Determination of the activation energy for the reaction of bromide and bromate ions in acid solution.

Objectives

1. To understand the chemistry of activation energy.

2. To determine the activation energy for the reaction of bromide and bromate ions in acid solution.

3. To understand the chemistry of rate of reaction.

Introduction

Activation energy is defined as the minimum energy that must be overcome in order for a chemical reaction to occur. It is usually denoted as E_a, and given in unit of kJ/mol. For a chemical reaction to proceed at a reasonably rate, there should exist an appreciable number of molecules with energy equal to or greater than the activation energy. The concept of activation energy also can be related to the collision theory. According to the collision theory, in order for a reaction to occur, the reactant particles must collide. However, not all collision brings about reaction. Only a certain fraction of the total collisions cause chemical change, which are called successful collision. The successful collisions have sufficient energy to overcome the activation energy barrier at the moment of impact to break the existing bonds and form new bonds, resulting in the production of products.

The activation energy of a reaction can be measured by using Arrhenius equation,

k = A

where k is the rate constant of a reaction with absolute temperature, T, is the energy of activation of the reaction and R is the gas constant. The pre-exponential term, A is the property of particular reaction related to the collision frequency of the reactive species and thus is temperature dependent. However, according to the equation, the dependence of k on temperature is dominated by the strong exponential term, so the dependence of A on temperature is usually ignored as a first approximation. By taking logarithms of both sides,

lo k = - /2.303RT + lo A

= - /2.303RT + constant

So, when a reaction has a rate constant that obeys Arrhenius equation, a plot of lo k versus 1/T gives a straight line. The gradient of the straight line is - /2.303R while the interception of the straight line on the y-axis of the graph can be used to determine the values of lo A.

Now, the rate of reaction is higher when the time taken for a fixed amount of reaction to complete is shorter. This makes the time taken, t to complete a fixed amount of reaction is inversely proportional to the rate constant, k.

t α 1/k

or t = constant/k

By taking logarithms of both sides,

lo t = - lo k + constant

= /2.303RT + constant

A plot of lo t versus 1/T gives a straight line as well and the slope of the graph is /2.303R. Thus, if t is measured at several temperatures then the energy of activation can be found.

In this experiment, the above method is applied to the reaction of bromide and bromate ions in an acid solution which occurs slowly at room temperature.

KBr + 5 KBr + 3 3 + 3 B + 3

or Br + 5 B + 6 3 B + 3

The time required for a fixed amount of the reaction to be completed, t is found by adding a fixed amount of phenol and some methyl red indicator to the reaction mixture. The bromine produced in the first reaction reacts very rapidly with the phenol to form tribromophenol.

+ 3 B OH + 3 HBr

When all the phenol has reacted, the bromine continuously produced in the first reaction will then react with the methyl red indicator and bleaches its colour.

Methyl red + colourless compound

Apparatus : 1 dm^-3 beaker, 3 100cm³ beakers, 2 boiling tubes, 1 5cm³ pipette, 1 10-cm³ pipette, thermometer (0 - 110°C), stopwatch

Material : 0.01 mole dm^-3 aqueous phenol solution, bromide/bromate solution (0.0833 mole dm^-3 potassium bromide and 0.0167 mole dm^-3 potassium bromate, equivalent to 0.05 mole dm^-3 bromine), 0.3 mole dm^-3 sulphuric acid, methyl red indicator.

Procedures

1. 10 cm³ of phenol solution and 10 cm³ of bromide/bromate solution were pipette into one boiling tube.

2. Four drops of methyl red indicator were added to the mixture.

3. 5 cm³ of sulphuric acid was pipette into another boiling tube.

4. The two boiling tubes were immersed in the water bath of (75 ± 1) °C.

5. The contents of the two tubes were mixed by pouring rapidly from one tube to the other twice and the stopwatch was started at the same time.

6. The boiling tube containing the reaction mixture was kept immersed in the water.

7. The time required for the red colour of the methyl red indicator to disappear was determined.

8. The whole experiment was repeated at 65, 55, 45, 35, 25 and 15 °C.

9. Ice was used to achieve the lowest temperature.

Results & Calculations

(Assume R = 8.314 J K^-1 mol^-1)

From Graph 1,

Slope of the graph = E_a / 2.303R

E_a / 2.303R = (2.94 – 2.20) / (3.35×10^-3 – 3.10×10^-3)

E_a /2.303R = 0.74 / (2.50×10^-3)

E_a = 2960*2.303R

E_a = 2960*2.303*8.314

E_a = 56676 J/ mol

Energy of activation, E_a = 56676 J /mol

Discussion

The reaction between bromide and bromate ions in acid solution is a slow reaction at room temperature. This may be due to the high activation energy of the reaction, which required 56.676 kJ of energy in order for a reaction to proceed favorably. The high activation energy is one of the factors that reduce the effectiveness of the collisions in bringing about reaction because according to collision theory, it is not sufficient that molecules simply collide can brings about reaction to occur. Colliding molecules without energy equal to or higher than activation energy will not cause any reaction to happen. There are also other factors that can reduce the effectiveness of collisions of molecules such as the order of colliding molecules which are not included in the objectives of this experiment.

According to the data obtained in Table 1, we can observed that the higher the temperature of the reactant species, the shorter the time taken for the disappearance of red colour of methyl red indicator. The shorter time taken indicates that the reaction is faster and higher rate of reaction. Although the activation energy for the reaction to occur remains the same at all the temperature and unaffected, but the rate of reaction increased as the temperature increased. This means that the rate of reaction is dependence on the temperature of the reactant species and this is indeed the case. This has also been proven in Arrhenius equation where rate constant, lo k is proportionally to 1/T with the slope of the graph, - /R and a constant of lo A.

lo k = - /2.303RT + lo A

The higher temperature caused the value of the - /2.303RT closer to the value of 0. The constant lo A will then minus off the value of /2.303RT and results in a larger value. The larger value will caused the value of the rate constant, k to become larger as well. Larger value of k will then results in faster reaction. Hence, it can be said that the higher the temperature of reactants, the reaction will proceed faster with the higher rate of reaction.

