Experiment 3 - Stoichiometry Reaction

Experiment 3

Title: Stoichiometry Reaction

Objectives:

To decompose sodium hydrogen carbonate (sodium bicarbonate) by heating, and to accurately measure the degree of completion of the reaction by analyzing the solid sodium carbonate product.

Introduction:

Stoichiometry is the calculation of quantitative (measurable) relationships of the reactants and products in a balanced chemical reaction. It can be used to calculate quantities such as the amount of products that can be produced with the given reactants and percent yield.

Stoichiometry calculations are based on the fact that atoms are conserved. They cannot be destroyed or created. Numbers and types of atoms before and after the reactions are always the same. This is the basic law of nature. From the atomic and molecular point of view, the stoichiometry in a chemical reaction is very simple. However, atoms of different elements and molecules of different substances have different weights.

We must be able to relate the amount of heat evolved in a laboratory scale reaction to that involved when two molecules react. The scaling factor used to relate readily useable quantities to the molecular scale is called the mole.

1 mole = 6.023 x 1023 molecules = Avogadro’s number of molecules

In this experiment, several reactions will be performed and physical measurements will be made that subsequently relate to the molecular scale. The thermal decomposition reaction proceeds above 270 °C according to the following equation:

2NaHCO3 (s) Na2CO3 (s) + CO2 (g) + H2O (g)

A known mass of NaHCO3 will be decomposed to Na2CO3 in this experiment. The weight of the Na2CO3 allowing a comparison with the percentage conversion of NaHCO3 into Na2CO3 can be calculated. Solution of Na2CO3 product will be titrated against HCL with known concentration. A second value of mass of Na2CO3 will be formed in the thermal decomposition.

Apparatus and Materials:

Sodium bicarbonate, hydrochloric acid (0.05 M), Thymol blue indicator, distilled water, electrical balance (4 decimal places), test tubes, test tube rack, Bunsen burner, volumetric flask (250 cm3), glass funnel, burette (50 cm3), retort stand, conical flask (250 cm3),

pipette (25 cm3).

Procedures:

Part 1: Thermal Decomposition of NaHCO3

4-decimal place balance was used to record the mass of a clean, dry test tube. 2.0g to 2.3g of NaHCO3 was added to the test tube.
With the aid of test tube holder, the tube and the contents were heated gently over a hot flame for not longer than 5 minutes.
The tube and the contents were allowed to cool to room temperature in a test tube rack and they were reweighed on the same balance.
Those masses were recorded and the amount of solid product was calculated.

(Na2CO3 + any unreacted NaHCO3)

Part 2: Titration of Na2CO3 with Hydrochloric Acid

A few cm3 of distilled water was added to the test tube. It was heated gently and shaken to dissolve the Na2CO3.
The solution was transferred to a 250 cm3 volumetric flask using a funnel. Few more cm3 of distilled water was added to the test tube and this was added to the volumetric flask.
The washing was repeated 2 or 3 times, and the funnel was washed with distilled water too to ensure all the Na2CO3 transferred to the volumetric flask.
The solution was made up to 250 cm3 mark. The volumetric flask was capped and inverted 6 or 7 times to homogenize the solution.
Burette was rinsed out with a little of the standardized 0.05 M HCl solution supplied. The burette was filled with the standardized HCl and was clamped vertically.
Three 250 cm3 conical flasks were rinsed out with distilled water, and 25.00 cm3 aliquots of Na2CO3 solution was pipetted into each.
2 or 3 drops of thymol blue indicator were added to each flask and the Na2CO3 solution was titrated with the HCl until the end-point is reached (grey-green colour). HCl was added drop-wise towards the end of the titration.

Results:

Table 1: Weights of test tube and NaHCO3.

Table 2: Weights of solid products (Na2CO3 + unreacted NaHCO3) after heating and

cooling the test tube contains NaHCO3.

Table 3: Initial and final burette readings, volume of HCL solution to neutralize the Na2CO3

solution.

Analysis and calculation:

From Table 1:

Weights of NaHCO3 = (weights of empty test tube + NaHCO3) – (weights of empty test tube)

= 32.5867g – 30.4367g

= 2.1500g

From Table 2:

Amount of solid products = (weights of empty test tube ...

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Results:

Table 1: Weights of test tube and NaHCO3.

