the condenser and continued to boil the acid benzoic solution. That acid was hot filtered to remove insoluble impurities with a Busne filter funnel and I obtained a colorless hot solution.By using a vacuum filter and operating quickly, I avoided crystallization in the flask.

Third, the resulting solution is divided into 2 cups and I observed them : A cup was put alone and covered at room temperature. The other was stirred with ice besides.

After all ,both cups were iced ,then I filtered out the crystals with a vacuum filter . I washed the crystals right on the funnel with a small amount of clean cold water. The resulting product was dried at 80 ° C.After the product was dried, I determined the melting point, weighed the amount of dried product, and calculated the yield.

Results:

At normal temperature, the needle-shaped crystal was formed slowly, the process mainly took place on the surface of the liquid. The crystals were quickly formed with cold stirring. That process took place all over the liquid.

The amount of acid used : m1= 5,02 (g);

The amount of acid obtained : m2 = 2,87 (g)

Melting point : t1 =123 ° C ; t2= 125 ° C

Refined performance : n= ( m2 * 100) :m1 = (2,87 *100) :5,02 = 57,17 (%)

Discussion:

The temperature was important factor which directly affected to the success of this experiment.Because Benzoic acid sublimates over 100°C , I dried it at 80° C to avoid differently heating in the oven. That resulted in a long drying time. The actual melting point was 123°C to 125 °C. That was slightly higher than the theoretical temperature. Thus, the obtained benzoic acid is not a high purity. The cause may be an error in the operation.

Conclusion:

The purified benzoic acid is not high purity and low efficiency because I could not obtain all of the product after I manipulated.

20193.117436.Midterm1.NVA-07

Experiment 1: Recrystallization of Benzoic Acid

I. Introduction

Recrystallization is one of the popular methods for organic compound separation and purification. This process involves dissolving the impure solid to be purified in an appropriate hot solvent and filtering to remove insoluble impurities [1]. In this investigation, benzoic acid was purified by precipitation from a saturated aqueous solution of acid benzoic. The purpose of this experiment is to determine the factors which affected the purity of refined acid benzoic by using the recrystallization method. The result of the experiment is also a practical basis to compare the effectiveness of this method with other purification methods.

II. Experimental Procedure

1. Chemicals and apparatus

Acid benzoic, distilled water were available in the laboratory. All the chemicals were laboratory grade and used without further treatment.

Some apparatuses that were used in this experiment included: 1-necked round-bottom flask, glass beakers, petri dish, stirring rod, Allihn condenser, Büchner funnel. They were possessed by the laboratory and washed before being used.

2. Procedure

Five grams of solid acid benzoic and 210 ml of water were put into the 1-necked round-bottom flask. Then, this flask was heated until the solid was completely dissolved in water, and the solution was boiled in five minutes. After hot filtration, the obtained solution was divided into two parts, labelled 1,2. Part 1 was slowly crystallized at room temperature. Part 2 was crystallized by stirring at 0 – 5-degree C. After that, the collected crystals in two parts were filtered by the Büchner funnel. Then, they were dried at 80 degree C in 30 minutes until the mass of the crystals was constant.

A small amount of pure solid acid benzoic was stuffed in a capillary tube. Briefly, this tube was inserted in a digital melting-point apparatus (IA9100). We recorded the temperature value range when observing the crystal beginning to melt until being completely clear.

III. Results

Crystals in part 1 were formed slowly, had needle-shaped, while in part 2, it was smooth, milky white, and was fast formed.

The mass of acid benzoic after refining was 3,22 grams; therefore, the refined efficiency was calculated by the following formula [2]:

In which:

m₁: the mass of unrefined acid benzoic

m₂: the mass of refined acid benzoic

After observing the state of the solid in a capillary tube, the melting temperature range which was recorded is from 113 to 124 degree C

IV. Discussion

The measured melting temperature of refined acid benzoic is higher than the value of pure acid benzoic shown in the handbook (122,35 degree C) [3]. This error might be caused because the amount of impurities is still considerable in the product.

Firstly, the impurities can be jumbled in the crystallization process. Benzoic acid is crystallized owing to the discrepancy of water solubility between different temperatures. Also, the forming rate and shape of crystals depend on the stirring speed [4]. In this experiment, the crystals in part 2 had formation time double as short as in part 1. Even though the time was reduced, the entrained impurities in the stirring process could be substantial.

Besides, impurities can be confused during the manipulation of the experimental steps. For example, the crystals might be contaminated by filter paper, or it might not be completely dried out. Furthermore, the efficiency of the process was average, which indicated that the impurities content in the raw matter was considerable, and the material lost significantly.

Furthermore, when comparing with other solid-purifying methods as extraction or sublimation, the recrystallization process does not require stringent environmental conditions such as solvent, pressure…The result of the experiment could illustrate that this method could refine a great amount of impure solid, but it still assured the purity in an acceptable tolerance.

Lastly, the efficiency of the experiment could be improved if the crystals can be stored safely and allowed to sit for a few days. The experiment could be repeated to get an accurate assessment.

V. Conclusion

Importantly, the results provide evidence for the influence of environmental conditions on the crystallization process. Besides, although the errors caused in each experimental step was significant, the obtained efficiency was satisfactory. Therefore, the recrystallization is an effective method of purifying nonvolatile organic solids.

20193.117436.Midterm1.NVA-08

Introduction

Chemical equilibrium is the state of the reversible reaction system where the rate of the positive reaction equals the rate of the inverse reaction and the concentration of substances does not change. Equilibrium is a quantity that specifies the equilibrium state of the system. For response reversible in solution, it is common to use the equilibrium constant with concentration (K_C). Regarding thermodynamic, the equilibrium state corresponds to the stable state of the system, at which the Gibbs function (or equipotential function) G of the system reaches a minimum value.

Turning to this report, the aim is to reexamine the effect of the parameters on the equilibrium solutions and gases. According to Le Chatelier's principle or Equilibrium law, three common parameters that affect the equilibrium state are concentration, temperature, and atmosphere. In this experiment, we reexamined the first two parameters.

Experimental Procedure

Chemicals and apparatuses

(The chemicals are used as received and without further treatment)

Saturated FeCl₃ solution

Saturated NH₄CNS solution

NO₂ gas

Essential apparatuses: Test tubes, communicating vessels, etc.

Effect of concentration

First of all, a drop of saturated FeCl₃ solution and a drop saturated NH₄CNS solution was added to a small cup with about 20 ml of distilled water inside. Then the solution was put into 4 test tubes (about 1ml for each tube or 10 drops of solution height in a test tube). Tube 1 was remained as a sample, while tube 2 was added 1 - 2 drops of the saturated FeCl₃ solution. 1 - 2 drops of saturated NH₄CNS solution were added into Tube 3.

The colors of the solutions in the above test tubes were observed and recorded. The result was used for comparing with the principle of equilibrium.

Effect of temperature

Two test tubes were connected by a pipe to make an apparatus called communicating vessels (described in Figure 1) containing NO₂ gas, which is reddish-brown. The valve K was unlocked on the pipe so that the colors of the two tubes are equal, then K has locked again. Tube 1 was dipped into a cooling mixture of ice and table salt. Then Tube 1 was immersed in hot water. Tube 2 was used for comparison. The color change of Tube 1 in ice and hot water was observed and recorded. The result was used for comparing with the principle of equilibrium.

Figure 1. An example of communicating vessels

Result

Effect of concentration

The color of Tube 1 remained the same. The color of Tube 2 and 3 turned darker instantly.

Effect of temperature

The color of Tube 1 in ice instantly became lighter than Tube 2 and disappeared. On the other hand, in hot water, the color of Tube 1 turned darker than Tube 2.

Discussion

Le Chatelier's principle claims that in an equilibrium system, if one of the parameters (temperature, concentration, pressure) changes, the equilibrium will shift in the direction against that change. In general, both experiments have proved the accuracy of the Equilibrium Law.

Effect of concentration

The reaction in this experiment:

Fe³⁺ + CNS^- ↔ Fe (CNS)²⁺ (1)

(bloody red)

In the result that was observed, when the concentration of equilibrium increased, the reaction moved toward the dimension that reduces the concentration of the substances. When the concentration of equilibrium reduces, the reaction moved toward the direction that increases the concentration of the substances. If FeCl₃ or NH₄CNS concentration increases, the equilibrium of the reaction (1) shifts in the forward direction, make the colors of the solutions darker.

Effect of temperature

The reaction in this experiment:

2 NO₂ ↔ N₂O₄ (2)

(brown) (no color)

In this experiment, when energy or heat was given to the reaction, the inversion of the reaction gives energy or heat. If the temperature decreases, the equilibrium will move towards the direction of the heat gain. If we lower the temperature, it will move against the direction of the heat gain. As a result, if the temperature increases, the equilibrium will shift in the inverse direction makes the brown darker. If we lower the temperature, the equilibrium will switch in the reaction direction that makes the color lighter. This reaction is exothermic.

Conclusion

To conclude, the statistic seems to exhibit the effect of conditions on the equilibrium state. The two experiments are used to reexamine the Le Chatelier’s principle or the Equilibrium law. The results proved the correctness of the law.

However, this experiment still contains some weaknesses. An apparent limitation of the experiment is the lack of samples, so the statistic is not accurate. This limitation would be fixed by having more sources of sample to analyze. Another limitation of this experiment is having no precise statistics in the result, which leads to lack of accuracy.

20193.117436.Midterm1.NVA-09

1. Introduction

Viscosity is the internal resistance of a liquid, which must overcome a force, with which the fluid flows.Viscosity plays an important role in manufacturing technology. Based on viscosity property, we can calculate the following problems such as: Molecular mass, structure, dissolved substance concentration, diffuser concentration and many other properties of polymeric substances. Solutions of polymers usually have a much greater viscosity than common solutions and colloidal systems of the same concentration. Therefore, the polymers used to measure the viscosity must be of dilute concentration

To study more about viscosity, I did an experiment with the main purpose is Determination of viscosity and characteristic viscosity of a solution of polyethylene glycol in water by research methods flux of solution through capillaries.

2. Experimental Procedure

a, Determine the flow time with the viscometer Ostwald.