Besides, the rate of reaction roughly doubles for every 10 °C rise in temperature. This is because increase in temperature increase the kinetic energy possessed by the reacting molecules, hence increase the speed of the reacting molecules. With the higher speed of moving, the reacting molecules will now collide more frequently and increase the amount of successful collision. As the kinetic energy of the molecules increase, the molecules can therefore easily overcome the activation energy barrier during the collision and bring about reaction to occur. Hence, the higher the temperature, the higher the kinetic energy possessed by the molecules, the faster the molecules move, the higher the frequency of collision, the easier the molecules to overcome the activation energy barrier at the moment of impact, the higher the amount of successful collision, the faster the reaction and therefore the higher the rate of reaction.

Fixed amount of phenol and methyl red indicator were added to the mixture contents every time the experiment repeated. This is because phenol can provides an intermediate state before the bromine molecules produced in the reaction between bromate and bromide ions in acid solution able to bleach the methyl red. In other words, phenol is used to observe the time taken for the bromine molecules to react completely with phenol and then bleach the colour of the methyl red indicator. The methyl red indicator was added to provide a distinct colour to ease the observation. This is because when the sulphuric acid was poured to the bromate/bromide ions solution, the methyl red indicator turns to pink colour. The bleaching effect from the bromine molecules caused the methyl red indicator to turn colourless after all the phenol is used up and this provides a clear difference of colour in order to stop the stopwatch. The amount of phenol added must be fixed in every repeated experiment in order to compare the time taken for the disappearance of colour of methyl red.

Phenol is chosen to be used to provide an intermediate state because phenol can react with bromine molecules very fast to produce tribromophenol and hydrogen bromide before the bromine molecules react with the methyl red.

This is very important because if other substances are chosen to replace phenol, the substance chosen might not react with the bromine so fast and caused the bromine to react straight away with the methyl red, bleaches its colour without going through the intermediate phase. The –OH group in the phenol has the effect of making the benzene ring much more reactive than it would be with a 2, 4-directing effect. This means that the incoming groups will normally go into the 2-position or the 4-position, but hardly go into the 3-position because of the 3-isomer is produced too slowly. The bromine molecules undergo substitution reaction in this reaction by substituting three hydrogen atoms from the benzene ring with three bromine atoms to maintain the aromaticity of the ring in phenol.

The reaction between bromate and bromide ions in acid solution is a redox reaction.

Br + 5 B + 6 3 B + 3

The potassium and sulphate ions act as spectator ions in this experiment and they are not involved in any of the redox reaction or changing of their state. The bromide ions undergo oxidation by donating one of its electrons to the bromine atom in the bromate ions. The bromine atom in the bromate ions then undergoes reduction by receiving electron from the bromide ions. The hydrogen ions and oxygen atom in bromate ions does not involve in increase or decrease in oxidation number but they were involved in changing the state from the ions in aqueous solution to the water molecule in liquid state.

On the other hand, there are some precautions steps that we need to take into considerations during the experiment. Care must be taken when using phenol and sulphuric acid as they can cause burn to the skin. Gloves and protective cloth must be worn at all times in the laboratory. The boiling tubes containing the mixture should be kept immersed in the water to ensure that the mixture remains its temperature while the reaction proceeding.

Determination of the enthalpy of reaction of a monobasic acid with sodium hydroxide

Objectives

1. To understand the enthalpy chemistry.

2. To determine the calorimeter constant.

3. To determine the enthalpy of reaction of acid-base reactions.

4. To study the exothermic reactions.

Introduction

Enthalpy is a measure of total energy of a thermodynamic system. It includes the internal energy of the system and the product of its volume multiplied by the pressure exerted on it by its surrounding.

H = U + p*V

H = Enthalpy p = Pressure

U = Internal Energy V = Volume

The enthalpy is normally measure with S.I. unit of Joule, J, although other units still in use such as Calorie, C and calorie, c. However, we often measure the change in enthalpy, ΔH instead of measuring the enthalpy, H because the total enthalpy cannot be measured directly.

Enthalpy change is defined by the following equation:

ΔH = H_Final – H_Initial

ΔH = Enthalpy change

H_Final = Final enthalpy of the system. In a chemical reaction, H_Final is the enthalpy of the products.

H_Initial = Initial enthalpy of the system. In a chemical reaction, H_Initial is the enthalpy of the reactants.

Enthalpy change is in positive values in endothermic reaction and negative values in exothermic reaction. Exothermic reactions involve energy released as heat into its surroundings, causing the temperature of the surroundings to rise while endothermic reactions involve energy acquired from its surroundings as heat, causing the surrounding temperature to drop. Change in enthalpy that occurs as a result of a chemical reaction is numerically equal to the heat of reaction under constant (atmospheric) pressure conditions (ΔH = q). The heat of reaction is conveniently measured adiabatically in a Dewar calorimeter by the rise or fall in temperature of the products produced by the reaction in solution. Dewar flask is used because it is designed to preserve heat and minimize heat loss to the surrounding. In addition, isolated system is also required in this experiment to obtain an accurate data. Since every Dewar flask has different calorimeter constant because of different substances used, so the calorimeter constant ( ) must first be determined, that is the quantity of heat required to increase the temperature of the calorimeter and its content by 1 °C.

The constant is measured by supplying the calorimeter and contents with a definite known quantity of heat. This can be done electrically or by adding a known amount of concentrated sulphuric acid.

Results & Calculation

Part 1 Calorimeter Constant

 

 

 

 

 

 

 

 

 

 

 

 

 

 

By using the values calculated, 0.384M in Graph 2, we can estimate the amount of heat liberated which is 2.775 kJ. Since the heat liberated by the dilution of sulphuric acid is then absorbed by the Dewar flask, so the value of the ΔH is in positive value.

