The standardized concentration of NaOH was used to identify the unknown acids, this value is shown is experiment 16.1. This mean standardized concentration was 0.0784 M. By titrating a known amount and concentration of this standardized base to the unknown acid, a titration curve and its derivative could be plotted in order to find the molar mass and acid dissociation constant.
The first unknown acid titrated, BGYW, was identified to be Maleic acid, C4H4O4. The identity of the acid was found by taking both the molar mass and pKa into account. However, one pKa value had a very high percent error. The molar mass is taken more into consideration than the pKa value as elaborated on in number 4. The shape of the curve shown in Figure 1 also corresponds
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Table 16-2-1 in the lab manual was used to compare the pKa and molar mass of the unknown acid to determine the identity. C4H4O4 was the only diprotic acid with comparable molar mass and pKa values with a 31% error for the molar mass and a 12% error for the second pKa value. However, the first pKA value comparison had a 157% error which could be due to experimental error (as shown in Table 1). Increments of NaOH added to reach the endpoint being too large could have also affected the molar mass difference between the unknown and the actual reference value. Since the increments were too large, the equivalence point in Figure 1 is slightly off since it was overshot. Therefore, the initial moles of acid and base were expected to be of lower value than those referenced. The dV uncertainty shown in the sample calculations page +/- 0.44 mL displayed this source of error between the calculations of moles of acid, the half equivalence point and the pKa; accounting partially for the percent error in the molar mass and pKa of the unknown acid versus Maleic (C4H4O4).
Both Figures 3 and 4 were used to identify unknown acid *B4QK as 3,5 Dimethoxybenzoic (C9H10O4). The titration curve only had one inflection point, thus one equivalence point as demonstrated in Figure 3; showing that the acid was monoprotic. The
The purpose of this experiment is to determine an unknown concentration of acid (hydrochloric acid) with a standard solution of a base (sodium carbonate) using titration method.
The purpose of this semester long experiment was to determine an unknown organic acid. An organic acid is an organic compound with acidic properties. A base reacts with acids to form salts. Titrations are used to determine the concentration of unknown substances. The purpose of the KHP experiment was to determine the molarity of NaOH. HCl titrations are mainly to check technique and used to verify the molarity of NaOH solution. The hypothesis is that this acid is C4H3OCOOH.
The purpose of this experiment was to determine the pKa, Ka, and molar mass of an unknown acid (#14). The pKa was found to be 3.88, the Ka was found to be 1.318 x 10 -4, and the molar mass was found to be 171.9 g/mol.
Table 2: Consists of color extract taken from a red cabbage for a natural indicator. The pH reading that was measured by using the pH meter and the result of the pH reading to determine whether the solution was acidic or basic.
Chemistry 102 is the study of kinetics – equilibrium constant. When it comes to the study of acid-base, equilibrium constant plays an important role that tells how much of the H+ ion will be released into the solution. In this lab, the method of titrimetry was performed to determine the equivalent mass and dissociation constant of an unknown weak monoprotic acid. For a monoprotic acid, it is known that pH = pKa + log (Base/Acid). When a solution has the same amount of conjugate base and bronsted lowry acid, log (Base/Acid) = 0 and pH = pKa. By recording the pH value throughout the titration process and determining the pH at half- equivalence point, the value of Ka can be easily calculated. In this experiment, the standardized NaOH solution has a concentration of 0.09834 M. The satisfactory sample size of known B was 0.2117 g. The average equivalent mass of the unknown sample was found to be 85.01 g, pKa was found to be 4.69, which was also its pH at half-equivalence point and Ka was found to be 2.0439×〖10〗^(-5). The error was 1.255% for equivalent mass and 0.11% for Ka. In other word, the experiment was very precise and accurate; the identity of the unknown sample was determined to be trans-crotonic by the method of titrimetry.
1. First, 3.0015 g of salicylic acid were measured using an electronic balance. 2. The salicylic acid was then transferred into a 150 mL beaker.
Identifying this organic acid was an extensive task that involved several different experiments. Firstly, the melting point had to be determined. Since melting point can be determined to an almost exact degree, finding a close melting point of the specific unknown can accurately point to the identification of the acid. In this case the best melting point
Looking at the table in the lab book, the acid with the closest pKa to my calculated pKa is trans-crotonic. It has a pKa of 4.69 while my calculated pKa was 4.62. I feel that these numbers are very close, as the percent error is only 1.49%. The percent error for my calculated equivalent mass of the acid is also small at 3.11%. Error could have come from reading the graph, as I had to estimate the equivalence point. I should have taken measurements at smaller intervals when approaching the inflection point. This would have allowed me to read the graph more accurately. Error could have also come from the pH meter, which may not have been calibrated precisely. Lastly, error could have resulted if I inaccurately read the buret while recording the
Observe and measure a weak acid neutralization and determine the unknown identity of an unknown acid by titration.
The guiding question of this ADI lab was, “What are the identities of the unknown compounds?” The goal of this lab was to understand the relationships between moles and molar mass to find the identity of unknown compounds. The mole can be used to measure small amounts of a substance or is used to convert from unit to unit using dimensional analysis. One mole is equivalent to the molar mass in grams of that substance. If you start with the moles of an unknown substance, multiply it by a given compound’s molar mass, and then divide it by however many moles are in the compound of your choice, you will get the mass of the compound. With that answer you can then compare with mass of the compound in the bag to determine its identity. We first started
One milliliter of 6.00-M phosphoric acid was placed into a 125-mL Erlenmeyer flask using a volumetric pipette. Using a slightly larger pipette, six milliliters of 3.00-M sodium hydroxide was transferred into a 50-mL beaker. Then a disposable pipette was used to slowly mix the sodium hydroxide into the phosphoric acid while the solution was swirled around. Then both the beaker and flask were rinsed with 2-mL of deionized water and set aside. A clean and dry evaporating dish was weighed with watch glass on a scale. Then the solution was poured into the dish and the watch glass was placed on top. The solution was then heated with a Bunsen burner to allow for the water to boil off to reveal a dry white solid. After the dish cooled to room temperature it was once again weighed and the new mass was recorded.
By using acid-base titration, we determined the suitability of phenolphthalein and methyl red as acid base indicators. We found that the equivalence point of the titration of hydrochloric acid with sodium hydroxide was not within the ph range of phenolphthalein's color range. The titration of acetic acid with sodium hydroxide resulted in an equivalence point out of the range of methyl red. And the titration of ammonia with hydrochloric acid had an equivalence point that was also out of the range of phenolphthalein.. The methyl red indicator and the phenolphthalein indicator were unsuitable because their pH ranges for their color changes did not cover the equivalence points of the trials in which they were used. However, the
First, three titration curves and three second derivative curves were created to determine the average pH at the half-equivalence point from the acetic acid titrations. Titration curves were used as visuals to portray buffer capacity. The graphs and a table, Table 1, that showcased the values collected were created and included below. The flat region, the middle part, of Figures 1, 2 and 3, showed the zone at which the addition of a base or acid did not cause changes in pH. Once surpassed, the pH increased rapidly when a small amount of base, NaOH, was added to the buffer solution. Using the figures below and
To keep this a fair test I have made sure that the acid used is the
An acid-base titration is the determination of the concentration of an acid or base by exactly neutralizing the acid/base with an acid or base of known concentration. This allows for quantitative analysis of the concentration of an unknown acid