The viscometer is placed in the temperature stabilizer, the water level should exceed level a of secion branch. Use pipette to add 10ml of liquid into first branch. Use the rubber tube connected to first branch to pump the liquid slowly up to excess a about 1 cm. Therefore, the liquid remains at the bottom of the large bulb in first branch . Allow the liquid to flow naturally and use the stopwatch determines how long the liquid will flow from a to b. Take the experiment for 3 times

Note: Before doing this experiment, we need to clean ostwald viscometer and rinse with distilled water.

b, Determine the viscosity of solution polyethylene glycol in water

The thermostat is set at a fixed temperature (preferably 20^oC or 25^oC) to determine the viscosity of pure water, and then determine the viscosity of solutions polyethylene glycol in water with the following concentrations, respectively:

C = 5;8;10;15;20;25 % mass of polyethylene glycol in water.

Note: Take each measurement 5 minutes apart to stabilize the whole solution.

c, From the results of the experiment, apply the formula to calculate: Viscosity (η), specific viscosity ( η_r ) and characteristic viscosity ( [η] ) of polyethylene glycol in water.

From experiment, we have: Flow time from a to b of polyethylene glycol (t), water(t_o), viscosity of water (η_o).

Determination of viscosity (η)

Replace η_o,t,t_o to formula 1:

η = η_ot/t_o (1)

Determination of specific viscosity (η_r )

Replace η,η_o to fomula 2:

η_r = η/η_o-1 (2)

Determination of characteristic viscosity ( [η] )

The characteristic viscosity is defined as:

So, the relation between η_r/C and C is as follows:

η_r/C = [ η ]+ aC

η_r/C is the first order function of C, plotting the dependency η_r/C according to C. I Will find the specific viscosity [ η ].

3.Results

After doing experiments and calculations, I have the wrong parameters and table.

Laboratory temperature (⁰C ): 37 ^oC

Water viscosity at laboratory temperature: η_o = 0,7 ( mN.s/m² )

Flow time of water: t_o = 23,16 (s)

Table 1: Determine the viscosity of solution polyethylene glycol in water.

Graph 1 : The dependence of the characteristic viscosity on the solution concentration

From the graph show that ηr/C= f© is a linear function that increases as solution concentration increases and at C is close to 0 then [ η ]= ηr/C = 3,2

4. Discussion

Due to the characteristic viscosity of polymeric solutions increasing with concentration. In the small concentration range, the characteristic viscosity increases proportionally with the high molecular substance concentration, so in the graph we can see the straight line of the dependence of the characteristic viscosity on the increasing concentration.

Experimental results are relatively similar to reality, uncertainly in the allowed area.

Uncertainly due to: Instruments, measuring process, taking calculation error

5. Conclusion

After setting up the table and graphing I found specific viscosity [ η ] = 3,2 when A reached 0.

The measurement of polymeric solutions viscosity plays a very important role, helping to determine the molecular mass, structure and many other properties of polymeric substances.

Solutions of polymers usually have a much greater viscosity than common solutions and colloidal systems of the same concentration. Therefore, the polymers used to measure the viscosity must be of dilute concentration.

20193.117436.Midterm1.NVA-10

Introduction

A titrate reaction is one of the most popular methods to identify the concentration of the solutions in analytical chemistry. The current experiment aimed to determine the concentration of HCl solution by a standard Na₂B₄O₇ solution and to find the tested NaOH solution concentration by the HCl solution. The used titrate reactions were acid – base reactions, and we used different indicators to know the endpoints (Phenolphthalein, Methyl red, Methyl orange). The indicators, due to their changed color range, were chosen to comply with the pH hops of each reaction. The results were calculated following the law of the equivalence effect.

Experiment Procedure

Materials and apparatus:

All chemicals used in this experiment like pure Sodium tetraborate decahydrate powder (Na₂B₄O₇.10H₂O), concentrated HCl solution, tested NaOH solution, indicators (Phenolphthalein, Methyl red, Methyl orange), and distilled water were available in the laboratory. Needed apparatus included one 50ml – burette, one 10ml - pipette, one 100ml – beaker, one 500ml – beaker, two 100ml – Erlenmeyer flasks, one 250ml – volumetric flask, and one rubber bulb.

Experiment procedure:

This experiment had two parts: First, we titrated the HCl solution by the Na₂B₄O₇ solution. Second, the tested NaOH solution was titrated with the same HCl solution in the previous part.

In the first part, two solutions – 250ml of a standard 0,1N Na₂B₄O₇ solution and 500ml of an approximate 0,1N HCl solution – were prepared. An estimated amount of 4,7675g Na₂B₄O₇.10H₂O powder was moved into the volumetric flask and dissolved with warm distilled water. About 4,2ml of the concentrated HCl solution were taken in the fume hood to dilute. Then, the HCl solution was titrated by 10ml of Na₂B₄O₇ solution, which was put into an Erlenmeyer flask with 2-3 drops of Methyl red. The titrate reaction stopped when the color of the Na₂B₄O₇ solution changed from yellow to red.

In the second part, the tested NaOH solution was titrated by the prepared HCl solution. 10ml of NaOH solution, which had 7-8 drops of Phenolphthalein or 1-2 drops of Methyl orange, was titrated by HCl solution until the color of the NaOH solution changed. If the solution contained Phenolphthalein, the reaction stopped when the color changed from pink to colorless; the solution with Methyl orange had yellow color converted to orange.

Every value of the HCl solution was noted down after titrating. Each part was conducted at least three-time with volume differences were no more than 0,2ml. Apparatus must be cleaned by distilled water and dry before starting the experiment. This procedure followed the instruction given in the manual lab.

Results

The results observed in this experiment are summarized in Table 1 and Table 2. Table 1 shows the results of the HCl solution titrated process. This process led to the HCl solution concentration was 0,09213N. Table 2 shows the dissimilar results when the tested NaOH solution was titrated by the HCl solution using two indicators. Because of the differences, the NaOH solution concentration was respectively 0,1219N (with Phenolphthalein) and 0,1250N (with Methyl orange).

Table 1: The HCl solution volume after three-time titrated with Na₂B₄O₇ solution.

Table 2: The HCl solution volume after three-time titrated with NaOH solution.

We used two formulas to calculate the results, one to determine the standard Na₂B₄O₇ solution concentration, the other to find the HCl solution, and the NaOH solution concentration.

The formula calculated the concentration of the standard Na₂B₄O₇ solution was:

N_Na2B4O7 = 0,010015(N)

In which: m = 4,7746 (g), the weight of taken Na₂B₄O₇ powder

E_{Na2B4O7.10H2O} = 381,40 (g/mol), equivalent weight of Na₂B₄O₇.10H₂O

V_o = 250 (ml), the volume of the volumetric flask

The formula determined the solutions concentration after the titrate reactions was:

N_solution (N)

In which: N: the known concentration of solutions (Na₂B₄O₇ solution, HCl solution)

V= 10ml, the titrated volume of solutions (Na₂B₄O₇ solution, NaOH solution)

Discussion

The titrating HCl solution by the standard Na₂B₄O₇ solution gave three results suitable with tolerances (the volume differences were no more than 0,2ml). The color-changing from yellow to red of Methyl red was exact with the guideline in the lab manual. So, this stage of the experiment had no unexpected problems.

The identification of the tested NaOH solution concentration, however, had unlike results. The volumes of HCl solution, when titrated the NaOH solution by two indicators, had 0,34ml differences, it led to the difference between two results was 0,0031N. The indicators’ color changed like expected. Although the instructor had agreed with the results, the process needed to rework more carefully.

Possible source of errors in this experiment contains both objective and subjective reasons. Objective reasons include instrument errors, the quality of used chemicals like indicators, or the tested NaOH solution. Subjective reasons are lack of practice, the limited ability of color discrimination and round off the calculated results. All of these sources affect the results and they should be discarded.

Conclusion

The results of this experiment clearly show how a titration reaction proceeded. Any minor mistake would influence the experiment. The law of equivalent formula can be used to determine the concentration of a solution. Indicators are used to recognize the end of the reaction.

20193.117436.Midterm1.NVA-12

Introduction

An alloy is a compound combined from metallic and non-metallic elements to form a better substance. One example of an alloy is brass. Brass is an alloy of copper and zinc. Blending ratios between copper and zinc give us a wide variety of different brasses. Brass is an alternative alloy, it is used in many areas such as decorations, welding materials, and electrical equipment.

The aim of the current experiment was to determine copper in its alloy.

The copper content was determined by using differential optical absorption spectroscopy method. When the analytical solution has a large concentration (A value will be very large), differential absorption spectroscopy method is used to reduce the measured value A, with the use of an empty solution containing substance. The concentration is needed to analyze, because then the measured value will be in a linear range [1].

The content of copper will be determined by one-point method and basing on properties of the optical absorbance. The concentration of solution to be analyzed was C and the concentration of comparison solution was C₁. The absorbance of solution to be analyzed was A, the absorbance of the comparison solution was A₁. Lambert- Beer Law: A= εbC. The measured value of absorbance was A’

A’ = A – A₁ = εb(C – C₁) = K(C – C₁) [1].

The copper content which determined by differential optical absorption spectroscopy method was in the form of Cu[(NH₃)₄]²⁺ complex solution. It had absorbance maximum at λ _max = 605 nm. The control solution and the standard solution have had the same slope K so I can find the concentration of the sample solution [2].

Experimental Procedure

The chemicals and apparatuses:

Cooper (II) sulfate (CuSO₄), Amonium hydroxide (NH₄OH) and distilled water were available in laboratory. Used apparatuses were volumetric flasks, pipette, cuvette and beakers. The absorbance of solutions was measured by the UV-VIS Spectrophotometer.

Procedure:

First, we washed apparatuses with tap water and rinsed them with distilled water.

After that, we prepared three Cu [(NH₃)₄]²⁺ complex solutions: The standard Cu [(NH₃)₄]²⁺ solution was prepared by using a pipette to add 5ml of 0.15M to volumetric flask, labeled 01,using graduated cylinder to add about 5ml of NH₄OH, distilled water was added so that the volume of standard solution was 25ml. For the blank solution, a pipette was used to add 5ml of 0.1M CuSO₄ to volumetric flask, labeled 02. Then we used graduated cylinder to add about 5ml NH₄OH, distilled water was added so that the volume of standard solution was 25ml. For the control solution, used a pipette to add 5ml CuSO₄ from sample of copper (II) sulfate CuSO₄ into volumetric flask, labeled 03. Then we used graduated cylinder to add about 5ml NH₄OH, distilled water was added so that the volume of standard solution was 25ml. The stopper of the three flask was placed and mix the solutions well by inverting and swirling.