Therefore, the calorimeter constant has a value of 0.338 kJ / °C.

Part 2 Enthalpy of Reaction

Part I

Amount of heat absorbed by the Dewar flask after adding nitric acid to the mixture (50 c of water + 50 c of 1M sodium hydroxide) is 2.535 kJ. This means the total enthalpy of reaction, ΔH has a value of -2.535 kJ because the heat is being released from the reaction. However, the enthalpy changes during the dilution of nitric acid also need to be considered to obtain more accurate result.

Part II

Therefore, the enthalpy of reaction of sodium hydroxide with nitric acid is -2.028 kJ.

Discussion

The calorimeter constant for the Dewar flask used in this experiment has a value of 0.338 kJ / °C. This means that for every 0.338 kJ of energy absorbed by the Dewar flask, the contents of the Dewar flask will increase by 1 °C. However, this value is not so accurate because of possible errors during the experiment. In order to obtain an accurate result, the Dewar flask must be an isolated system. This is to ensure that all the heat released to or absorbed from the surrounding remain inside the Dewar flask but not from outside of the flask. However, there are no ideal Dewar flask in this world that guarantee no heat will be loss from the Dewar flask. This is because heat can travel through vacuum by radiation and radiation cannot be controlled. So, the design of the Dewar flask can only prevent conduction and convection of heat but only able to minimize the radiation of heat.

Besides, the top of the Dewar flask is also let to open in this experiment to ease the measurement of temperature of the contents in the Dewar flask. This further caused the error in data because the system is not isolated system anymore but became an open system. Open system allowed heat loss to or acquired from the surroundings and caused the temperature difference, ∆T to drop. During the titration, methyl orange was introduced into the solution before start titration. Methyl orange is used as an indicator of the equivalence point of the titration because it has a sharper end point and the change of colour is easier to be observed. The methyl orange is red in colour when the pH value is below 3.1 and yellow when above 4.4. Any pH value in between this range will show a mixture of those two colours.

From the experiment, we know that both the reaction between sodium hydroxide with sulphuric acid and sodium hydroxide with nitric acid are exothermic reaction. This is shown in the temperature differences where the temperature of the mixture increased after both reactions. This means that heat is being released from the system to its surrounding and caused the surrounding temperature to increase. In part 2, we do not take the value of Δ as the enthalpy of reaction between sodium hydroxide and nitric acid because dilution of nitric acid also can cause change in enthalpy. Instead, we take into consideration of enthalpy change of dilution of nitric acid to obtain a more accurate result.

Besides, some of the precaution steps have to be carried out such as wearing gloves and goggle during the experiment. The bottle of the concentrated nitric acid and sulphuric acid has to be closed when not in use and put in the fume hood because the white fumes of the nitric acid or sulphuric acid can be toxic.

Conclusion

Enthalpy change, ΔH can be used to measure the heat of reaction because they are numerically equal. (ΔH = q) Dewar flask is used to create an isolated system to ensure that the data more accurate. Since every Dewar flask has their calorimeter constant, so we need to measure that constant before the experiment start. The calorimeter constant of the Dewar flask used in this experiment has a value of 0.338 kJ / °C.

In part 2 of the experiment, enthalpy of reaction between sodium hydroxide and nitric acid were to be determined. The enthalpy of dilution of nitric acid need to be considered as it can affect the results. Enthalpy of dilution of nitric acid is -0.507 kJ. After deduct this enthalpy of dilution from the total enthalpy change during the reaction, we can get the enthalpy of reaction of sodium hydroxide with nitric acid which is -2.028 kJ. By measuring the enthalpy of reaction, the reaction between sodium hydroxide and nitric acid is to be known as exothermic reaction since the energy is being released to the surrounding. Dilution of nitric acid and sulphuric acid is also known as exothermic reaction as well.

References

1. Darrell D. Ebbing & Steven D. Gammon (2009). General Chemistry Ninth Edition. Boston, MA & New York, NY: Houghton Mifflin Company. Chapter 18, page 732.

2. Raymond Chang (2005). Physical Chemistry for the Bioscience. Edwards Brothers, Inc. Chapter 3, page 46 – 47.

3. Hans Kuhn, Horst-Dieter Forsteling & David H. Waldeck (2009). Principles of Physical Chemistry. New Jersey: John Wiley & Sons, Inc. Chapter 17, page 559.

Appendix

Part 1 Calorimeter Constant

Table 1 Temperature of 100 c of distilled water in Dewar flask

Time, minutes	1	2	3	4	5
Temperature, °C	22.8	22.8	22.8	22.8	22.8

Table 2 Temperature of solution after adding 2 c of concentrated sulphuric acid

Time, seconds	15	30	45	60	75	90	105	120	135	150
Temperature, °C	31.0	31.0	31.0	31.0	28.5	28.0	28.0	28.0	28.0	28.0

Time, seconds	165	180	195	210	225	240	255	270	285	300
Temperature, °C	28.0	28.0	28.0	28.0	28.0	28.0	28.0	28.0	28.0	28.0

Table 3 Heat liberated when various quantities of concentrated sulphuric acid are added to 100 c of water

Acid added, c	Molarity of solution, M	Heat liberated, kJ
0.60	0.108	0.802
0.75	0.138	1.016
1.50	0.276	1.987
2.30	0.421	3.016
2.50	0.459	3.293

Table 4 Titration of 25 c of sulphuric acid solution against 1M sodium hydroxide.