Table 2: Weights of solid products (Na2CO3 + unreacted NaHCO3) after heating and

cooling the test tube contains NaHCO3.

Table 3: Initial and final burette readings, volume of HCL solution to neutralize the Na2CO3

solution.

Analysis and calculation:

From Table 1:

Weights of NaHCO3 = (weights of empty test tube + NaHCO3) – (weights of empty test tube)

= 32.5867g – 30.4367g

= 2.1500g

From Table 2:

Amount of solid products = (weights of empty test tube + solid products) – (weights of

empty test tube)

= 31.8689g – 30.4367g

= 1.4322g

From Table 3:

Average volume of 0.05M HCl solution required to neutralize Na2CO3 solution during titration

= [(26.8 + 26.7 + 26.5) / 3] cm3

= 26.67 cm3

The degree of completion of Na2CO3 decomposition is calculated shown as below:

Na2CO3 + HCl NaHCO3 + NaCl

No. of mole of HCl = concentration x volume

= 0.0500 M x 0.02667L

= 0.00133 mol

No. of mole of Na2CO3 = No. of mole of HCl

= 0.00133 mol

Mass of Na2CO3 = no. of mole of Na2CO3 x molar mass of Na2CO3

= 0.00133 mol x [23(2) + 12 + 16(3)]

= 0.00133 mol x 106

= 0.14098g

Degree of completion of Na2CO3 decomposition

= 93.44%

Discussion:

In this experiment, electrical balance with 4-decimal place was used. This is to increase the accuracy of the results during weighing process, so that we can gain the weights of solid products which contains Na2CO3 and unreacted NaHCO3 more accurately. Table 1 shows the weights of empty test tube and NaHCO3. After gaining the weights of NaHCO3, the test tube with sodium hydrogen carbonate, NaHCO3 was heated not more than 5 minutes to decompose it in order to get the solid sodium carbonate product. The distilled water was added in the test tube contains Na2CO3. The test tube was heated gently and was shaken at the same time to ensure all the solid can be dissolved completely.

In the titration process, 0.05M of HCl was added drop-wise into Na2CO3 solution in order to gain the end-point of the solution more accurately. The endpoint is defined by the choice of indicator as the point at which the colour changes. In this experiment, the end-point was reached when the solution turned grey-green colour. To increase the accuracy in observing the change of colour, we can increase the amount of thymol blue indicator added in the solution, so that the colour changes can be more obvious. The grey-green colour was gained just as the blue colour of the solution disappears. The turning colour of solution into yellowish should be avoided because the yellow colour indicates the solution is becoming acidic already and the titration should be repeated until the correct green-grey colour is obtained. In this experiment, the balance equation of the titration is as follows:

Na2CO3(aq) + HCl(aq) NaHCO3(aq) + NaCl(aq)

The degree of completion of Na2CO3 decomposition is calculated, which is 93.44%. This shows that the decomposition is not completed. The incomplete of decomposition is due to the time taken in heating the test tube is not enough to decompose all the NaHCO3 into Na2CO3.

Questions:

Use of indicator in the reaction.

An indicator is any substance in solution that changes its colour as it reacts with either an acid or a base. Selecting the proper indicator is important because each indicator changes its colour over a particular range of pH values. Indicators are either weak acids or weak bases. We choose indicator depends on several factors. One of them is the pH value at equivalent point. The pH value can be vary depend on the titrated substance and titrant used.

Different indicator will changes colour at different pH and so are for different purpose. A phenolphthalein indicator is used when it is a weak acid-strong base titration; thymol blue indicator in strong acid-strong base reactions; and methyl orange indicator for strong acid-weak base reactions. In this experiment, thymol blue indicator is used. It is a reddish-brown crystalline powder that is used as an pH indicator. It transitions from red to yellow at pH 1.2-2.8 and from yellow to blue at pH 8.0-9.6. It appears as grey-green colour when the solution is neutral.

Generally, the faster the pH changes in the range where the indicator changes colour, the sharper the endpoint of the titration will be and the more different indicators will be suitable for the titration. Knowing which pH ranges cause specific indicators to appear which colours allows us to use a set of one or more indicators to act as a crude pH meter to estimate the pH of a solution and the pH can then be used for further calculations. Besides estimate the pH value, indicator can also used to detect the equivalence or neutral point in the titration. The point in the titration at which the colour changes is known as the “end point”. Depending on how quickly the colour changes, the endpoint can occur almost instantaneously or be quite wide. The equivalence point is a single point defined by the reaction stoichiometry as the point at which the base (or acid) added exactly neutralizes the acid (or base) being titrated.