Next, we used UV-Vis spectrophotometer to measure the absorbance of three solutions in flasks 01; 02; 03. We turned on the instrument and set up the experiment you want to perform in the spectrophotometer software .The absorbance of three solutions was determined at the wavelength of 605nm. We added solution to the cuvette until enough for light to pass through, do not waste solution or risk spills by over-filling the cuvette. For each subsequent measurement, empty and rinse cuvette, shaking out as much of the rinse solvent as possible [2].

Result:

The results in this experiment are summarized in Table 1.

Table 1. Concentration and absorbance of three solutions

Calculation of data: M_Cu²⁺= 63.546 g / mole

C3 = 0.135M

C_Cu²⁺= C₃. M_Cu²⁺= 0.135 .63.546= 8.579 g/l

Discussion:

Measurement of absorbance at large values was prone to error in determination of concentration. Using differential optical absorption spectroscopy method helped reduce it because blank solution was a solution of known concentration replaces distilled water.

The wavelength at 605 nm was chosen as the incident ray wavelength because the complex has an absorption maximum at 605 nm. Measured at 605 nm for high sensitivity and accuracy.

Possible sources of error in this experiment include making mistakes during dilutions and other types of calculations and spilling chemicals during transfer, process of estimating a measurement. Errors occurred when we added solution to beakers or flasks and pipette. This experiment could be improve by repeating each measurement and using carefully apparatuses.

Conclusion:

From the results of above, this experiment shows the quantity of copper in brass: 8.79 g/l. We can conclude that the aim of this experiment has being achieved. Differential optical absorption spectroscopy method was successfully used to determine the mass of copper in alloy.

20193.117436.Midterm1.NVA-13

Introduction

The continuous variation method has been used frequently to determine the formulas and formation constants of complexes since its introduction by Job. The purpose of the experiment was to determine the complex composition of Cu (II) – Nitrozo-R salt. The reaction equation between M and R reagent is

M+ nR ↔ MRn (1. 1)

Equilibrium constant is Keq = (1 .2)

From equation (1.2), if the concentration of R reagent increases, the concentration of the complex will increase. If a series of solutions are mixed, in which the total volume is changeless but the volume ratio changes, there will be a solution with the maximum complexity concentration. The total component concentration is the same in each sample. A solution with a maximum complex composition is a solution with the ratio of the concentration of M, R equal to their composition in the complex. The method is based on plotting the measured absorbance. The composition of the complex gets from the point of intersection of the tangents to the curve. The reaction equation between Cu2+ with Nitro-R reagent is

Acid medium Cu2+ + HR ↔ [CuR]^- + H⁺

Weak acidic and neutral medium Cu2+ + 2HR ↔ [CuR₂]^- + 2H⁺

Base medium Cu2+ + 3HR ↔ [CuR₃]^- + 3H⁺

When using the method of continuous variation, we need attention:

Ensure the correctness of Lambert-Beer law.

Keep the ionic force constant and pH stable (buffer solution can be used).

Experimental Procedure

2.1 Chemicals and Apparatus

10 ^-3M Nitrozo-R-sol, 10 ^-3M CuSO₄, 0,05M Na₂SO₄, H₂SO₄ with pH=4 were used in the experiment. All solutions were available in the laboratory. UV-VIS spectrophotometer (Shimadzu UV-2600) was used to measure the solution absorbance.

2.2 Procedure

Firstly, nine volumetric flasks with volume 25 ml were prepared to precise dilutions with the following composition:

Table1. The volume of solutions

The solution was diluted to the mark with H₂SO₄ solution, again ensuring that the solution was thoroughly mixed.

Secondly, we measured the absorbance of each solution with a cuvette 1 cm thick at גּ_max= 490nm. The blank sample is distilled water. Finally, the relationship graph of A with V_Cu2+ was built. We found the maximum absorption and determined the composition of the complex. This procedure followed the experiment guide line.

Results

The results are summarized in Table 2. The data are presented graphically in figure 1

Table2. The absorbance of each solution

The solution was diluted to the mark with H₂SO₄ solution, again ensuring that the solution was thoroughly mixed.

Figure 1: A continuous variations plot

From the experimental data obtained:

A_max=0,737, V_Cu2+=5 ml, V_Nitrozo-R=5 ml

The chemical coefficient is the ratio of volume Cu²⁺ to volume Nitrozo-R. From the graph, the chemical coefficient is n==1. Therefore, the chemical formula is CuR.

Discussion

Figure 1 shows the relationship between absorbance and volume Cu2+. The concentration of R reagent increases, the absorbance increases, because the concentration of the complex increases. The complex formation of Cu2+ and Nitrozo-R depends on the pH of the medium and concentration. In an acid medium, CuR complex mainly exists. In a weak acidic and neutral medium, CuR2 complex mainly exists. In a base medium, CuR3 complex mainly exists. The existence of all forms of the complex depends on the medium. Experimental results are true with the hypothesis. When determining the complex composition, the medium should be closely examined. Possible sources of error in this experiment include solution preparation error, instrument error, and graph error.

Conclusion

Experimental results are true with the hypothesis. The method of continuous variation is very popular to determine the chemical coefficients in complex reactions. The complex formation depends on the pH of the medium. Through this experiment, I learned how to measure the absorbance with a spectrophotometer and determine the complex composition.

20193.117436.Midterm1.NVA-14

Titrating sodium hydroxide with hydrochloric acid

Introduction

In the laboratory, the reaction of an acid with a base to make a salt and water is a common reaction. The overall goal of this experiment was to determine the concentration of the sodium hydroxide by titration. “A titration is a technique where a solution of known concentration is used to determine the concentration of an unknown solution” [1]. The reaction will be observed as follows:

NaOH + HCl = NaCl + H₂O

The method of this experiment is based on the titration of a strong acid with a strong base. To do this titration, we have to determine the Equivalence Point in which the acid has completely reacted with the base by using an indicator. The indicator will change color at the endpoint. Phenolphthalein indicator and methyl orange indicator is used for determining neutralization occurs.

Experimental Procedure

1. Materials

Materials for a Titration Procedure are:

Burette, pipette, Phenolphthalein indicator and methyl orange indicator, Erlenmeyer flask, Hydrochloric Acid, and sodium hydroxide.

All chemicals were used without further purification.

2. Titration Procedure

Before each experiment, we rinsed the burette with the standard solution (HCl), the pipette with the unknown solution (NaOH), and the Erlenmeyer flask with distilled water.

Experiment 1: Titrating with a phenolphthalein indicator

First, the standardized solution (0,0914 (N) of Hydrochloric Acid) is poured into the burette adjusted the level of the liquid to exactly zero marks before titration. Second, an accurately 10,00 ml of the sodium hydroxide was moved into the Erlenmeyer flask using the pipette, along with 7-8 drops of Phenolphthalein indicator. The color of the solution was reddish violet. Then titration by Hydrochloric Acid until the indicator changed color from reddish violet to colorless, recorded the value on the burette. Titration was performed at least three times, take the average volume.

Experiment 2: Titrating with a methyl orange indicator

First, the standardized solution (0,0914 (N) of Hydrochloric Acid) is poured into the burette and adjusted the level of the liquid to exactly zero marks before titration. Second, an accurately 10,00 ml of the sodium hydroxide was moved into the Erlenmeyer flask using the pipette, along with 1-2 drops of and methyl orange indicator. The color of the solution was yellow. Then, titration by Hydrochloric Acid until the indicator changed color from yellow to orange, recorded the value on the burette. Titration was performed at least three times, take the average volume.

Result

We describe the result of titration which showed the color-change of the indicator at the endpoint. Color of the solution was colorless when we used Phenolphthalein indicator and orange when we used methyl orange indicator

The results are summarized in Table 1:

Table 1. The dependence of the volume of Hydrochloric Acid on indicators

Using phenolphthalein indicator: N_{NaOH =} = 0,09778 (N)

Using methyl orange indicator: N_{NaOH =} = 0,09752 (N)

Discussion

The concentration of the sodium hydroxides which are given by the teacher are 0,105 (N) (using phenolphthalein indicator) and 1,07 (N) (using methyl orange indicator). Compare with my results, the errors of the experiment are 6,8 % and 8.8%. Because of the experiment, I make some mistakes. One of the reasons which cause errors in the titration is the unclear change of color in the indicator. Therefore, we need to pay attention to avoid over-titration of the equivalent point. Besides, several factors can cause errors in titration, such as misreading volumes and faulty technique. If the reading is taken from the higher sections of the burette or the pipette, the volume measurement will be in error. Therefore, we have to read the liquid level using the bottom of the meniscus at eye level.

In this experiment, we chose the Phenolphthalein indicator and methyl orange indicator because the pH leap of titration with error ±0.1% is from 9.7 to 4.3. Meanwhile, the color-change interval of the phenolphthalein indicator is 8-10, and the color-change interval of the methyl orange indicator is 4,4-6,2. Therefore, both phenolphthalein and methyl orange could be indicators for the reaction.

Conclusion

The titration of a strong acid with a strong base is the simplest of an analytical chemistry experiment. This will allow information about the unknown solution to be determined. The experiment makes some mistakes causing errors in titration such as over-titration of the equivalent point, misreading volumes …so we need more care.

20193.117436.Midterm1.NVA-15

Experiment 1: The saturation curve method for the determination of the complex composition Cu (II)-Nitrozo-R salt

I. Introduction

One of the basic methods of the photometric method for the determination of the complex composition is the saturation curve method. The characteristic of the method is based on a solution's composition-property graph to determine the chemical coefficients.

The aim of the current experiment was to determine the complex substance composition Cu (II) - Nitrozo-R salt. The reaction creates a complex substance:

Cu²⁺ + nR ↔ CuR_n

Equilibrium constant is Keq = . With “n” is the ratio coefficient when the complex substance reaction reaches equilibrium.

To determine “n”, the saturation curve method is used. In this method, the concentration of one constituent is constant (this is usually a metal ion) while the concentration of the other is changed. Then, we measure the absorbance of the series solutions and graph the relationship A-C_R/C_Cu2+. If the complex substance is stable, the obtained graph will have a break point corresponding to the abscissa C_R/C_Cu2+= n. [1]

The Cu metal ions and HR reagent act together for complexes depending on the environmental conditions and concentration. In the acidic environment, there is mainly CuR complex. In weak acidic medium or in the neutral medium, CuR₂ complex mainly exists. In the base medium, CuR₃ complex are mainly formed. [1]

II. Experimental procedure

Chemicals and Instruments:

Chemicals used were Nitrozo-R-sol of 10 ^-3M, CuSO₄ of 10 ^-3M, Na₂SO₄ of 0,05M, H₂SO₄ with pH=4. All chemicals were available in the laboratory. In which, Nitrozo-R-sol and CuSO₄ were used to form complex substance. A 0,05M Na₂SO₄ solution was used to maintain a constant ionic force because it was a strong electrolyte with a large concentration. In addition, sulfuric acid with pH=4 was used to keep the pH stable.

Used instrument was a spectrophotometer with a cuvette of 1 centimeter to measure the absorbance.

Procedure:

We prepared eight 25 ml volumetric flasks to mix a series of solutions with the following composition:

Table 1. Amount of the solutions for the experiment

The solution was added to the volume mark and mixed well

Then, we proceeded to measure the absorbance of the complex substance at wavelength was 490nm with the blank sample of distilled water. The solution absorbance was recorded by a spectrophotometer.

Next, we graphed the relationship between absorbance and composition.

From this graph, we computed and determined the composition of complexes according to the saturation curve method.

This procedure followed the documentation’s instructions. [2]

III. Results

The results obtained in this experiment are presented graphically in Figure 1.

Figure 1. Diagram of “Absorbance-Composition”

As the results, when the Nitrozo-R-sol volume increased from one to eight milliliters, the absorbance increased from 0,165 to 0,532. The maximum absorption was 0,532.

From this graph, we had a ratio coefficient n= = 0,96 ≈ 1.

Consequently, the chemical formula was CuR.

IV. Discussion

From the experimental results, when the concentration of HR reagent increased, the absorbance increased due to the increased amount of complex. In addition, as the volume ratio of R and Cu augmented from 0,25 to 1, the curve augmented rapidly. When the volume ratio of R and Cu increased from 1 to 2, the curve was almost parallel to the horizontal axis. Because at the time HR acts sufficiently with metal ions, the absorbance does not increase anymore (the reaction reaches saturation).

Figure 1 shows the composition of the complex substance Cu (II) - Nitrozo salt:

was one. This is expected since complex substance formation is dependent on environmental conditions and concentration. In acidic environment, complex substance exists mainly in the form of CuR.

Possible sources of error in this experiment include instrumentation error, inaccuracy in graphing and the H2SO4 solution added was not at the correct volume mark.

V. Conclusion

To sum up, the results of this experiment may have error, however it supports the hypothesis. It shows that the complex substance composition Cu (II) - Nitrozo-R salt is ≈ 1 in acidic environment by using the saturation curve method. From this experiment, I knew how to measure the absorbance with a spectrophotometer, practiced skills in the lab and applied photometric method especially the saturation curve method to determine the complex composition.

20193.117436.Midterm1.NVA-16

Limited miscibility of two liquids.

I. Introduction

The purpose of this report is to determine the critical miscibility of two liquids at different temperatures. In this report, I built a temperature - composition diagram, then used the diagram to determine composition of any heterogeneous mixture of those two liquids.

Limited miscibility of two-liquid systems lie between the miscibility of completely dissolved systems and the miscibility of completely insoluble systems. Such some systems are phenol-water, buthanol-water, trimethylamine-water, etc.

For the phenol-water system at fixed temperature, when adding phenol to the water, at first phenol is completely mixed in water, forming a single (homogeneous) phase. If phenol is added up to some concentration, they no longer mixe and the system is splited into two phases: the water-saturated phenol phase (lower) and the phenol-saturated water phase (upper). These two phases of liquid are called conjugate, and when shaken vigorously, the mixture mixes and becomes turbidity.

At each temperature, the solubility of phenol in water and of water in phenol is a fixed value. As the temperature increases, solubility increases.

The graph below shows the affection of temperature to composition

Image 1: Composition-temperature graph of the phenol-water system.

The curve aKb divides the graph into two parts: the area inside the curve represents the heterogeneous system (2 phases) while the other one corresponds to the homogenous system. The curves aK and Kb show the effect of phenol on water (water phase) and water on phenol (phenol phase), respectively. K is the limited solubility point where the composition of the two phases is equal. Tc is called the critical dissolution temperature.

A composition-temperature diagram can be built using one of the two following ways:

Isothermal method:

Keep the temperature of the system constant, change its composition many times by adding gradually phenol to the water. Then determine at which composition the state of the system changes from homogenous to heterogeneous phase, and vice versa.

Shake the vial containing these two liquids vigorously and soak it into a thermostatically fixed thermostat until the mixture is completely divided into two phases (layers). Then quantitatively analyze these two layers.

Multi-heat method:

For mixtures containing m components (the system is still turbid), gradually increase the temperature until the mixture becomes clear. If the temperature is continued to rise, the mixture remains clear. Thus, base on the temperature at which the mixture starts to clear or turbid to determine the point b.

Repeat the experiment with mixtures of different compositions will eventually obtain the aKb curve.

In this report, we used the multi-heat method.

II. Experimental procedure

1. Building the phase diagram of the phenol-water system

“First, we took 10-12 test tubes containing different compositions of the phenol-water mixtures. Next, we immersed the test tubes in a glass of water and heat it up gradually. Then, the system was still very turbid, we gave the temperature a quick rise and being stirred gently. After that, we observed to note that when the mixture is about clear, the temperature must be risen very slowly and the mixture must be stirred more vigorously. Finally, we recorded the temperature at which the mixture started to become clear, then we let the temperature decrease slowly and observed to record the temperature corresponding to the mixture state starting to be cloudy.

These two temperature values must differ very little, not more than 0.5 degrees C

We took: t_average = (c_lear+ c_loud)/2

Since we didn’t immerse the thermometer directly in the mixture, it must be near the transition temperature, so the increase in temperature must be very slow to be considered that the temperature on the thermometer was equal to that of the mixture.”

2. Determination of the composition of the phenol-water system.

“By performing the above experiment, we found out the temperature. Base on the temperature-composition graph, we determined the composition of the mixture.”

Results

Figure 2: The temperature- composition graph of phenol- water system.

The effect of temperature variation on the degree of miscibility in this system was described by the graph of temperature versus composition at constant pressure.

The vertical axis on the graph represented the temperature and the horizontal one represented the corresponding composition.

IV. Discussion.

The graph is an upside-down ring. We saw this graph that when the concentration of phenol was increased from 10% to 35%, the miscibility increased and consequently the miscibility temperature increased from 42ºC to 67.5ºC. However, further increase of concentration of phenol from 35% to 70% made the miscibility decreased and consequently the miscibility temperature decreased.

After the experiment, we can build up the temperature-composition diagram of any two-liquid system.

During the experiment, there were errors in the results. They include instrumentation error, faulty measurement process. And errors were occurred when the experimenter read incorrect the clear starting temperature and the cloudy starting temperature.

V. Conclusion

After doing this experiment, students are provided with theory and practical knowledges on the topic of limited miscibility of two liquids. This is an important basic knowledge of physics and chemistry. It is a prerequisite for helping students grasp better other subjects of chemical engineering.

VII. Appendix

Data and experimental results.

The test temperature: 22^oC

Table 1. The experimental result.

20193.117436.Midterm1.NVA-17

The Method of Continuous Variation: Determination of complex composition Copper (II)-Nitrozo-R Salt

Introduction

The method of continuous variations, also called Job’s method, is used to determine the stoichiometry of a metal-ligand complex. We prepare a series of solutions such that the total moles in each solution is the same. The concentration of the complex is determined by the limiting reagent, with the greatest concentration occurring when the metal and the ligand are mixed stoichiometrically. “Nitrozo-R salt (symbol HR) is the common name of the two sodium 3-hydroxy-4 nitroso-2,7 naphthalene disunfonic acid-salt. Meⁿ⁺ metal ion and HR reagent react together for complexes MeR, MeR2, MeR3. They depend on the environmental conditions and concentration. In the acidic environment, it mainly exists in the MeR form. In the weak or neutral acidic environment, it mainly exists in the MeR2 form. In the alkaline environment, it mainly exists in the MeR3 form. Because the existence of complex depends on the environment so the determination of complex composition form must be checked prudishly” [1]. The absorbance (A) of a complex was measured and the maximum of absorbance (Amax) is a value that the ratios of Cu²⁺ and R are equal. The purpose of this experiment was to determine the complex composition of Copper (II)-Nitrozo-R salt. One way to investigate it was the method of conituous variation. The main advantage of this method is common and exact. The general conclusion is the main complex components in salt exist in the MeR form.

Experimental procedure

Chemicals

All chemicals are used as received. Chemicals are labeled in the laboratory. There are four solutions. 1,00x10-3M stock solution of copper complex was prepareed from 0.2g CuSO4.10H20, and H2SO4 pH4 to nearly 1 liter. Concentration of CuSO4 had to check by titrating PAN and adjusting to obtain 10-3M concentration. Nitrozo R salt was prepared from 0.373g pure Nitrozo-R and distillated water to 1 liter. H2SO4 pH4 was prepared from 1ml 0.1M H2SO4.and nearly 2 liter distillated water. PH of solution was adjusted to obtain 3.9 +/- 0.1 by adding drop by drop 0.1M H2SO4 or 0.1M NaOH. 0.05M Na2SO4 was prepared by 16.1g Na2SO4 .10H2O, 7.7g anhydous Na2SO4 and 1liter distillated water. It adjusted to pH4.0+/-0.1 by adding drop by drop 0.1M H2SO4 or 0.1M NaOH.

Experimental setup

The first, the experiments were performed with a series of nine solutions. They were prepared in 25 mL volumetric flasks. The composition of the nine solutions were presented in table:

Each flask is made up to volume mark and shake well

The second, the absorbance was measured at 𝞴 = 490 nm with the blank solution (distilled water) by UV-Vis. A continious variation plot depends on the absorbance (A) with the volume of Cu²⁺. They were found Amax and determined the composition of the complex. Used apparatus were 25mL voumetric flasks and 1,2,10 mL pipet.

Results

The results obtained in this experiment are the absorbances each volumetric flask. We summarized in table 1. This data is presented graphically in Figure 1.

Table 1. The absorbance of each solution in each flask.

Each flask is made up to volume mark and shake well

The results of Table 1 build the continious variation plot. 2 tangents are constructed. They intersect at one point. From that intersection, we descend on the horizontal axis and obtain the volume of Cu²⁺

The determination of VCu²⁺:

VCu²⁺ = 4,8ml

The determination of n:

n = (10-4,8)/4,8 = 1

The form of main complex components in salt exist in CuR.

Figure 1. The relationship between volume of Cu²⁺ and A

Discssion

Figure 1 shows the relationship between volume of Cu²⁺ and A. This relationship is summarized by the graph. The performance line gradually increased to maximum and then decreased. If the concentration of the initial solution is changed, on the absorbance component graph of the solution, the maximum positions remain identical. The displacement occurs only when the pH of the solution is changed. Errors in this experiment process include can be attributed to measurement of solutions in flasks and from solution acquisition, graphing and tangent computer drawing. In addition, the error is due to the calculation.

Conclusion

The results of the experiment cleay show that for the main complex form and survey the method of conituous variation. At 𝞴=490 nm with Amax of the ingredient form is the same. Since when the absorbance is maximum, the sensitivity and accuracy are highest, we can measure the smallest analyte concentration and the obtained results are clearest and most accurate. Base on the method conituous variation, the main complex components in salt exist in the CuR form. In conclusion, the concentration of the complex is determined by the limiting reagent, with the greatest concentration occurring when the metal and the ligand are mixed stoichiometrically.

Further reading

There are many ways to determine the composition of a complex. This method is popular and easy to manipulate but has some limitations. for example, it is the process of the continious variation plot and determining maximum absorbance. Construction of a continuous variation plot for each condition leads to the deduction of the molecular formula of the complex composition of Copper (II)-Nitrozo-R salt. This experiment demonstrates the concept of stoichiometric relationships between ions in the formation of compounds and serves to give students practice in calculations involving limiting reagents.

20193.117436.Midterm1.NVA-18

PRECIPITATION TITRATION: DETERMINATION OF CHLORIDE ION

1.Introduction

1.1 Purpose:

The aim of this experiment is to determine the Chloride ion by two different methods, and then compare the results with each other.

1.2. Background:

Titration is a very common laboratory method of quantitative chemical analysis, it is used to determine the concentration of an unidentified-concentration substance. Silver Chloride is very insoluble in water and it is regarded as a precipitate. Firstly, with the Mohr’s method, as the Silver Nitrate solution is slowly added, a white precipitate of Silver Chloride is formed: [1]

Ag⁺ + Cl^- AgCl.

Indicative reaction:

2Ag⁺ + CrO₄^2- Ag₂CrO₄

When the Silver Chloride is all precipitate, the additional Silver ion react with the indicator. Basing on the turning red-brown of the solution, the equivalence point of the titration could be determined.

On the other hand, the determination of Chloride ion can be carried out by using an adsorbent indicator - Fluorescein. The equivalence point is typically signaled by a color change that is influenced by an adsorbent mechanism.

2.Experimental Procedure

Ensure that all the apparatuses were clean, the standard Silver nitrate solution was filled in the burette:

Mohr’s method: We pipetted a 10.00mL sample of NaCl into a conical flask and added about 0,5 - 1mL of Potassium Chromate 5%. After that, the sample was titrated with 0,1N Silver Nitrate solution that was filled in burette until the solution turned red-brown, then we stopped the titration and noted the value of Silver Nitrate solution volume that had reacted. We repeated the titration two more times.

The method using adsorbent indicator: We pipetted a 10,00mL sample of NaCl into a conical flask and added about 2 – 3 droplets of Fluorescein 0,5%, the solution then turned fluorescent. After that, the sample was titrated with 0,1N Silver Nitrate solution that was filled in the burette until the fluorescence faded away. The titration was stopped and then the value of Silver Nitrate solution volume was noted. We repeated the titration two more times.

3.Results

Standard:

V_NaCl = 10,00 (mL)

V_AgNO3 = 0,05N

Table 1: The measured volume of AgNO₃.

Results

Mohr’s method

Using adsorbent indicator

From the results in the above table, I can calculate the concentration of the Chloride ion which is supposed to be determined:

Table 2: The calculated concentration of Cl^-

4.Discussion

As what is being shown in the tables, the results of the two different methods are quite close to each other. Agreed titre is within 0,1mL, the error can be caused by instrument error or titration movement.

The adsorbent mechanism can be explained by the following balance: [1]

HFl = H⁺ + Fl^-.

Before the equivalence point, the precipitate AgCl has the orientation to adsorb its own ion from the solution – Cl^-. The solution has the fluorescence of Fl^-. After the equivalence point, the odd Ag⁺ is adsorbed on the surface of the precipitate, the solution no more has the fluorescence.

The pH condition of the two titrations must be neutral: pH = 6,5 – 8,5 [2]. It’s because Silver oxide is formed at high pH, and Silver Dichromate is formed at low pH. This will reduce the concentration of Chromate ions, and delay the formation of the precipitate. In the second method, Fluorescein is a weak organic acid with pK_a=6,5, so the pH must be higher than 6,5 for the indicator to change the color. On the other hand, the pH must be not too high to prevent Silver Oxide from precipitating.

The concentration of the Ag₂CrO₄ must be not too high with the calculated value about 5% [2] in order that the Silver Nitrate has all precipitated when the solution suddenly turns red-brown. If the concentration is high, the red-brown Silver Chromate is formed during the Silver Chloride is precipitating, that is also cause the error of the titration.

5.Conclusion

Both two methods are relatively simple, they allow us to determine the equivalence point accurately. The mohr’s method is used more often because the chemical is easy to find.

20193.117436.Midterm1.NVA-20

Glass corrosion of HF acid

Introduction

The objective of this experiment was to determine whether HF acid can be stored in a glass bottle. The purpose was to understand the reasons that cause the chemical reaction of HF acid on glass and find out a engraving solution on the glass surface. I hypothesise that with the acidity of HF acid, it can’t be stored in a glass bottle and will cause corrosion effect.

Methods and Materials

The following steps were followed to conduct the HF acid on glass experiment:

Step 1: preparing glass paraffin solution: In this first step of the experiment, we need to prepare a paraffin solution, one piece of glass and an object which is sharp enough create a scratch on the paraffin surface of the glass which is a screw in this case.

Step 2: paraffin application on glass: The glass surface was applied a thin layer of paraffin and waited till dry. Using the screw, we try to engrave on the paraffin surface of the glass. Note that we must clearly engrave on the glass for the best results at the end of the experiment.

Step 3: test HF acid solution on glass: A pipette was used to take a small amount of 0.05M of HF acid solution and drip it into the etched glass plate. Then carefully covered that chemically treated glass and waited for about an hour to see the result.

Step 4: phenomenon happens: We took the glass out and rinsed it through water, scraped off the paraffin layer on the glass surface, observed the phenomenon happen.

Noted : Because HF acid is very toxic, gloves should be used when performing the experiment and do not put acid wire solution on hands.

Result

After performing the experiment, areas that were engraved which has no paraffin protection were left the scratches as we created. In contrast, areas of the glass that covered with paraffin didn’t appear to have any chemical reaction.

SiO₂ + 4HF → SiF₄↑ + 2H₂O

Discussion

HF acid solution has caused corrosion where it contacted directly with glass surface which were not protected by paraffin. These scratches were due to the reaction of SiO₂ with HF acid solution. This chemical reaction generated heat and engraved the scratches applied to it.

This experiment appeared to have the same result with previous theories where acid was in contact with glass material. This experiment went out smoothly, there were no problems and limitation encountered during the experience.

Conclusion

The result of this experiment indicates that a glass bottle could not be used to store HF acid solution because the acidity of HF solution can corrode glass material. In addition, we found an effective way to for engraving solution on the glass thanks to the glass corrosion of HF acid.

20193.117436.Midterm1.NVA-21

Determination the concentration of an Fe³⁺ solution by K₂Cr₂O₇

I. Introduction

Titration concept

During a titration process, the titration is added slowly to the substance to be titrated (with known volume) until the reaction is complete. From the volume of titration solution consumed, we can calculate the concentration of the substance to be analyzed. The titration is usually performed by a burette.

The Lab’s purpose

The purpose of the experiment was to determine the concentration of an Fe³⁺ solution by K₂Cr₂O₇. The objective of the Lab was to know how to titrate and determine the concentration of the Fe³⁺ solution by K₂Cr₂O_7.

Background

First, we reduced Fe³⁺ to Fe²⁺ by Zn:

2Fe³⁺ + Zn = 2Fe²⁺ + Zn²⁺

Reducing was performed in HCl acid environment. Then, we titrated the Fe³⁺ solution by K₂Cr₂O_7. In strongly acidic environments:

1 Cr₂O₇^2- + 6e– + 14H+ ⇋2Cr³⁺ + 7H₂O E⁰ (Cr₂O₇^2-/2Cr³⁺) = 1,36V

6 Fe²⁺ – 1e– ⇋ Fe³⁺ E⁰ (Fe³⁺/Fe²⁺⁾ = 0,77V

Cr₂O₇^2- + 6Fe²⁺ + 14H+ = 2Cr³⁺ + 6Fe³⁺ + 7H₂O

To know the right indicator, we needed to know the potential step of the process.

Example: We titrated 50ml FeSO₄ 0,1M solution by K₂Cr₂O₇ 0,01M, pH= 0. The potential jump with error ±0,1 corresponded to a volume of K₂Cr₂O₇ from 49,95ml to 50,05ml. When we added 49,95ml the K₂Cr₂O₇ solution:

E_dd = E⁰ (Fe³⁺/Fe²⁺) + 0,05916.log = 0,77 + 0,05916.log = 0.95V

[Cr³⁺] = 0,10000M × = 0,01666M

[Cr₂O₇²] = × 0,01000M = 5.10^-6M

E_dd= E⁰(Cr₂O₇^2-/Cr³⁺) + log= 1,36 + log= 1,34V

So the potential jump of the titration was 0,95V to 1,34V. During the reaction, H₃PO₄ was added to easily determine the equivalence point. Therefore, the new potential jump would be calculated. Assuming that under pH, we titrated α(Fe³⁺) concentration (α(Fe³⁺) = 10^-2).

E_dd = E⁰(Fe³⁺/Fe²) + 0,05916log = 0,68 + 0,05916log + 0,05916log10^-2

= 0,74V

So the potential jump of the titration in the presence of H₃PO₄ was from 0,74V to 1,34.

Hypothesis

If there were the diphenylamine indicator in solution, the solution would change from cyan green to purple- blue. Because the diphenylamine indicator has a potential value in the potential jump (E⁰ = 0,76V).

II. Experimental procedure

Materials and apparatus

Chemicals

Used chemicals were the diphenylamine indicator, the K₂Cr₂O₇ solution, a mixed solution (H₂SO₄ + H₃PO₄) 6N, Zn grain, a HCl 1:1 solution and the Fe³⁺ solution.

Apparatus

Used apparatus were a beaker, a copper pipe, a 100ml Erlenmeyer flask, a 250ml Erlenmeyer flask, a funnel, a filter paper, a burette, 10ml pipette, a rubber bulb, a hot plate, an electric stove.

All apparatus and chemicals were available in the laboratory. We took them because they were used for this experiment.

The procedure

The procedure for conducting the experiment is below. First, we rinsed the apparatus. Then we put K₂Cr₂O₇ solution into burette, expelled any air bubbles and adjusted to the “0” mark before titration. Next, we took exactly 10ml of Fe³⁺ solution into the conical flask, added about ten grains of Zn and 5ml of 1:1 HCl acid to the conical flask (working in a hood). Then we boiled it to almost until the color was completely lost. If a white precipitate appeared, we would add HCl to dissolve it. We did not boil it because the Fe compound evaporated and made the wrong result. After that, we spayed distilled water around the mouth of the conical flask to the volatile HCl on the wall to go down to the solution, quickly cooled the conical flask to room temperature. Then we filtered out the excess Zn with filter paper, collected a few large 250ml conical flask. Next, we rinsed the bottle with distilled water many times, until the volume of solution in the conical flask is 100ml.Then we added 5-7ml a mixture of two acids and 1-2 drops of the diphenylamine indicator. After that, we performed titration with K₂Cr₂O₇ until solution appeared purple – blue and recorded the volume of K₂Cr₂O₇ consumed. Last, we did it at least three times (deviation not more than 0.2ml) and took the average value.

III. Results

The solution changed from cyan green to purple- blue.

V_Fe3+ was taken for titration: 10ml

Table 1: V_K2Cr2O7 was determined in the titrations.

Concentration of Fe³⁺ solution:

N­_Fe3+ =N_k2Cr2O7. _K2Cr2O7 / V_Fe3+ = 0,05.17,83/10 = 0,089 (N)

C_Fe3+ = N_Fe3+.Đ_Fe3+ = 0,082.55,845 = 4,97 (g/l)

Number of grams of Fe contained in the amount of sample analyzed:

m = M_Fe.N_Fe3+ = 55,845.0,089 = 4,97 (g)

IV. Discussion

The comment

The experiment is true from the hypothesis: the solution changed from cyan green to purple blue. The titration results of the following two times were more accurate than the first. (the first time was 17,7ml and the last two times was 17,9ml). It was the error due to misread. Also, because the next time we did more carefully. For more accurate results, we needed to clean the apparatus when doing the experiment. In addition, we needed to work more seriously.

The role of substances

The process of reducing Fe³⁺ by Zn in HCl environment.

The reactions have occurred during Fe³⁺ reduction:

2Fe³⁺ + Zn 2Fe²⁺ + Zn²⁺

Zn + HCl ZnCl₂ + H­₂

The role of 1:1 HCl:

1:1 HCl reacted with Zn to create new born H₂, acting as a reducing agent. HCl took excess to prevent Fe²⁺ from being reduced to Fe. HCl created a dark yellow FeCl₄^- complex which identified the end of the reduction process.

The end of the reduction process:

Solution was boiling, then lost yellow color and did not have a white precipitate of the Zn base salt.

Condition for the reduction to obtain Fe²⁺:

Zn and HCl used residual. In addition, we boiled it until it was nearly boiling, not completely. The filtration process quickly rinsed so that no oxidation of Fe²⁺ in the air occurred.

Titration procedure for Fe²⁺ using K₂Cr₂O₇ solution.

Titration reaction:

Cr₂O₇^2- + 6Fe²⁺ + 14H⁺ 2Cr³⁺ + 6Fe³⁺ +7H₂O

Conditions of titration reaction:

Environment: strong acid, PH = 0

Indicator: diphenylamine

The color of the solution changed from colorless ( because the [Fe(HPO₄)₂]^- complexon has no color) to light yellow ( because Cr³⁺ is yellow ) and at the end of the titration the solution turned purple blue ( because the indicator changed color ).

The role of H₂SO₄ and H₃PO₄ acids:

H₃PO₄ created a colorless [Fe(HPO₄)₂]^- complex for easy equivalence determination. H₂SO₄ created a strongly acidic medium, kept PH = 0 during the titration.

V. Conclusion

From the experiment, we knew how to determine the concentration of a solution of Fe³⁺ by K₂Cr₂O₇. The experimental procedure was to reduce Fe³⁺ to Fe²⁺ then we titrated Fe²⁺ by K₂Cr₂O₇. From there we had quite accurate experimental results. However, the titration still had the error. Therefore, we needed to do carefully. So this experiment was important that we needed to know. It helped us to have more experimental skills and experience.

20193.117436.Midterm1.NVA-22

Determination of EDTA concentration by ZnSO4 solution

Introduction

Titration is an analytical technique that allows the quantitative determination of a dissolved substance (analyst) in a sample. This is a method widely used in the field of chemistry. Definition of some the substances were used in the experiment. ‘EDTA is a multi – acid (hex aprotic), acid atoms are the atoms lost during complexing complexon(III) symbol is Na₂H₂Y, this is a substance of sufficient purity for diluting standard solutions. In the experiment we also use an indicator called ETOO, this is a pH-dependent color indicator. The aim of the current experiment was to determine the concentration of EDTA by ZnSO4 solution’.

The obtained parameter is volume, from the volume, we determine the concentration. The basis of the method:

Titration equations:

Na₂H₂Y = 2Na⁺ + H₂Y^2-

Zn²⁺ + H₂Y^2- ⇋ ZnY^2- + 2H⁺

Eriocromden T complexing indicator (ETOO) denoted H₃Ind. Because the color of this indicator is dependent on pH, therefore, we use buffer NH₄OH + NH₄Cl to ensure the discoloration of the indicator at pH = 8-10. The change from the red color of the ZnInd complex- to green by Hind^2- , free state indicates the end of the titration.

Experimental procedure

Materials and apparatus

Chemicals were available in the laboratory : A 0.05N ZnSO4 solution and unknown concentration of complexon solution were prepared. Besides the two main solutions above, there were also a buffer solution (NH₄Cl + NH₃) and ETOO indicator.

Used apparatus were beaker, pipette, burette and Erlenmeyer flask

The order of conducting the experiment

‘First, the apparatuses were cleaned, inside pipette and burette should be rinsed with the solution to be titrate. Then we loaded complexon (III) solution on the burette and removed all bubbles, set the burette to zero. Next, we took exactly 10ml of ZnSO4 solution into an Erlenmeyer flask. After that five to seven ml of NH₄CL+NH₃ buffer solution and some ETOO indicator were added. Then the solution was shaken until a purple pink color appeared. Next, we proceeded with the titration until the solution turned clear blue and recorded complexon (III) volume. Last, we titrated at least three times and took the average volume’.

Results

10ml of 0.05N Zn²⁺ solution was taken for titration; we obtained the volume of H₂Y^2- after three titrations respectively: 11.6ml, 11.5ml, 11.5ml. The average volume of H₂Y^2- after three titrations were 11.53ml. From the above average volume, we could determine the concentration of H₂Y^2- was 0.043(N).

Discussion

The results of the experimental process had errors compared with the theory; however, this error was still in the increments. The cause of the error was the volumetric apparatus error, the experimenters did not get the exact amount of chemicals. Objective reasons from recognizing the time of color change to determine equivalence points were not correct.

During the titration you need to maintain the pH of the medium at 8-10, to ensure discoloration of the indicator. We ensure that a conditional stable constant of complexes made up of metal ions and complexon is large, to completely react. Keep the free form of ETOO green in the signal to indicate the end of the titration

The discoloration of the solution is on the basis: when complexons are added, first it reacts with Zn²⁺, when no Zn²⁺ complexons are added to replace the Hind^- indicator. The change from purple to clear blue in solution is indicative of the end of the titration.

Conclusion

The titration method is an interesting and popular method. It helps us to determine the composition and concentration of an unknown substance very easily. The operations are quite simple and give reliable results. Along with that, the titration method also helps us to save more time than other methods. This method is very convenient for checking whether the theoretical calculation results on our paper is true or false.

20193.117436.Midterm1.NVA-23

Titration of the NaOH solution with oxalic acid

Introduction

Titration of solutions are understanded as solutions of known to exact concentration. It is used to determine the concentrations of different solutions.

Preparation was precised carefully, the titration would be the better. The aim was

prepared a few solutions from a stock standard of solutions and the concentration was confirmed by titration. ^[2] The titration allowed to determine the concentration of unknown NaOH with an error of less than 5%. ^[1]

Experimental Procedure

Chemicals

The 96% NaOH available in the laboratory. Oxalic acid crystals H₂C₂O₄.2H2O are prepared as standard solutions with a concentration of 10^-3 M. Phenolphthalein indicator was placed in a glass bottle on the laboratory rack.

Apparatus

The apparatus includes: a technical scale, an analytical balance, a cup and volumetric flask of 100ml, 250ml, a glass funnel, a rubber bulb, and three pipettes with volume values ​​of 2ml, 5ml, 10ml, three 150 ml conical flasks were rinsed for the laboratory.

Procedure

Preparation NaOH solution from 4.2 grams of solids, stir and pour into a 100 ml volumetric to flask. After that, we added to prepare 0.1 N H₂C₂O₄ solution from 0.63 grams of H₂C₂O₄.2H₂O, dilute the solution and then titration.

Figure 1: Preparation of H₂C₂O₄ solution. ^[3]

Figure 2: Preparation of NaOH solution^{. [3]}

The titration 20 ml of NaOH with 10^-3M H₂C₂O₄ added two drops of phenolphthalein indicator gently to shake. Oxalic acid solution is loaded on the burette which expelled air bubbles. The titration of prepared solution has started until color was lost then stop.

Results and Discussion

The results are shown in the following table.

Table 1: The volume H2C2O4 consumed during titration. ^[1]

From the table above, the consumption volume of oxalic acid solutions is 17 ml. Call C1, C2 and V1, V2 the concentration and volume of the solution before and after dilution is respectively. The concentration of the solution is constant, so C1.V1 = C2.V2 or N_H2C2O4.V_H2C2O4 = N_NaOH.V_NaOH. The concentration of NaOH can be determined from this formula. The concentration of NaOH was found 8.5.10^-4 M.

From the pH jump was determined the concentration of the above solution was determined through the equivalent volume of oxalic acid.

Figure 3: The pH-volume relationship of oxalic acid. ^[1]

The relationship between volume of solution and pH has been shown in figure 3.

This process shows the change in color of the solution during the titration.Where the horizontal axis represents the volume of the oxalic acid solution and the horizontal axis represents the pH value of the solution. When a base is titrated with an acid its pH value gradually decreases the color of the solution from pink to colorless. From the graphics, we can determine the equivalence point. Discoloration of there the most obvious was occurred at that time. The pH plot depending on the volume of the oxalic acid solution is a smooth descending curve. These are expressed in the following formula: pH = -log [H ⁺].

Conclusion

The method has been relatively to accurate, it saved time. The reliability increase and accuracy in order to increase, it will be further to titrate with another indicator and comparation of two methods from there. In short, the experiment helped to understand more about the titration of method and the concentration of an unknown solution was determined. Instrument preparation, a burette reading also leads to error process. It is acceptable to ignore.

20193.117436.Midterm1.NVA-24

SYNTHESIS OF ETHYL ACETATE

Introduction

The aim of this experiment is synthesis of ethyl acetate. In the laboratory, the ethyl acetate is synthesized from ethanol and acetic acid. The reaction synthesis of ester between ethanol and acetic acid happen in condition solution to have pH less 7 and high temperature. In this experiment, sulfuric acid is used catalyst for the reaction. This is equation synthesis of ester:

CH₃COOH + C₂H₅OH H₂SO₄, 120^oC CH₃COOC₂H₅ + H₂O

Ethyl acetate ester, also known as 3-hydroxybutanone or acetyl methyl carbinol, an organic compound with the formula C₄H₈O₂. It is a colorless liquid with a pleasant, buttery odor. Density, boiling point, molar mass of ethyl acetate is: 0.902 g/cm³, 77.1^oC, 88.106 g/mol. [1]

Figure 1. Structure formula of ethyl acetate

Ethyl acetate is a widely used solvent in industry. It is used primarily for paints, lacquers, varnishes [2]. Besides, it also is used for cleaning mixtures, perfumes and nail polish removers.

Experimental Procedure

All chemicals were used as received: 95% ethanol, glacial acetic acid, sulfuric acid, 2% Na₂CO₃, 50% CaCl₂, CaCl₂ dried at laboratory.

Figure 2. The modulation of ethyl acetate [3]

Method:

First, mixture include: 5,00 ml ethanol and 5,00ml sulfuric acid were added three- necked round-bottom flask. At the same time, the mixture was cooled, stirred. Next, mixture include: 10,00ml ethanol and 10,00ml glacial acetic acid were added separating funnel. After that, laboratory instruments were assembled as figure 1. Three-necked round-bottom flask was heated up to 115^oC-125^oC on hot place. The mixture (CH₃COOH +C₂H₅OH) in separating funnel was dripped gradually down; the solution dripped to the end. After that, the mixture was heated ten minutes more time.

The distilled solution was neutralized with 5% NaCO₃ to pH=6-7. The liquid layer below the separating funnel was removed completely. 50% CaCl₂ was added the ester layer to remove ethanol. After that, CaCl2 dried also was added the ester flask to remove water.

The main product was distilled into two segments: before 75^oC and from 75^oC to 78^oC.

Figure 3 . The water-repellent distillation [3]

Result

Volume of pure ethyl acetate is 13,6ml. The efficiency (%mass) of reaction was calculated from volume of ethyl acetate.

%H= 13,6 x 0,902 = 40,06%

0,348 x 88

Discussion

After finishing the experiment, the product was obtained pure ethyl acetate. The efficiency of reaction is 40,06% and this efficiency is low. This esterification is a reaction reversible. As the volume of ethanol or acetic acid increased, the efficiency increased. Temperature of distillation process is maintained stably because some other reaction may happen, then efficiency of esterification will reduce.

C₂H₅OH H₂SO_4, 170^oC C₂H₄ + H₂O

2C₂H₅OH H₂SO₄, 140^oC C₂H₅OC₂H₅ + H₂O

Conclusion

Esterification process happens with two main stages. Efficiency of reaction may change dependent on reactant.

20193.117436.Midterm1.NVA-25

Refining of Benzoic Acid

Introduction

Benzoic acid is one of the necessary materials used in chemical industry. Specifically, it is element of gum, pesticides, analgesic and artificial preservative. In standard conditions, benzoic acid is the crystal solid that is colorless, odorless. Its melting point is 122.4^oC and its boiling point is 249^oC (1). Refining means remove the unwanted ingredients from chemistries. Basing on different solubility of different matters in one solvent, impurities are removed, and crystal is collected. The aim of this experiment is confirming the theory and refining of benzoic acid.

Experiment Procedure

The experiment was proceeded with these following steps. First, Five grams benzoic acid and 210 milliliters distilled water were put into a 250-mililiter-round bottom flask. Then, this was designed in such a way that described in Picture.1 (2)

Picture.1: Utensils installation diagram of making saturated solution

Utensils were set in order from bottom to top. Allihn condenser was put on round bottom flask and both were placed above and one centimeter from hot plate. When water flowed in condenser, the solution in the flask was heated by heater. After 5 minutes, if there was still a large amount of solid, adding mount of water into the flask through the condenser. This mixture was boiled continuously until the solid was completely dissolved. Saturated solution was collected after this stage.

In order to remove impurities, the saturated solution was refined by Buchner funnel in hot environment condition. To avoid early crystallization, vacuum and quick action was required. This solution was divided into two beakers: one was proceeded in room temperature without mixing; one was stirred up and cooled out the beaker by a mixture of ice and salt to watch and record phenomena.

The mixture in both beakers was aggregated, got cool and refined by vacuum filter. The collected crystal was cleaned in Buchner funnel by a small amount of cold water. The production was in drying oven at 80^oC until dry. The refined benzoic acid was determined its melting point and weighted. The crystal was stuffed into glass capillaries and measured with a Digital Melting Point Apparatus.

Results and Discussion

The result of the refining process is that the mass of benzoic acid was 2.51 grams and melting temperature was 124^oC. This result demonstrates two things.

First, crystallization of one matter in one solution depends on external conditions, specifically in this experiment, is temperature and stirring. It is much easier for a matter to crystallize in cool environment than hot one. Mixture causes convection of solution that breaks the undercooling.

Second, the result highlight that little is known about the efficiency of refining is calculated by the recipe

. (2)

With is the mass of benzoic acid in the result and is at the beginning. However, the refined benzoic acid was incompletely pure. Because the melting point of the collection was 124^oC that was higher than the theoretical temperature being 122.4^oC. The difference between these temperatures could be explained by the error during experimental process.

Conclusion

This experiment elucidates the way to refine benzoic acid. This conclusion follows from the fact that environmental conditions can influence formation of crystals. In addition, these findings provide information about the process of making benzoic acid purer with the efficiency was 50.2%. Nevertheless, the purity of the refined sample was not 100% for the reason that the produce depended on the manipulations of the performer which were not standard.

20193.117436.Midterm1.NVA-26

“Determination the NAOH solution concentration by the HCL solution”

Introduction

The purpose of this experiment was to identify the concentration of a base solution by determining the volume of acid consumed to neutralize the base. The experiment was determined using by an acid-basic titration method. Acid-basic titration is an analytical technique that allows the determination of the volume acid solution reacts adequately with a basic solution. The experiment uses an indicator: phenolphthalein and methyl orange. In the titration, these two indicators are used to determine the timing of a solution's colour change. Consequently, we can measure the medium volume of hydrochloric acid.

Experimental Procedure [1]

2.1. Laboratory apparatus and chemicals

All chemicals used in the experiment include tested NaOH solution, HCl standard solution (0.1N). All laboratory instruments include Volumetric flask 250ml, Erlenmeyer flask, Beaker 100ml, Burette, Pipettes 10ml, wash bottle.

2.2. Performing a titration of the solution (NaOH) with an acid solution (HCl)

Preparation 0.1N HCl solution from a concentrated HCl solution

(A thick acid HCl d=1.9 g/ml; C% = 38%; MHCl= 36.5 g/mol)

Firstly, a volume of VHCl = 4.2 mL was obtained by diluting 500 ml of HCl solution with concentration = 0.1N. Take a measuring cylinder of about 4,2 mL of concentrated hydrochloric acid from the bottle (work in a fume hood). Fill a beaker containing both of 500 mL of distilled water. Then stir well with a glass rod for a homogeneous solution.

The second, performing a titration of the solution (NaOH) with an acid solution (HCl). The tool was been washed. Load prepared HCl solution into the burette, expel air bubbles, correct zero marks before standardization degree. The tested NaOH solution (10 ml) was taken from the test sample and was added 7-8 drops of phenolphthalein (pink solution) to the conical flask. The reaction stopped when the colour changed from pink to colourless. If the indicator was used methyl orange, the above would repeat operations. In this test, add 1-2 drops of methyl orange (yellow solution) to the conical flask. The reaction is stopped when the solution changes colour from yellow to orange. Each reaction was carried out experimentally at least three times (error0.2%). The result was read the lower surface of the solution touching the burette mark. During the titration, we need to pay attention to the colour of the solution in the conical flask to avoid equivalent point over-titration.

Analysing data

The determination medium volume HCl using the phenolphthalein indicator. Then, the equivalent concentration of NaOH (NNaOH) and mass concentration of NaOH (Cg/l (NAOH)) were determined

When the methyl orange indicator was used. The data was NNaOH(N)=0.07 mL and CNaOH(g/L) = 2,7998 g/L

Results

After the experiment, the results were obtained in table 1 and graph 1

Table 1. The volume of HCl acid consumed in the titration

Discussion

The titrate reaction between NaOH solution and HCl solution: this is the titration reaction between a strong acid and a strong base. The mechanism the colour change of methyl orange: when the solution has a PH greater than 4.4, methyl orange is in a basalt medium and has a yellow colour. During the reaction, the PH abruptly changes after passing the equivalence point. At this time, the PH is less than 3.1, Methyl Orange is acidic and changes colour to orange. With phenolphthalein, when the solution has a pH greater than 8, the solution is pink. When the reaction occurs, the pH gradually decreases so the colour of the solution fades. After the equivalence point, the NaOH disappears when the solution becomes colourless.

Therefore, whichever indicator was used, the average volume of the hydrochloric acid solution after the three experiments was approximately the same. The results obtained different depending on the preparation of the test device and titrated the operation. Before the titration, one can predict that the average volume of acid consumed using the two indicators is approximately the same. Our job is to do experiments to see if this prediction is correct.

In fact, possible sources of error in this experiment include the reading result, device error and the colour change at the equivalence point. This affects the volume of the HCl solution. A wrong volume value would impact the concentration of a solution.

Conclusion

In sum up, how did the experiment help us to understand the titration method? Instrument preparation, reading burette and experimental manipulation should a great will influence on titration results. And how to choose the right indicator of the reaction to recognizing discolouration. Determining the average volume of hydrochloric acid. Then the determining the equivalent concentration of the base solution (NaOH)

20193.117436.Midterm1.NVA-27

Chemicals in the lab are pure substances without impurities. In order to get pure chemicals, it is necessary to collect, separate, filter and recrystallize; in which stages recrystallization to obtain pure products is very important. For that reason, I wrote this report to introduce to everyone about the recrystallization method as well as its notes and how to carry out this method, especially with benzoic acid.

ORGANIC CHEMISTRY REPORT: REFINING OF BENZOIC ACID BY THE RECRYSTALLIZATION METHOD

Introduction

Benzoic acid

Benzoic acid is the simplest aromatic carboxylic acid, which has a white (or colorless) solid with the formula C₆H₅CO₂H.

Physically [1]: Benzoic acid is a white (or colorless) solid, no scent; slightly soluble in cold water and dissolve a lot of hot water. The melting temperature is T_t = 122, and the melting point is T_m = 249 ℃

Chemistry [1]: Benzoic acid has properties of weak acidity and benzene (S_E)

Benzoic acid occurs naturally in many plants and serves as an intermediate in the biosynthesis of many secondary metabolites. Salts of benzoic acid are used as food preservatives. Benzoic acid is an important precursor for the industrial synthesis of many other organic substances so we need to refine it to get the purest usable crystals.

Recrystallization method [2]

Recrystallization method is one of the most basic method to crystallize. We could base on the principle that different substances have different solubility in the same solvent so we can eliminate impurities and obtain pure crystals.

The selection principles solvent: 5 rules

Do not react with compounds to be crystallized

The compound to be crystallized must be well soluble in hot solvent and insoluble in cold solvent

Impurities must be insoluble in solvents

The solvent must have a boiling point less than the melting point of the substance to be refined

The solvent must have a low boiling point for easy removal after purification

Because of benzoic acid properties and the requirement of a solvent for the recrystallization process, we could choose water as the solvent for the recrystallization of benzoic acid.

Procedure [3]

Chemical: Chemicals used in the experiment are available in the laboratory. There was impure benzoic acid: 5,0 g; distilled water: 210 ml and activated carbon: 0,1 g

Laboratory instruments

Produce

First, we prepared adequate instruments and laboratory chemicals, observed the test instrument installation diagram and made chemical phase preparation.

Secondly, 5-gram-benzoic acid and 210-ml-water were poured into the reaction equipment that described in the picture. The condenser was installed as long as it makes a straight with the main reaction equipment which was set up at the bottom of the condenser. The distance between the electric stove and the bottom of the flask must be 1cm. A flow of water was passed through the condenser, turned on the electric hot plate then a saturation liquid of benzoic acid was created. During heat, we needed observe the dissolving process of benzoic acid and draw any comments. After 5 minutes, if the number of solids remains much, add solvent to dissolve the solids by through the condenser (note: turn off the hot plate before you added)

If the solvent in main reaction equipment was impure then erased it if needed.

In order to erase the unwanted color, turn of the electric stove for a moment, then unbolt the condenser, after that 0.1 gram of finely ground, the activated carbon would be added to the main reactor. Condenser would be reinstalled and boiling process was continuing.

The color in the liquid was erased then the liquid needed to be filtered to separate impurities. The filter process would be conducted with Buckner funnel. The gained liquid was hot, crystal-clear and with no color. (We should boost the speed up in filter process to avoid crystallization process)

Obtained liquid is divided into two cups. The first one would be covered at room temperature. The other one was cooled and stirred. Observed the crystallization process in both cups and made a comparison and conclusion

Combining two cups and cooling outside, crystal would be obtained by the filtering process with the vacuum filtration equipment. An amount of cold water would be used to make the crystal cleaner in the Buckner funnel.

Drying the product at 80 °C until absolutely dry then used the final product to detect the melting point. Weighted the product and found the efficiency.

Results

After the experiment, benzoic acid crystals initially switched from opaque white to white (the color of pure benzoic acid crystals) and crystal mass obtained was 2,36(g). It was 47.2% from the total amount of raw crystals had been used.

The melting point was 122℃ for the first time and 123℃ for the second. It was necessary to determine the melting point of the obtained crystals to know whether the crystal was pure or not

Discussion

During the experiment, we have seen the properties of benzoic acid and the principle of the recrystallization method.

The first amount of acid is pure, it has 47,2% acid and 52,8% impurities. This was a relatively good result.

The determined melting temperature is also relatively accurate, the first being 122℃ and the second being 123℃ (near form 122,4℃). Pure Benzoic acid is crystalline solid, with no color and melting point is 122,4°C. Looking at the results obtained, we could see the deviation. The reason for this deviation may be due to errors in the lab and the obtained benzoic acid is still incomplete.

Conclusion

The experiments were taken place right under the guidance and results within acceptable error.

Through the exercise we have seen the properties and applications of benzoic acid, from which, we could understand the purpose of obtaining pure benzoic acid as well as the recrystallization method.

Moreover, we also need to pay more attention to the chemicals and carry out also the error is minimal

20193.117436.Midterm1.NVA-28

Experiment 1: Chemical Equilibrium

Introduction

Chemical equilibrium is a state of a reversible reaction, the number of molecular product substances is formed from the initial substances as the number of molecular product substances react with each other to create the initial substances in a time unit.

The objective of this experiment is to determine the equilibrium constant of the homogenious reaction in the solution:

2Fe³⁺ + 3I^- = 2Fe²⁺ + I₃^-

by the method of measuring the concentration directly at equilibrium state.

In the water solution, Iron (III) can oxidize iodine (I^-) to create Iron (II) and Iodine (I₂). In this redox reaction, if it proceeds at PH ≈ 0 [so that iron (III) can exist in ion Fe³⁺] and the redudant I^- (so that I₂ which is created can exist in ion I_3- and avoid losing Iodine in the sublimation), we have a fundemental equilibrium equation:

2Fe³⁺ + 3I^- = 2Fe²⁺ + I₃^- (1)

The equilibrium constant of the reaction (1) is determined:

K = [Fe²⁺]_cb ²[I₃^-]_cb / [Fe³⁺]_cb ²[I^-]_cb (2)

By the method of the titration of the volumn, we can determine the concentration of constituents at equilibrium state.

Experimental Procedure

Materials

200ml erlenmeyer flask, 150ml ground stopper

10 ml pipette

25 ml burette

Water bath

Distilled water dispenser

All of them are availble in the laboratory.

Experiment Procedure

Redetermine the concentration of ~ 0,03M FeCl₃

We use the pipette to get 5ml of ~ 0,03M FeCl₃ into an erlenmeyer flask 150ml. We add 5ml of distilled water and ~ 0,5 KI, then shake them slightly to dissolve. After 15 minutes, we make the titration of I₂ produced by Na₂S₂O₃ solution 0,01M from buret elenmeyer flask to canary-coloured. We add 3 drops of starch into the erlenmeyer flask.

From using 0,01M Na₂S₂O₃ solution, we can determine the exact concentration of FeCl₃ and write down the determined value C⁰_Fe³⁺.

Redetermine the concentration of I_3- after its reaction

We put FeCl₃ solution (determined the concentration above) and 0,03M KI into 4 150ml erlenmeyer flasks (mark them in order 1, 2, 3 and 4) following the ratio of the volumn in this table:

Pouring the solution in the flask 1 and 2, we regard it as the starting reaction point of compounds (1+2).

After 30 minutes from its reaction time, we use the pipette to get 5ml of the compounds (1+2) to put them into the elenmeyer flask with 40ml of distilled water which was already refrigerated (about 10^ºC) and then we make the titration of Iodine produced (from I₃^-) by 0,01M Na₂S₂O₃ solution with the starch above.

We repeat the process of titration above, about 15 minutes per time, until the volumn of 0,01M Na₂S₂O₃ solution uses the equally adjacent titration twice.

Results and Discussion

We redetermine the concentration of FeCl₃ solution:

0,01M the volumn of Na₂S₂O₃ uses the titration 5ml FeCl₃: V_Na2S2O3 = 15 (ml)

C⁰_Fe³⁺ = 0,03 (M)

We determine [I₃^-] at equilibrium: [KI]₀ = 0,03 (M)

Overall, from 2 tables, we can see that at the same time, the volumn of Na₂S₂O₃ solution in each table has an insignificant difference, the flask (1+2) has 6,6 ml and 6,4 ml whereas the other one (3+4) has 6,4 ml and 6,3 ml which corresponds to 30 minutes and 45 minutes, respectively. In addition to that, at the same time in the flask (1+2) and (3+4), we have the volumn of Na₂S₂O₃ solution at 60 minutes, 6,3 ml and 6,2ml; however, at 75 minutes,the volumn of that one is not defined.

Handling the figures in the flask (1+2)

[I₃^-] = 3,22.10^-3 (M)

[Fe²⁺] = 2.[I₃^-] = 6,44.10^-3 (M)

[Fe³⁺] = 6,44.10^-3 = 8,56.10^-3 (M)

[I^-] = – 3.3,22.10^-3 = 5,34.10^-3 (M)

K= [Fe²⁺]_cb ²[I₃^-]_cb / [Fe³⁺]_cb ²[I^-]_cb = = 11,97.10³

Handling the figures in the flask (3+4)

[I₃^-] = = 3,15.10^-3 (M)

[Fe²⁺] = 2.[I₃^-] = 6,3.10^-3 (M)

[Fe³⁺] = 6,3.10^-3 = 0,0117 (M)

[I^-] = 3.3,15.10^-3 = 2,55.10^-3 (M)

K= [Fe²⁺]_cb ²[I₃^-]_cb / [Fe³⁺]_cb ²[I^-]_cb = 55,08.10³

Conclusion

The experiment of measuring the concentration of the substances directly at equilibrium state helps us determine the equilibrium constant of the reaction:

2Fe³⁺ + 3I^- = 2Fe²⁺ + I₃^- . The cause of the error of the results is the early process of the titration or late than theory, the error of meters,….