Initial burette reading, c	15.0
Final burette reading, c	34.2

Part 2 Enthalpy of Reaction

Part I

Table 5 Temperature of mixture of 50 c of water and 50 c of 1M sodium hydroxide

Time, minutes	1	2	3	4	5
Temperature, °C	22.0	22.0	22.0	22.0	22.0

Table 6 Temperature of the mixture after adding 5 c of 10M nitric acid

Time, seconds	15	30	45	60	75	90	105	120	135	150
Temperature, °C	29.5	29.5	29.5	29.5	29.5	29.5	29.5	29.5	29.5	29.5

Time, seconds	165	180	195	210	225	240	255	270	285	300
Temperature, °C	29.5	29.5	29.5	29.5	29.5	29.5	29.5	29.5	29.5	29.5

Part II

Table 7 Temperature of 100 c of distilled water

Time, minutes	1	2	3	4	5
Temperature, °C	21.5	21.5	21.5	21.5	21.5

Table 8 Temperature of the water after adding 5 c of 10M nitric acid

Time, seconds	15	30	45	60	75	90	105	120	135	150
Temperature, °C	23.0	23.0	23.0	23.0	23.0	23.0	23.0	23.0	23.0	23.0

Time, seconds	165	180	195	210	225	240	255	270	285	300
Temperature, °C	23.0	23.0	23.0	23.0	23.0	23.0	23.0	23.0	23.0	23.0

The Use of Volumetric Flask, Burette and Pipette in Determining the Concentration of NaOH Solution

Objectives:

1. To carry out titration of a strong acid (HCl) with a strong base (NaOH).

2. To determine the end point of titration with the use of phenolphthalein as indicator.

3. To determine the concentration of base when the concentration of acid is known by doing calculations related to titration.

Introduction:

The main purpose of this experiment is to determine the unknown concentration of a known reactant. Volumetric analysis is the most common method used is titration which is a method of quantitative chemical analysis because volume measurements play a key role in titration. The technique is carried out by using a reagent of a known concentration (standard solution) and volume to react with a solution where the concentration is unknown. There are various types of titration carried out for different purposes such as acid-base titration and redox titration. Within the acid-base titration, there are four types of reactions:

(1) titration involving a strong acid and a strong base

(2) titration involving a weak acid and a strong base

(3) titration involving a strong acid and a weak base

(4) titration involving a weak acid and a weak base

The titration carried out in this experiment is the acid-base titration which is based on the neutralization reaction that occurs between an acid and a base to produce salt and water. The base is added slowly into the conical flask with acid until there are all exactly reacted and this is called the end point or also known as the equivalence point of the reaction.

The equivalence point of the neutralization reaction is the point at which both acid and base have been consumed and neither is in excess. In other words, the hydrogen ion and hydroxide ion concentrations are equal. In order to determine the equivalence point in a titration, acid-base indicators need to be added to the acid solution before the titration start. The end point of a titration is indicated when the indicator changes color. An indicator is a weak organic acid or base that has distinctly different colors in its non-ionized and ionized forms. This will be discussed further in discussion section. Different indicators show different colour changes at the same pH, therefore choosing of indicator for a particular titration depends on the acid and base used.

Hydrochloric acid (HCl) and sodium hydroxide (NaOH) is used as the reactants in this experiment. HCl is a monoprotic acid which dissociate to give out one H⁺ ion. Monoprotic acids have acid dissociation constant, K_a, which indicates the level of dissociation in water. Hydrochloric acid has large K_a value as it is a strong acid and dissociate completely in water. NaOH is a metallic base and ionic which composed of sodium cation and hydroxide anion. The hydroxide ion makes sodium hydroxide a strong base which reacts with acids to form water and corresponding salts.

Procedure:

Results & Calculations:

Burette readings (mL)	Titration 1	Titration 2	Titration 3
Initial reading	0.00	0.00	1.10
Final reading	26.00	24.80	26.40
Volume of NaOH solution used (mL) (Final reading – Initial reading)	26.00	24.80	25.30

 

	Volume of NaOH used (mL)
Titration1	26.00
Titration 2	24.80
Titration 3	25.30

 

NaOH _(aq) + HCl _(aq) NaCl _(aq) + H₂O _(l)

Information given:

Volume of NaOH used = = 0.0253667 dm³

Volume of HCl used = = 0.025 dm³

Concentration of HCl = 0.01M

No. of moles HCl, n =

=

=

Based on equation, 1 mole of HCl reacts with 1 mole of NaOH.

Hence, 0.00025 mole of HCl reacts with 0.00025 mole NaOH.

(Concentration of diluted NaOH)

Dilution, M₁V₁ = M₂V₂

M₁ (5) = (0.00986) (250)

M₁ =

(Initial concentration of NaOH)

 

Discussion:

The titration of a strong acid and strong base in this experiment can be illustrated with a graph called a titration curve. It is a graph of pH versus volume of the solution titrated. The figure below represents the pH versus volume data of the titration curve for the HCl-NaOH titration. From the graph we may explain the chemical changes happen during titration and decide which indicators best to be used to determine the endpoint which matches the equivalence point of the neutralization.

Based on the graph, the pH has a low value at the beginning of the titration which shows the concentration of the HCl in conical flask. As the titration proceeds, the pH changes slowly until it reaches just before the equivalence point. At the equivalence point, the pH rises sharply by just adding only two drops of base. This is because when the equivalence point is reached, the number of moles of hydrogen ions and hydroxide ions is equal to each other; therefore a slight addition of base will result in a steep increase of pH. Beyond the equivalence point, the pH again rises only slowly. According to the graph, any acid-base indicator whose color changes in the pH range from about 4.0 to 10.0 is suitable for HCl-NaOH titration.

The acid-base indicator that is used for the titration of HCl-NaOH is phenolphthalein which is situated within the pH range of 8.3 to 10.0. Other indicators only have pH range within 1.0 to 8.8 which does not include the pH range beyond 9.0 as phenolphthalein where the pH range is until 10.0. Based on the interpretation of the graph, indicator whose color changes in the pH range from 4.0 to 10.0 is only suitable for the HCl-NaOH titration. Therefore phenolphthalein is chosen rather than the other indicators.