Notice that the titration is being stopped at the thymol blue end-point around pH 9.6 to 8.0. The equation for this reaction is: Na2CO3 + HCl NaHCO3 + NaCl

If more HCl is added to this solution at the thymol blue end-point it will react, what will be the products of this new reaction?

Sodium carbonate, Na2CO3 will react with hydrochloric acid, HCl in two stages. In the first stage, Na2CO3 react with HCl to form sodium bicarbonate, NaHCO3 and sodium chloride, NaCl. The balance equation is shown as below:

Na2CO3 + HCl NaHCO3 + NaCl

In second stage, more HCl is added. Sodium bicarbonate, NaHCO3 react with HCl to form sodium chloride, water and carbon dioxide. The balance equation is shown as below:

NaHCO3 + HCl NaCl + H2O + CO2

Hence, the product of this new reaction is sodium chloride, NaCl.

What do you understand by the expression “H2CO3 is a diprotic acid”?

A diprotic acid is an acid that contain within its molecular structure of two hydrogen atoms, which are capable of dissociating in water. H2CO3 is a diprotic acid, because it contains two hydrogen atoms and it dissociate in water to form 2 moles of hydrogen ions, H+.

H2CO3 H+ + HCO3-

HCO3- H+ + CO32-

Two methods are used in Part1 and 2 of this experiment to determine the amount of Na2CO3 formed from the thermal decomposition of NaHCO3. What are they?

In Part 1, thermal decomposition method is used. It involved the heating process of the NaHCO3 in test tube over the Bunsen burner followed by cooling process. Decomposing of NaHCO3 can be readily achieved through heating. The solid product of sodium carbonate, Na2CO3 is obtained and its amount can be determined through weighing. Weighing process is important to know the accurate weight of the Na2CO3 formed in the reaction. Analytical balance with 4 decimal places is used to weight the Na2CO3 in order to get the more accurate result.

In Part 2, titration method is used. Na2CO3 is titrated with HCl solution in the burette until the end-point is reached. The Na2CO3 reacts with HCl to form sodium bicarbonate, which is the original reactant used in the beginning of this experiment. The average volume of Na2CO3 is calculated from the 3 times of titration in order to calculate the concentration of Na2CO3.

Why is it important to transfer all of the Na2CO3 from the test tube to the volumetric flask (in the titration)?

Transferring all the Na2CO3 from the test tube to the volumetric flask in titration is considered as a precaution step as it will affect the accuracy of the readings obtained. When transferring the Na2CO3 from the test tube, distilled water is added to rinse the test tube for a few times to ensure there are no any remaining of sodium carbonate stick at the inner side of the test tube. If this process is not carry out, this will affect the accuracy of the concentration sodium carbonate used in the titration process.

What is Hess’s Law?

Hess’s Law is a relationship from physical chemistry named Germain Hess. Hess’s Law states that the heat evolved or absorbed in a chemical process is the same whether the process takes place in one or in several steps. This is also known as the “law of constant heat summation”. It can be used to predict energy changes that are not easily measured.

Hess's law allows the enthalpy (ΔH) for a reaction to be calculated even when it cannot be measured directly. This is accomplished by performing arithmetic operations on chemical equations and known ΔH values. When an equation is multiplied by a constant, its ΔH must be multiplied by the same number as well.

Hess’s Law also stated that the enthalpy changes are additive. Thus, the ΔH for a single reaction can be calculated from the difference between the heat of formation of the products minus the heat of formation of the reactants.

ΔH° = Σ(ΔHf°products) - Σ(ΔHf°reactants)

Addition of chemical equations can lead to a net equation. If enthalpy change is included for each equation and is added, the result will be the enthalpy change for the net equation. If the net enthalpy change is negative (ΔHnet < 0), the reaction is an exothermic reaction. On the other hand, a positive ΔH value shows an endothermic reaction.

Conclusion:

The decomposition of sodium hydrogen carbonate can be readily accomplished by heating NaHCO3 in a test tube over Bunsen burner. The degree of completion of Na2CO3 decomposition was calculated, which is 93.44%.

Reference: