Acids And Bases Lab Report
Acids & Bases: Reactions, Standardizations, & Titrations Experiments 21 & 22
Experimental Overview:
The procedure for this experiment was carried out as instructed in the laboratory manual, Experiments in General Chemistry, 4th ed., S.L. Murov, Experiment 21, Acids and Bases: Reactions and Standardizations, and Experiment 22, Acids and Bases: Analysis. There were modifications made by the instructor to dilute the 6M NaOH to 0.1M in 300mls instead of 500mls in Part B of Experiment 21.
One objective for performing these two experiments was to observe qualitatively the reactions between common acids, bases, and the indicators phenolphthalein, methyl orange, bromothymol blue, and red cabbage extract. Another focus of this experiment was to perform a …show more content…
standardization of sodium hydroxide solution using the titration technique with potassium hydrogen phthalate. By titrating the standardized NaOH, determine the percent by mass of acetic acid in an unknown vinegar solution as well as determine the molecular mass of an unknown acid.
Background Information:
When an acid such as HCl and base such as NaOH are mixed with each other they neutralize and produce salts and react as follows:
HCl + NaOH ³ NaCl + H2O + Heat
Heat can be quantitatively measured using a thermometer as a result of the reaction forming water. Other times, combining an acid such as HCl with a base such as Na2CO3 will result in the following:
2HCl(aq) + Na2CO3 (aq)³ 2 NaCl + CO2 (g) + H2O (l)
Another unique property of acids and bases are their high reactivity with certain metals. Often their reactions will often be violent.
NaOH + Al(s) + ³ AlO (s) + H2 (g) + Na+ (aq)
Chemical indicators of acids or bases work because their molecules change shape in the presence of ions. The following indicators were used in this lab:
Methyl Orange is often used as an indicator for strong acids. In the presence of H+ ions the indicator turns a bright red, but in a pH of about 4 it is orange, and in more basic solutions it turns yellow. Bromothymol blue is a more diverse indicator because it turns yellow below pH 6, green at neutral 7, and blue at 8 and above. Red cabbage juice also turns different colors in different pH. It turns red in an acid and green in a base.
Phenolphthalein is an indicator for bases that turns from clear to pink in the presence of OH- ions. This is due to fact that phenolphthalein molecules are clear, but when they ionize in solution, their H+ ions that make them a weak acid, are pink.
The titration procedure is a quantitative analysis which is used to determine the concentration of a known reactant. A calibrated burette is used to add a reagent to determine the exact amount. The titration is stopped when the solution slightly changes due to the indicator. Indicators such as phenolphthalein are useful in determining the molarities of bases such as NaOH. Through titration, NaOH is carefully added into a specific amount of diluted KHP (potassium hydrogen phthalate 204.22 g/mol) with a few drops of phenolphthalein. When the system can no longer equilibrate the NaOH being added to it because all the KHP, will be used up, there will be an excess of NaOH which the phenolphthalein will indicate by turning pink. By knowing the amount of NaOH titrated, we can calculate the average molarity of OH- reacted because it will react in a one to one ratio with the monoprotic KHP. Titration is a necessary method because it is inaccurate to calculate the amount of OH- ions using the molarity of NaOH, because each dilution of the solution contains a different amount of ions. This process of determining the concentration of a solution is called standardization. Vinegar is produced by the fermentation of sugar to ethanol and then bacteria catalyze the oxidation of the ethanol to acetic acid. The legal minimum is 4% by mass acetic acid in vinegar. Using the standardized NaOH, it is also possible to determine the percent by mass of acetic acid contained in vinegar as well as the molar mass of and unknown monoprotic acid. Like KHP, acetic acid, HC2H3O2, is an example of a monoprotic acid because in solution, it only releases one H+ ion, becoming H+(aq) and C2H3O2-(aq).
Results and Calculations:
Table 1 summarizes the experimental results of experiment 21 part A. 1 through 4 and includes the reagents added for each step. The reagents which were added are listed on the left, as well as their balanced equations. Observations are displayed on the right hand side. This table summarizes the common reactions of acids and bases as well as the reaction time for an indicator to work.
Table 1. Results of Reactions with Acids and Bases.