As we known an indicator is usually a weak organic acid or base that has distinctly different colour in its non-ionized and ionized forms, but what are the characteristics of an indicator that make let us determine the endpoint of a titration by showing different colours. Acid-base indicators exist in two forms, a weak acid represented as HIn and having one colour and its conjugate base represented as In- and having a different colour. The indicator does not affect the pH of the solution if only just a small amount of indicator is added to a solution. However, the ionization equilibrium of the indicator is affected by the concentration of H₃O⁺ in the solution. When in a solution, the acid ionizes to the following ions:

HIn_(aq) + H₂O ↔ H₃O⁺_(aq) + In^-_(aq)

acid colour base colour

Based the Le Châtelier’s principle, increasing [H₃O+] in the solution shifts the equilibrium to the left, increasing the amount of HIn and thus showing the acid colour. On the other hand, decrease in [H₃O+] in a solution shifts the equilibrium to the right, increasing the amount of In^- and hence displaying the base colour. In general, if 90% or more of an indicator is in the form HIn, the solution will show the acid colour. If 90% or more is in the form In^-, the solution takes on the base colour. If the concentrations of HIn and In^- are about equal, the indicator is in the process of changing from one form to the other and has an intermediate colour which is the mixture of acid and base colour.

Based on the results obtained, within the three titration carried out only Titration 3 is less than 3 and within the range. However, Titration 1 and 2 is more than 3 and is out of the range. The causes of the results to be out of range can be due to human errors. First of all, the NaOH solution could have been diluted as the burette used to fill the NaOH solution is rinse with distilled water and not with NaOH before use. This causes the concentration of NaOH to be slightly different from the standard solution that has been prepared. The same possibility does happen to the conical flask which is used to fill the HCl which will be titrated against NaOH where the flask only rinse with distilled water not with HCl. Thus, the concentration of HCl may be less than 0.01M. Besides that, the reading on the burette could have some minor error because during the recording of readings the meniscus shown on the burette is not clearly view. In order to correct the error, a white paper should be situated behind the burette in order to have a clearly view on the position of the meniscus so that a more accurate readings can be obtained.

Precaution steps:

Firstly, is the determination of the titration end point which is based on the colour changing of the indicator added to the conical flask. Confusion arises about when to stop the titration as the colour changes is difficult to be observed. Therefore, a white tile or a piece of white paper should be placed at the bottom of the conical flask so that the changes of colour can be easily seen. Next, NaOH solution will react quickly with the carbon dioxide in the air. Therefore NaOH should be cover when it is not in use. This is the reason the prepared standard solution of NaOH is closed with the cap. Lastly is the dilution or the preparation of the standard solution of NaOH. The standard solution is prepared in a 250mL volumetric flask by adding 5mL of NaOH and distilled water should be added to the graduated line in the flask. During the adding of distilled water, water could have overshoot the line and cause the concentration of the standard solution to be different from the expected concentration. Thus, use dropper to fill the water into flask when the meniscus level approaches to the graduated line of the flask to avoid the overshooting of distilled water during the preparation of standard solution.

Thin-Layer Chromatography and Column Chromatography

Objectives

Part I

1. To learn the technique of TLC and the visualization of colorless components.

2. To identify an unknown drug by a TLC comparison with standard compounds.

Part II

1. To learn the technique of column chromatography.

2. To separate the mixture of pyrene and p-nitroaniline by column chromatography.

Introduction

Chromatography is a method for separating complex samples into their constituent parts, and it is the most important procedure for isolating and purifying chemicals. There are few types of chromatography used to separate different complex samples. Types of chromatography commonly used included gas chromatography (GC), high performance liquid chromatography (HPLC), thin layer chromatography (TLC), and column chromatography (CC). In this experiment, we are only using the TLC and CC method to separate our sample mixtures. In chromatography, there are two phases involved in separating the samples, which are the stationary phase and mobile phase. Stationary phase is the part of the chromatographic system though which the mobile phase flows where distribution of the solutes between the phases occur. The stationary phase may be a solid or liquid that is immobilized or adsorbed on a solid. Examples of substances used as stationary phase included silica gel or alumina used in the TLC and CC, and filter paper was used as stationary phase in the paper chromatography. Alumina is generally used for chromatography of less polar compound whereas silica gel is better for compounds containing polar functional group. Mobile phase is the part of the chromatographic system which carries the solutes through the stationary phase. The mobile phases are either liquid or gases. The mobile phase is often known as eluent.

Solvent that can be used as eluent in TLC and CC can be categorized into three category: non-polar, moderate, and polar solvent. The following diagram shows some examples of solvent with different polarity

According to the principle of chromatography, different compounds will have different solubility and adsorption to the two phases between which they are to be partitioned. The samples are put on the stationary phase and are then carried along by the mobile phase. Since different compounds have different degrees of adsorption, the migration rate for each compound will be different and thus allow separation of compounds. The adsorption of compounds onto the stationary phase is depends on different types of interactions between the samples and the stationary phases. Examples of types of interactions included ion-dipole, dipole-dipole, hydrogen bonding, dipole-induced dipole, and Van Der Waals forces. In this experiment, silica gel was used as the stationary phase. The following diagram shows the structure of silica gel used as the stationary phase.

Silica gel is a porous form of SiO₂ and the surface of gel contains Si-OH and Si-O-Si functional groups. As shown in the diagram above, the silica gel consists of polar functional group. Hence, the dominant interactive forces between the adsorbent (silica gel) and the materials to be separated are of dipole-dipole type. Highly polar molecules will interact fairly strong with the polar Si-O bonds in the silica gel and will tend to adsorb onto the fine particles of the adsorbent. The stronger adsorption will caused the material passes through the system to be slower. On the other hand, weakly polar molecules will held less tightly with the silica gel and thus tend to move through the adsorbent more rapidly than the polar species. The adsorption strengths of each compound having the following types of functional groups are arranged in the order of increasing group polarities:

However, variation may occur depending on the overall structure of each specific compound. The adsorption of polar molecules can be altered by using polar solvent. The more polar the eluent, the greater the eluting power, thus allowing compounds to move faster over the adsorbent surface. This is because polar solvent will compete for the silica adsorption sites, thus caused the polar molecules not to adsorb strongly on the silica gel. This caused all the compounds to travel at a faster rate but the order in which the compounds move remains the same. However, this will caused the separation between the non-polar compound and polar compound to become smaller and poor results might be obtained. Thus, a mixed eluent might be better for separating a mixture of samples consists of non-polar and polar compounds but the polarity of the mixed eluent must be controlled. Mixed eluent can be prepared by mixing low polarity and high polarity solvents and hence creating any eluting power needed.

In the TLC technique that we are going to applied in this experiment, a mixture of solvent ethyl acetate : hexane, 1 : 3 was used as the eluent. The silica gel (TLC plate) was used as the stationary phase. A capillary spotter was used to spot the samples on the TLC plate about 1.5 cm from the bottom of plate. The plate was then placed into a closed developing chamber which has a shallow layer of solvent that does not submerge the spot. The chamber is lined with a folded piece of filter paper to ensure a uniform and saturated atmosphere of solvent vapor. After the solvent has reached about 0.5 cm from the top of the plate, the plate was removed from the developing chamber. The capillary action of the solvent causes the initial spot to be separated into individual components that may be visualized by color identification or with the following techniques for colorless compound: (a) irradiation with ultraviolet light, (b) reversible staining with iodine vapor, and (c) spraying a reagent that irreversibly colors the spots. The rate at which a compound moves in respect to the solvent front, R_f, is characteristic of that compound under standard conditions. The R_f values can be calculate by using the following equation:

R_f = Distance of Compound / Distance of Solvent

TLC can be used to separate a very small amount of sample easily and rapidly without any costly equipment. Thus, it is often used to monitor the progress of a reaction by running the crude sample beside the reaction sample on the same plate. It also can be used to determine the best eluent for subsequent separation by column chromatography. Some of the common uses of TLC included:

(a) To determine the number of components in a mixture.

(b) To determine the identity of two substances.

(c) To monitor the progress of a reaction.

(d) To determine the effectiveness of a purification.

(e) To determine the appropriate conditions for a column chromatographic separation.

(f) To monitor column chromatography.

Column chromatography is a technique that uses an adsorbent packed in a glass column, and a solvent that moves down slowly through the packed column. Similar to TLC, silica gel was used as the stationary phase. The eluent used also the same which is the mixture of ethyl acetate : hexane in a ratio of 1 : 3. The column is packed with silica gel and the solvent was allowed to drain as the silica packs tightly. Small layer of sand was applied on top of the silica gel to avoid disturbance of silica gel during the adding of fresh solvent. The mixture of compounds was dissolved in small amount of solvent and poured into the column. The more polar compound will adsorb strongly on the silica gel and only move a little with every addition of fresh solvent. The less polar compound will move faster down the column due to less interaction with the silica gel. Fresh solvent is continually added to the top of the column until each band resolves and is carefully collected. The movement of solvent down the column due to the gravitational forces will caused the mixture of compounds to be separated. When the substances are colored, we can directly observe and collect them as they drain off. However, when the compound is colorless, several small fractions of the eluting solvent must be collected and testing each by thin layer chromatography.

Apparatus & Materials

Part I

UV lamp, capillary tube, 250 mL beaker, aspirin, acetaminophen, caffeine, unknown A, unknown B, TLC plates, ethyl acetate, hexane, and iodine.

Part II

Glass column, UV lamp, capillary tube, 250 mL beaker, test tubes, glass funnel, pyrene, p-nitroaniline, TLC plates, ethyl acetate, hexane, and iodine.

Procedures

Part I Analysis of Analgesic Drugs

Part A Spotting of the TLC plates

1. A TLC plate was obtained from the instructor.

2. The plate was set down on a clean, dry surface, and then a line was drawn lightly across the plate about 1.5 cm from the bottom and 0.5 cm from the top of the plate with the aid of a 2B pencil.

3. Five 2-3 mm lines was made, spaced about 0.6 cm apart and running perpendicularly through the line across the bottom of the TLC plate.

4. The first and the last line were drawn such that they are about 0.8 cm from the edge of the plate.

5. Separate capillary tube was used for spotting each sample on the TLC plate.

6. Acetaminophen solution was first spotted, then the caffeine, then the unknown A, then aspirin, and lastly the unknown B.

7. The spots are made as small as possible to avoid “tailing” when the plate was developed.

8. The plate was examined under the ultraviolet (UV) light to see that enough of each compound has been applied; if not, more sample were added.

Part B Developing the TLC plates

1. A developing chamber was developed using a 250 mL beaker as the chamber, a half-piece of filter paper inside, and aluminium foil to cover.

2. The eluent, 1 : 3 mixture of ethyl acetate : hexane, was poured into the beaker to a depth of under 1 cm (about 15 mL).

3. The prepared TLC plate was placed in the developing chamber and the solvent level was ensured to be below the pencil line.

4. After the solvent has risen to near the top of the plate (about 0.5 cm from the top), the plate was removed.

5. The solvent was allowed to evaporate from the plate in the fumehood.

Part C Visualization

1. The colorless compounds are visualized by illumination of the plate with an ultraviolet (UV) lamp.

2. The spots were outlined with a 2B pencil.

3. The spots may also be visualized by putting the plate in an iodine chamber for a couple minutes.

Part D Comparison of the unknown with reference standards

1. The plate was sketched in the notebook and the R_f values for each spot was calculated.

2. The unknown drug was determined based on the R_f value.

Part II The Separation of Pyrene and p-Nitroaniline by Column Chromatography

Part A Column preparation

1. A 40 cm chromatography column, 15 g of deactivated silica gel and 110 mL of the developing solvent mixture (ethyl acetate : hexane, 1 : 3) were obtained.

2. Slurry of the adsorbent (silica gel) was prepared with 50 mL of solvent in a 250 mL Erlenmeyer flask.

3. The column was clamped in a vertical position.

4. The stopcock of the column was closed and 15 mL of solvent was poured in.

5. After setting, all of the slurry was quickly decant through a funnel into the column.

6. The stopcock was opened and solvent was allowed to drain while the wall of the column was tapped with the ends of a folded piece of rubber tubing.

7. Once the solvent level is within 6 cm of the top of the adsorbent, a 0.5 cm level layer of sand was added on the adsorbent.

8. Excess solvent was drained off until its level is precisely on top of the sand and the stopcock was closed.

Part B Separation and collection of pyrene and p-nitroaniline

1. The pyrene (0.02 g) / p-nitroaniline (0.02 g) mixture was added with a few drops of ethyl acetate to dissolve as much of the mixture as possible.

2. The solution was carefully transferred to the top of the sand layer with a dropper.

3. The solvent was drained off until the mixture solution is just below the top of the sand.

4. The wall was rinsed with about 1 mL o fresh solvent (ethyl acetate : hexane, 1 : 3) and drained until the level is once again below the top of the sand.

5. The rinsing of the wall was repeated until the solvent above the silica gel is virtually colorless.

6. The column was filled with fresh solvent very carefully and the solvent was allowed to drain.

7. The separation of bands was observed as the column developed.

8. The colorless band of pyrene was collected into four fractions.

9. When the edge of the yellow band (p-nitroaniline) reached the lower part of column, a new test tube was replaced and the yellow band was collected into four fractions.

10. Each fraction was concentrated to a small volume by evaporation for analysis by TLC.

Part C Analysis of the fractions

1. The fractions were spotted on a TLC silica gel plate along with the reference pyrene and p-nitroaniline.

2. The TLC plate was developed in a developing chamber containing a 1 : 3 micture of ethyl acetate : hexane.

3. The TLC plate was visualized with the ultraviolet (UV) light to determine which fractions are pure pyrene and which are pure p-nitroaniline.

4. The chromatogram was drawn in the lab notebook.

5. The R_f values for pyrene and p-nitroaniline were calculated.

Results & Calculations

Part I Analysis of Analgesic Drugs

Distance of solvent travelled = 8.0 cm

Distance travelled by compounds: (a) Acetaminophen = 0.50 cm

(b) Caffeine = 0.35 cm

(c) Unknown A = 0.50 cm

(d) Aspirin = 2.40 cm

(e) Unknown B = 2.20 cm

R_f values of compounds: (a) Acetaminophen = 0.0625

(b) Caffeine = 0.0438

(c) Unknown A = 0.0625

(d) Aspirin = 0.3344

(e) Unknown B = 0.2750

Diagram of acetaminophen, caffeine, and unknown A, aspirin, and unknown B travelled on the TLC plate

Conclusion: Unknown A belongs to the acetaminophen whereas unknown B belongs to the aspirin.

Part II The Separation of Pyrene and p-Nitroaniline by Column Chromatography

Distance of solvent travelled = 8.0 cm

Distance travelled by compounds: (a) Pyrene: (i) First fraction = 6.10 cm

(ii) Second fraction = 6.10 cm

(iii) Third fraction = 6.10 cm

(iv) Fourth fraction = 6.10 cm

(b) p-Nitroaniline: (i) First fraction = 1.60 cm

(ii) Second fraction = 1.70 cm

(iii) Third fraction = 1.70 cm

(iv) Fourth fraction = 1.70 cm

R_f values of compounds: (a) Pyrene: (i) First fraction = 0.7625

(ii) Second fraction = 0.7625

(iii) Third fraction = 0.7625

(iv) Fourth fraction = 0.7625

(b) p-Nitroaniline: (i) First fraction = 0.2000

(ii) Second fraction = 0.2125

(iii) Third fraction = 0.2125

(iv) Fourth fraction = 0.2125

Diagram of pyrene and p-nitoraniline travelled on the TLC plate

Conclusion: Pyrene is present in the 1^st to 4^th spots while p-nitroaniline is present in the 5^th to 8^th spots.

Discussion

The eluent used in this experiment is a mixture of ethyl acetate : hexane in a ratio of 1 : 3. Ethyl acetate, CH₃COOCH₂CH₃, is a polar solvent because of the oxygen bonded to carbon in the compounds has higher electronegativity than carbon. So, the oxygen atom in the ethyl acetate will become partial negative while the carbon atom will acquire a partial positive charge. Hexane, CH₃(CH₂)₄CH₃, is an organic solvent which is non-polar. Mixture of these two solvents gives the eluent polarity properties and allows the polar compounds to be separated as well. If the polarity of the eluent is too low, the polar compounds will not be able to carry by the eluent and will not be separated. However, if the polarity of the eluent is too high, the polar compound will travel so fast that the separation between non-polar compound and polar compound to become so small, which result in poor separation. Hence, the polarity of the eluent must be optimal for the separation of the samples during chromatography. A few drops of acetic acid have been added to the eluent to prevent the deprotonation of the compounds during the separation of compounds. This is an important step for retain the identity of the compounds and hence, a pure compounds can be separated out.

In Part I of this experiment, TLC plate have been used for the separation of the compounds. Silica gels were coated on the TLC plate and act as the adsorbent. Silica gel consists of polar functional group such as Si-O-Si and Si-OH which results from the difference of electronegativity of the of oxygen atom with the silicon atom. Since the silica gel is polar, polar compounds will then interact strongly with the silica gel and adsorb on the silica gel strongly. This will caused the polar compounds to travel only a short distance from the bottom of the TLC plates as the eluent move upward. As for the non-polar compounds, they interact weakly with the silica gel and thus can be carried easily by the eluent. As the eluent moving upwards, the non-polar compounds can travel further away from the bottom of the TLC plate. The moving of eluent along the TLC plate is due to the capillary action of the solvent, allowing the solvent to travel in the gap present in the three dimensional network of silica gel. This causes the initial spot to be separated into individual components as the samples were carried by the eluent. A folded filter paper was inserted into the developing chamber to help the eluent move upwards quicker. This is because when eluent travels up along the filter paper, it will vaporize easily due to the larger surface area. Since the developing chamber is covered with aluminium foil, the eluent vapor will remain in the chamber. When the chamber is saturated with eluent vapor, the eluent will then move upwards quicker along the silica gel without vaporization. This promotes the rate of eluent travelled along the silica gel without affecting the separation of the compounds.

Three compounds which were spotted on the TLC plate are acetaminophen, caffeine and aspirin. The following diagrams show the structure of these compounds.

              Acetaminophen                        Caffeine                                 Aspirin

As the diagrams shown, all these three compounds are polar compounds. However, the polarities of the compounds may differ from each other. This is because of the electronegativity difference between the atoms in the compound. Nitrogen and oxygen atom present in these compounds have higher electronegativity compared to the carbon atom and hydrogen atom. Thus, nitrogen and oxygen atom will acquire a partial negative charge while the carbon and hydrogen atom will acquire a partial positive charge.

According to the results obtained, aspirin can travel the farthest away from the bottom of the TLC plate, followed by acetaminophen and caffeine. This is shown by the R_f values of each compound in which the higher R_f value indicate that the compound can travel further away and thus is less polar. The R_f value for aspirin is 0.3344, followed by acetaminophen, 0.0625 and finally the caffeine, 0.0438. Therefore, aspirin is the least polar compounds among those three compounds, followed by acetaminophen and finally the most polar compound is the caffeine. This may be due to the presence of large number of electronegativity atom such as nitrogen and oxygen atom in the compounds, thus increase the polarity of the caffeine compound. Since the compounds with similar polarities will travel with the same distance along the TLC plate, we can identify the unknown compounds by comparing the R_f values of the unknown with the standard compounds. Unknown A and B have the R_f values of 0.0625 and 0.2750 respectively. Unknown A has the same R_f value with the acetaminophen, thus we can conclude that unknown A is acetaminophen. Unknown B has almost similar R_f value with aspirin, so we can deduce that unknown B belongs to the aspirin.

In Part II of this experiment, slurry of silica gel was used in column chromatography. The function of silica gel in the column is similar to the silica gel on the TLC plate, which act as the stationary phase for the materials to be separated to adsorb. The two compounds which were mixed together and separated by using column chromatography technique are pyrene and p-nitroaniline. The structures of the compounds are shown in the following diagram.

                      Pyrene                                                   p-Nitroaniline

According to the structure of pyrene, it is a non-polar compound. This is because only carbon and hydrogen atom present in the compound and the difference in electronegativity of carbon and hydrogen atom is very small only. On the other hand, p-nitroaniline consists of highly electronegativity atoms in its compound which is nitrogen and oxygen atom. This makes the p-nitroaniline a highly polar compound and thus will adsorb strongly to the silica gel. It was then can be predicted that pyrene will be collected first from the column before p-nitroaniline. This is because pyrene which is non-polar compound is weakly adsorb onto the silica gel and will be carried by the eluent easily as the eluent drained off from the column due to gravitational forces. The mixture of pyrene and p-nitroaniline was dissolved in minimum volume of ethyl acetate before added to the column. If too much of the solvent used, it will results in a poor separation of the compounds. This is because when the compounds dissolved in high volume of solvent, certain portions of the pyrene will be on top of the column while the other will be at the bottom of the column. This happened similarly to the p-nitroaniline and caused the separation of the pyrene and p-nitroaniline not complete, which results in poor separation.

According to the results obtained, the first four fractions of eluent collected in test tubes consist of pyrene only and the last four fractions contain p-nitroaniline only. This is proven by doing the TLC on the fractions of pyrene and p-nitroaniline collected. The first four fractions of test tubes containing pyrene were spotted on the TLC plate with four separate spots. Similar procedures were done to spot the p-nitroaniline in the last four fractions of test tubes on the same TLC plate. It was shown that pyrene move farther on the TLC plate compared to the p-nitroaniline. This proved that p-nitroaniline is more polar compared to the pyrene. However, the first and the fourth spot which are the spots of pyrene have a very fade color under the ultraviolet light. This might be due to low concentration of pyrene at the beginning and at the end of collection of colorless band before the yellow band was collected. At the beginning, most of the pyrene still moving from the top of the column and haven’t reach the bottom of the column. So the concentration of the pyrene is lower. At the end of collection of colorless band, since most of the pyrene already drained off, so there is lower concentration of pyrene as well. Same thing happened to the fifth and the eighth spot which are the spots of p-nitroaniline. The yellow color of p-nitroaniline is fader for the fifth and eighth spot because of lower concentration of p-nitroaniline. It was also observed that the four spots of pyrene on the TLC plate have mixed together during the TLC was carried out. This may be because of the pyrene is non-polar and thus interact weakly with the silica gel. When eluent passed through the spots, they were carried easily by the eluent and thus dissolved together to form a large spot on the TLC plate. Another reason might be due to human error in which the diameter of the spots were too big and thus the pyrene can easily interact with each other and mixed together.

Precaution Steps

1. Do not look directly at the ultraviolet lamp.

2. Do not interrupt the beaker after TLC plate has been introduced into the developing chamber.

3. Do not bend the silica gel excessively since the silica adsorbent may flake off.

4. Do not stop the collection of fractions of eluting solvent too long as the compounds may settle down to the bottom of the column and mixed together.