Reagents added Observations
2mls of 3M HCl + 2mls of 3M NaOH HCl + NaOH ³ NaCl + H2O + Heat The initial temperature of the 3M HCl was measured at 24 C . When 3M NaOH was added and stirred, the temperature was found to be 28 C.
1mL of 3M HCl + 1mL of 1M Na2CO3 in a large test tube
2HCl + Na2CO3 ³2 NaCl + CO2 + H2O Slight fizzing when the solutions were combined. When a lighted splint is inserted into the tube, the flame get smaller and goes out.
3mLs of 3M NaOH + wad of aluminum foil
2NaOH + 2Al + ³ 2AlO (s) + H2(g) + 2Na+ (aq) A lot of loud fizzing, gas can be seen forming. When lighted splint is inserted, the flame immediately is sucked in and a loud pop noise is made.
Solutions A+B
Solutions A+C It took 13 seconds for the solution to turn pink
It took 32 seconds for the solution to turn pink
Table 2 shows the color change of four indicators when they are added to three solutions with different pH¡¦s . The original colors of the indicators are included on the left.
Table 2.
Color Chart of Indicators in 0.01M HCl, pH 7 Buffer, and 0.01 M NaOH
0.01M HCl pH 7 Buffer 0.01 NaOH
Phenolphthalein
(clear) Cloudy precipitate No change Pink
Methyl orange
(yellow-orange) Reddish pink No change Yellow
Bromothymol blue
(green) No change No change Blue
Red-cabbage extract (purple) Pink No change Green
Table 3 includes data collected throughout the process of standardizing the sodium hydroxide solution. Information from all three titrations of sodium hydroxide into KHP is recorded in order to calculate the average molarity of the NaOH solution in the end. Calculations used to derive the numbers in the table will be shown after the table.
Table 3. Data collected from three titrations of 0.1M NaOH into 50mls KHP with 3 drops of phenolpthalein.
Trial 1 Trial 2 Trial 3 Average
Mass of KHP 1.001g 1.002g 1.001g
Moles of KHP 0.00490mols .00491mols .00490mols
Initial Buret Reading 30.0mL 30.0mL 30.0mL
Final Buret Reading 9.8 mL 10.0mL 10.0mL
Volume of NaOH Solution 20.2mL 20.0mL 20.0,L
Molarity of NaOH solution 0.243M 0.246M 0.243M 0.2439M
Deviation From average 0.0009 0.0016 0.0009
.00113
The moles of KHP were determined as follows:
Mass of KHP x Molar Mass of KHP = number of moles KHP
From trial 1:
The moles of KHP and the moles of NaOH are equivalent therefore the molarity of NaOH is found by dividing the number of moles NaOH (same as moles of KHP) by the volume of NaOH titrated. This calculation was repeated for all three trials.
Molarities of NaOH:
From Trial 1:
The average molarity from all three trials is then calculated as 0.2439M.
Further analysis is done by determining each trial¡¦s deviation from the average.
Average Molarity ¡V Experimental Molarity = Deviation from Average
From Trial 1:
Table 4 summarizes the data collected from three titrations of the standardized NaOH into vinegar with an unknown concentration of acetic acid (unknown #1).The buret readings indicate the volume of NaOH titrated. In the end, the average mass percent of acetic acid in the unknown vinegar was calculated. Calculations are shown below.
Table 4. Data from three titrations of standardized NaOH into unknown vinegar #3.
Trial 1 Trial 2 Trial 3 Average
Mass of Flask 98.3414g 96.8691g 74.8863g
Mass of Flask + solution 101.0999g 101.2024g 79.6695g
Mass of solution 2.7585g 4.3333g 4.7632g
Density of vinegar 0.5517g/ml 0.8666g/ml 0.9566g/ml
Molarity of NaOH used 0.2439M 0.2439M 0.2439M
Initial Buret Reading 30.0mL 30.0mL 30.0mL
Final Buret Reading 15.7mL 15.2mL 15.5mL
Volume of NaOH solution 14.3mL 14.8mL 14.5mL
Moles of NaOH 0.00349mols 0.00361mols 0.00354mols
Moles of Acetic Acid 0.00349mols 0.00361mols 0.00354mols
Molarity of Acetic Acid 0.698M 0.722M 0.708M 0.709M
Deviation of each molarity from average 0.011 0.013 0.001M 0.0083
Mass Percent of Acetic Acid in Vinegar 7.59% 5.00% 4.44% 5.68%
From Trial 1: Since acetic acid is monoprotic, it also reacts in a 1:1 ratio with NaOH, so the number of moles of NaOH was equated to the moles of acetic acid.
From Trial 1:
The mass percent of acetic acid in vinegar was calculated as follows:
From trial 1:
Table 5 includes data collected from three titrations of NaOH into unknown acid #1. For each trial, a molecular mass is determined and then averaged in the last column. Calculations are shown below.
Table 5. Data from three titrations of standardized .2439M NaOH into an unknown acid #1 with phenolphthalein indicator.
Trial 1 Trial 2 Trial 3 Average
Mass of Unknown 0.5023g 0.5021g 0.5082g
Initial Buret Reading 30.0mL 30.0mL 30.0mL
Final Buret Reading 20.5mL 20.4mL 20.0mL
Volume of NaOH titrated 9.50mL 9.57mL 10mL
Moles of unknown acid (assuming monoprotic) 0.00232mols 0.00233mols 0.00233mols
Molecular Mass of acid 216.50g/mol 215.49g/mol 215.79g/mol 215.93g/mol
Deviation of MM from avg 0.57 0.44 0.14 0.38
The moles of the unknown acid were calculated using the molarity of the standardized NaOH multiplied by the liters of NaOH titrated into the unknown acid, which is equal to the moles of monoprotic acid reacted.
For trial 1:
Since molecular mass is in units of g/mol, the mass of the unknown acid is divided by the moles of acid calculated above.
For trial 1:
Discussion
Three experiments involving reactions between acids and bases were done to test and observe their unique properties. These included neutralization, gas evolutions, as well as a basic redox reaction. Various common indicators were also tested out by combining them with acids, bases, and a neutral solution as well to witness how they work. The standardization of NaOH was also done for use in determining the percent of acetic acid in vinegar and the molecular mass of an unknown acid.
When HCl is added to NaOH, the result was a rise in temperature by 4 degrees because when water is formed during neutralization, heat is also released as a product therefore raising the temperature of the solution. The results of my experiment proved that this is a true neutralization reaction. When 1mL of 3M HCl was added to 1M Na2CO3 and a lighted splint was inserted, the flame went out because of the CO2 being liberated from the reaction, much like how a fire extinguisher works. However, in the following experiment where HCl was combined with Na2CO3, the flame that was inserted immediately got pulled into the test tube making a loud pop because of the vacuum being formed. This occurred because hydrogen gas was being released from the reaction, which is flammable. My results agreed with what happens when a flame is inserted into a flammable gas and a nonflammable gas.
When NaOH and solid aluminum were combined, the reaction was so violent because aluminum has a high reactivity with acids and bases. The aluminum becomes oxidized as it goes from a 0 to a +2 oxidation state to form solid aluminum oxide, which is the dark precipitate that formed. The chemical equation 2NaOH (aq)+ 2Al(s) + ³ 2AlO (s) + H2(g) + 2Na (aq) corresponds with the results of my experiment.
During the mixture of the two solutions A+B and A+C, the method used for timing was too arbitrary, because I had to keep glancing up at the clock, whose second hand was not in direct view to me which resulted in error in the time it took to react. I am unsure why a precipitate formed when phenolphthalein was combined with HCl, unless this is just how phenolphthalein reacts in an acid, but to be sure I would test it using different acids. All of the indicators remained the same color in the pH 7 buffer because the buffer solution is neutral, because the indicators only react when there are ions present. Bromothymol blue was the only indicator which didn¡¦t change color in acid because it is unreactive with the H+ ions.
Three titrations of NaOH were used to determine an average concentration to ensure a more accurate number since each batch of NaOH collected from the stock solution will always have different concentrations of OH-, in order to help reduce error. Practice using the buret was also helpful in greatly reducing error during titration, because I found that adding even the slightest drip would taint the results and turn the solution dark pink. By starting the titration at 30.0mLs in each buret for all trials, I think I also reduced a source of error. The molarity of my NaOH solution turned out to be 0.2439M. It is shown in Table 3 that trial 2 had the greatest deviation from the average, which may have been caused by a larger mass of KHP, 1.002g initially, compared to 1.001g. There may have been an error during the weighing of the flasks in Table 4 for trial 3 because the weight of flask 3 was so much lower than the other two even though I used the same type of flask. According to the averages of my three trials, unknown vinegar #3 had 0.709mols of acetic acid per liter. However, this number was not used in calculating the average mass percent of acetic acid in vinegar, because I thought it would be more accurate to calculate for each trial and then take the average, which is 5.68%. Since it is a small percentage, I think that my experiment and answers are fairly accurate because the legal minimum is 4% by mass. In determining the molecular mass of an acid, error may have occurred during the titration process, from drops that could not be controlled accurately using the buret. In table 5, it can be seen that the volumes of NaOH titrated varied slightly. Although I could see no other major sources of error, I am doubtful about my result of an average molecular mass of 215.93g/mol is correct, because it is so large, I am unsure of what monoprotic acid it could be. Acid/base reactions are important because they are present in both chemical and biochemical reactions in everyday life. These include heat from neutralization, qualitative properties from gas liberating reactions, and oxidation-reduction reactions when reacting with certain metals. Many of these reactions can be found in household cleaners, and it is important to know what chemicals should be mixed or not. Indicators are also very useful because they are often used to test the pH of an environmental solution that living organisms are in, like fish. There are many types of titrations other than the acid-base titration that was performed in this experiment, but the basic goal overall is to determine the concentration of an unknown substance using one that is known.
Conclusion:
Overall, the results of my experiment coincided with acid/base theories as well as with my expectations. It was predicted that adding a strong acid to a strong base will result in a neutralization reaction with a lot of heat given off, because the temperature of the small test tube increased by four degrees, it proves that this theory is true. When the indicators were tested in various pH¡¦s, they all turned to their predicted colors, which shows that the experiment was performed accurately. The standardization of NaOH resulted in a concentration of 0.2439M NaOH averaged from three trials. However, I think if I were to perform more titration trials, this average might change slightly. With the standardized NaOH I found that the unknown #3 had an average of 5.68% by mass acetic acid, which is not under the legal limit, nor it is over, so it seems to be within the correct range. Lastly, my molecular mass for monoprotic unknown #1 was 215.93g/mol which I did not expect, since it is so large, the correct molar mass was not provided, so I am unable to deviate my results from the true value.
Experimental Overview:
The procedure for this experiment was carried out as instructed in the laboratory manual, Experiments in General Chemistry, 4th ed., S.L. Murov, Experiment 21, Acids and Bases: Reactions and Standardizations, and Experiment 22, Acids and Bases: Analysis. There were modifications made by the instructor to dilute the 6M NaOH to 0.1M in 300mls instead of 500mls in Part B of Experiment 21.
One objective for performing these two experiments was to observe qualitatively the reactions between common acids, bases, and the indicators phenolphthalein, methyl orange, bromothymol blue, and red cabbage extract. Another focus of this experiment was to perform a …show more content…
standardization of sodium hydroxide solution using the titration technique with potassium hydrogen phthalate. By titrating the standardized NaOH, determine the percent by mass of acetic acid in an unknown vinegar solution as well as determine the molecular mass of an unknown acid.
Background Information:
When an acid such as HCl and base such as NaOH are mixed with each other they neutralize and produce salts and react as follows:
HCl + NaOH ³ NaCl + H2O + Heat
Heat can be quantitatively measured using a thermometer as a result of the reaction forming water. Other times, combining an acid such as HCl with a base such as Na2CO3 will result in the following:
2HCl(aq) + Na2CO3 (aq)³ 2 NaCl + CO2 (g) + H2O (l)
Another unique property of acids and bases are their high reactivity with certain metals. Often their reactions will often be violent.
NaOH + Al(s) + ³ AlO (s) + H2 (g) + Na+ (aq)
Chemical indicators of acids or bases work because their molecules change shape in the presence of ions. The following indicators were used in this lab:
Methyl Orange is often used as an indicator for strong acids. In the presence of H+ ions the indicator turns a bright red, but in a pH of about 4 it is orange, and in more basic solutions it turns yellow. Bromothymol blue is a more diverse indicator because it turns yellow below pH 6, green at neutral 7, and blue at 8 and above. Red cabbage juice also turns different colors in different pH. It turns red in an acid and green in a base.
Phenolphthalein is an indicator for bases that turns from clear to pink in the presence of OH- ions. This is due to fact that phenolphthalein molecules are clear, but when they ionize in solution, their H+ ions that make them a weak acid, are pink.
The titration procedure is a quantitative analysis which is used to determine the concentration of a known reactant. A calibrated burette is used to add a reagent to determine the exact amount. The titration is stopped when the solution slightly changes due to the indicator. Indicators such as phenolphthalein are useful in determining the molarities of bases such as NaOH. Through titration, NaOH is carefully added into a specific amount of diluted KHP (potassium hydrogen phthalate 204.22 g/mol) with a few drops of phenolphthalein. When the system can no longer equilibrate the NaOH being added to it because all the KHP, will be used up, there will be an excess of NaOH which the phenolphthalein will indicate by turning pink. By knowing the amount of NaOH titrated, we can calculate the average molarity of OH- reacted because it will react in a one to one ratio with the monoprotic KHP. Titration is a necessary method because it is inaccurate to calculate the amount of OH- ions using the molarity of NaOH, because each dilution of the solution contains a different amount of ions. This process of determining the concentration of a solution is called standardization. Vinegar is produced by the fermentation of sugar to ethanol and then bacteria catalyze the oxidation of the ethanol to acetic acid. The legal minimum is 4% by mass acetic acid in vinegar. Using the standardized NaOH, it is also possible to determine the percent by mass of acetic acid contained in vinegar as well as the molar mass of and unknown monoprotic acid. Like KHP, acetic acid, HC2H3O2, is an example of a monoprotic acid because in solution, it only releases one H+ ion, becoming H+(aq) and C2H3O2-(aq).
Results and Calculations:
Table 1 summarizes the experimental results of experiment 21 part A. 1 through 4 and includes the reagents added for each step. The reagents which were added are listed on the left, as well as their balanced equations. Observations are displayed on the right hand side. This table summarizes the common reactions of acids and bases as well as the reaction time for an indicator to work.
Table 1. Results of Reactions with Acids and Bases.
Reagents added Observations
2mls of 3M HCl + 2mls of 3M NaOH HCl + NaOH ³ NaCl + H2O + Heat The initial temperature of the 3M HCl was measured at 24 C . When 3M NaOH was added and stirred, the temperature was found to be 28 C.
1mL of 3M HCl + 1mL of 1M Na2CO3 in a large test tube
2HCl + Na2CO3 ³2 NaCl + CO2 + H2O Slight fizzing when the solutions were combined. When a lighted splint is inserted into the tube, the flame get smaller and goes out.
3mLs of 3M NaOH + wad of aluminum foil
2NaOH + 2Al + ³ 2AlO (s) + H2(g) + 2Na+ (aq) A lot of loud fizzing, gas can be seen forming. When lighted splint is inserted, the flame immediately is sucked in and a loud pop noise is made.
Solutions A+B
Solutions A+C It took 13 seconds for the solution to turn pink
It took 32 seconds for the solution to turn pink
Table 2 shows the color change of four indicators when they are added to three solutions with different pH¡¦s . The original colors of the indicators are included on the left.
Table 2.
Color Chart of Indicators in 0.01M HCl, pH 7 Buffer, and 0.01 M NaOH
0.01M HCl pH 7 Buffer 0.01 NaOH
Phenolphthalein
(clear) Cloudy precipitate No change Pink
Methyl orange
(yellow-orange) Reddish pink No change Yellow
Bromothymol blue
(green) No change No change Blue
Red-cabbage extract (purple) Pink No change Green
Table 3 includes data collected throughout the process of standardizing the sodium hydroxide solution. Information from all three titrations of sodium hydroxide into KHP is recorded in order to calculate the average molarity of the NaOH solution in the end. Calculations used to derive the numbers in the table will be shown after the table.
Table 3. Data collected from three titrations of 0.1M NaOH into 50mls KHP with 3 drops of phenolpthalein.
Trial 1 Trial 2 Trial 3 Average
Mass of KHP 1.001g 1.002g 1.001g
Moles of KHP 0.00490mols .00491mols .00490mols
Initial Buret Reading 30.0mL 30.0mL 30.0mL
Final Buret Reading 9.8 mL 10.0mL 10.0mL
Volume of NaOH Solution 20.2mL 20.0mL 20.0,L
Molarity of NaOH solution 0.243M 0.246M 0.243M 0.2439M
Deviation From average 0.0009 0.0016 0.0009
.00113
The moles of KHP were determined as follows:
Mass of KHP x Molar Mass of KHP = number of moles KHP
From trial 1:
The moles of KHP and the moles of NaOH are equivalent therefore the molarity of NaOH is found by dividing the number of moles NaOH (same as moles of KHP) by the volume of NaOH titrated. This calculation was repeated for all three trials.
Molarities of NaOH:
From Trial 1:
The average molarity from all three trials is then calculated as 0.2439M.
Further analysis is done by determining each trial¡¦s deviation from the average.
Average Molarity ¡V Experimental Molarity = Deviation from Average
From Trial 1:
Table 4 summarizes the data collected from three titrations of the standardized NaOH into vinegar with an unknown concentration of acetic acid (unknown #1).The buret readings indicate the volume of NaOH titrated. In the end, the average mass percent of acetic acid in the unknown vinegar was calculated. Calculations are shown below.
Table 4. Data from three titrations of standardized NaOH into unknown vinegar #3.
Trial 1 Trial 2 Trial 3 Average
Mass of Flask 98.3414g 96.8691g 74.8863g
Mass of Flask + solution 101.0999g 101.2024g 79.6695g
Mass of solution 2.7585g 4.3333g 4.7632g
Density of vinegar 0.5517g/ml 0.8666g/ml 0.9566g/ml
Molarity of NaOH used 0.2439M 0.2439M 0.2439M
Initial Buret Reading 30.0mL 30.0mL 30.0mL
Final Buret Reading 15.7mL 15.2mL 15.5mL
Volume of NaOH solution 14.3mL 14.8mL 14.5mL
Moles of NaOH 0.00349mols 0.00361mols 0.00354mols
Moles of Acetic Acid 0.00349mols 0.00361mols 0.00354mols
Molarity of Acetic Acid 0.698M 0.722M 0.708M 0.709M
Deviation of each molarity from average 0.011 0.013 0.001M 0.0083
Mass Percent of Acetic Acid in Vinegar 7.59% 5.00% 4.44% 5.68%
From Trial 1: Since acetic acid is monoprotic, it also reacts in a 1:1 ratio with NaOH, so the number of moles of NaOH was equated to the moles of acetic acid.
From Trial 1:
The mass percent of acetic acid in vinegar was calculated as follows:
From trial 1:
Table 5 includes data collected from three titrations of NaOH into unknown acid #1. For each trial, a molecular mass is determined and then averaged in the last column. Calculations are shown below.
Table 5. Data from three titrations of standardized .2439M NaOH into an unknown acid #1 with phenolphthalein indicator.
Trial 1 Trial 2 Trial 3 Average
Mass of Unknown 0.5023g 0.5021g 0.5082g
Initial Buret Reading 30.0mL 30.0mL 30.0mL
Final Buret Reading 20.5mL 20.4mL 20.0mL
Volume of NaOH titrated 9.50mL 9.57mL 10mL
Moles of unknown acid (assuming monoprotic) 0.00232mols 0.00233mols 0.00233mols
Molecular Mass of acid 216.50g/mol 215.49g/mol 215.79g/mol 215.93g/mol
Deviation of MM from avg 0.57 0.44 0.14 0.38
The moles of the unknown acid were calculated using the molarity of the standardized NaOH multiplied by the liters of NaOH titrated into the unknown acid, which is equal to the moles of monoprotic acid reacted.
For trial 1:
Since molecular mass is in units of g/mol, the mass of the unknown acid is divided by the moles of acid calculated above.
For trial 1:
Discussion
Three experiments involving reactions between acids and bases were done to test and observe their unique properties. These included neutralization, gas evolutions, as well as a basic redox reaction. Various common indicators were also tested out by combining them with acids, bases, and a neutral solution as well to witness how they work. The standardization of NaOH was also done for use in determining the percent of acetic acid in vinegar and the molecular mass of an unknown acid.
When HCl is added to NaOH, the result was a rise in temperature by 4 degrees because when water is formed during neutralization, heat is also released as a product therefore raising the temperature of the solution. The results of my experiment proved that this is a true neutralization reaction. When 1mL of 3M HCl was added to 1M Na2CO3 and a lighted splint was inserted, the flame went out because of the CO2 being liberated from the reaction, much like how a fire extinguisher works. However, in the following experiment where HCl was combined with Na2CO3, the flame that was inserted immediately got pulled into the test tube making a loud pop because of the vacuum being formed. This occurred because hydrogen gas was being released from the reaction, which is flammable. My results agreed with what happens when a flame is inserted into a flammable gas and a nonflammable gas.
When NaOH and solid aluminum were combined, the reaction was so violent because aluminum has a high reactivity with acids and bases. The aluminum becomes oxidized as it goes from a 0 to a +2 oxidation state to form solid aluminum oxide, which is the dark precipitate that formed. The chemical equation 2NaOH (aq)+ 2Al(s) + ³ 2AlO (s) + H2(g) + 2Na (aq) corresponds with the results of my experiment.
During the mixture of the two solutions A+B and A+C, the method used for timing was too arbitrary, because I had to keep glancing up at the clock, whose second hand was not in direct view to me which resulted in error in the time it took to react. I am unsure why a precipitate formed when phenolphthalein was combined with HCl, unless this is just how phenolphthalein reacts in an acid, but to be sure I would test it using different acids. All of the indicators remained the same color in the pH 7 buffer because the buffer solution is neutral, because the indicators only react when there are ions present. Bromothymol blue was the only indicator which didn¡¦t change color in acid because it is unreactive with the H+ ions.
Three titrations of NaOH were used to determine an average concentration to ensure a more accurate number since each batch of NaOH collected from the stock solution will always have different concentrations of OH-, in order to help reduce error. Practice using the buret was also helpful in greatly reducing error during titration, because I found that adding even the slightest drip would taint the results and turn the solution dark pink. By starting the titration at 30.0mLs in each buret for all trials, I think I also reduced a source of error. The molarity of my NaOH solution turned out to be 0.2439M. It is shown in Table 3 that trial 2 had the greatest deviation from the average, which may have been caused by a larger mass of KHP, 1.002g initially, compared to 1.001g. There may have been an error during the weighing of the flasks in Table 4 for trial 3 because the weight of flask 3 was so much lower than the other two even though I used the same type of flask. According to the averages of my three trials, unknown vinegar #3 had 0.709mols of acetic acid per liter. However, this number was not used in calculating the average mass percent of acetic acid in vinegar, because I thought it would be more accurate to calculate for each trial and then take the average, which is 5.68%. Since it is a small percentage, I think that my experiment and answers are fairly accurate because the legal minimum is 4% by mass. In determining the molecular mass of an acid, error may have occurred during the titration process, from drops that could not be controlled accurately using the buret. In table 5, it can be seen that the volumes of NaOH titrated varied slightly. Although I could see no other major sources of error, I am doubtful about my result of an average molecular mass of 215.93g/mol is correct, because it is so large, I am unsure of what monoprotic acid it could be. Acid/base reactions are important because they are present in both chemical and biochemical reactions in everyday life. These include heat from neutralization, qualitative properties from gas liberating reactions, and oxidation-reduction reactions when reacting with certain metals. Many of these reactions can be found in household cleaners, and it is important to know what chemicals should be mixed or not. Indicators are also very useful because they are often used to test the pH of an environmental solution that living organisms are in, like fish. There are many types of titrations other than the acid-base titration that was performed in this experiment, but the basic goal overall is to determine the concentration of an unknown substance using one that is known.
Conclusion:
Overall, the results of my experiment coincided with acid/base theories as well as with my expectations. It was predicted that adding a strong acid to a strong base will result in a neutralization reaction with a lot of heat given off, because the temperature of the small test tube increased by four degrees, it proves that this theory is true. When the indicators were tested in various pH¡¦s, they all turned to their predicted colors, which shows that the experiment was performed accurately. The standardization of NaOH resulted in a concentration of 0.2439M NaOH averaged from three trials. However, I think if I were to perform more titration trials, this average might change slightly. With the standardized NaOH I found that the unknown #3 had an average of 5.68% by mass acetic acid, which is not under the legal limit, nor it is over, so it seems to be within the correct range. Lastly, my molecular mass for monoprotic unknown #1 was 215.93g/mol which I did not expect, since it is so large, the correct molar mass was not provided, so I am unable to deviate my results from the true value.