Titration Of Hcl Lab Report

Buffers can be used to resist change in pH in an acid or base solution. The purpose of the lab is to understand the properties of buffers in relative concentrations, in the presence of a strong acid/base and dilution of buffer components the ammonia/ammonium buffer system. In this case, these laboratory techniques associated with buffers are performed to determine the pH values in diluted solutions. This is done by observing the Henderson-Hasselbalch equation. This represents the pH value by comparing the p value and the concentrations of a weak acid and conjugate base of a buffer solution, as long as assumptions are valid. From there, a pH vs log can be used from the values obtained. Then, the theoretical pH of solutions can be determined with the addition of a strong acid (HCl) and base (NaOH). From this, the effect of dilution can be viewed when plotting the calculated pH values vs dilution on the same graph.

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 makeup of a compound drastically changes its intermolecular forces (IMFs) with the polar silica gel of a TLC plate; this concept is responsible for the variability of Rf values observed throughout the course of lab. The weaker the IMFs, the further a compound will travel through the silica. For instance, ionic interactions are the strongest IMF, but were not present during this experiment. H-bonding IMFs had the greatest impact for our specific compounds. Resorcinol and 3-chlorobenzoic acid exhibited low Rf values, due to its ability to H-Bond to silica’s hydrogen donors and oxygens acceptors (See graph 1). Both were adept to strongly H-bonded to silica because the compounds contained at least 1 H-donor and 2 H-acceptors. Thus, stronger/more

pH was recorded every time 1.00 mL of NaOH was added to beaker. When the amount of NaOH added to the beaker was about 5.00 mL away from the expected end point, NaOH was added very slowly. Approximately 0.20 mL of NaOH was added until the pH made a jump. The pH was recorded until it reached ~12. This was repeated two more times. The pKa of each trial are determined using the graphs made on excel.

We only added a small amount of HCl to the water and sodium chloride. We did not continue to add more HCl after a significant drop in pH was recorded. We added a total of 2 mL of HCl to both H20 and NaCl before the pH changed. The 1 gram solution of sodium acetate and acetic acid changed after a 8 mL, and the other two never dropped before we reached our total of 10 mL HCl.

When using different methods to measure pH levels there are some tools that can be useful. Some more than others but by putting into action the different methods it may determine which tools will work best and give the best results when testing the pH within a solution. The pH, which stands for the proportion of hydrogen ions in a solution, could be acidic (acidosis), neutral or basic (alkaline). The pH scale goes from numbers 1 through 14. A pH of 7 is neutral;

Methods: First, a burette, ring stand, clamp, and an empty flask were obtained. The burette, with the valve closed, was attached to the ring stand with a clamp, and the empty flask was placed below the burette. Next, 50mL of the NaOH solution were poured into the burette, and a small bit was drained into the empty flask to ensure that the tip of the burette was also full of NaOH solution. The volume of the NaOH in the burette was recorded. Next, approximately 0.6 grams of KHP were massed poured into an empty 125mL flask. Two drops of an indicator solution were added to the KHP

ii. The second part of the titration series involves titration of NaOH with Hydrochloric acid (HCL). Again, three reps of titration and a blank titration have to be completed. A volumetric pipet is used to measure 10.00mL of HCL into three labeled conical flasks. Then the flasks are filled with deionized water until about the 50mL mark. A buret is

The purpose of the experiment was to determine how a buffer works and how to use an acid-base indicator. The way a buffer works was determined by observing the changes in pH of solutions of different concentrations weak acids and their conjugate bases to determine how a buffer affects the pH change. The solution of 10 mL of 0.20 M CH3COOH and 10 mL of 0.20 M CH3COONa had slighter changes in pH than the solution of 10 mL of 0.0020 M CH3COOH and 10 mL of 0.0020 M CH3COONa. Both of these solutions were buffers, shown because they had slighter changes in pH than the solutions with only the weak acid or conjugate base and water. The determination of how buffers work was also tested with observing that the solution of NaC4H3O4 and Na2C4H2O4 had smaller

Its concentration was found through a quantitative and qualitative titration testing. Because HCl is acidic, a 1.07M NaOH solution was used to titrate an aqueous solution of HCl (the unknown solution). The PASCO program gave the final results shown in Figure 1. A typical titration curve can be described as starting at the initial pH of the solution, then some acid/base is added, until a sharp curve either up or down takes the pH past a neutral of 7, and towards the pH level of whatever the added acid’s/base’s pH is (Clark6). The stronger the acid or base being titrated, the faster the change occurs, which means a sharper slope of the titration graph. Also, it is important to note the equivalence point on this titration graph. At the equivalence point “the correct amount of standard solution must be added to fully react with the unknown concentration” (Xavier3). In Figure 1, this equivalence point is shown by where the sharpest slope occurs in the curve, which is seen at about a pH of 5.0

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

The purpose of this lab was to use process titration to find concentration of an aqueous solution of Hcl(aq) , using KOH(aq) as the titrant.

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

For this experiment, a pH meter was used so this part of the experiment began with the calibration of the pH meter with specified buffers. The buret was then filled with the standard HCl solution and a set-up for titration was prepared. 200g of the carbonate-bicarbonate solid sample was weighed and dissolved in 100 mL of distilled water. The sample solution was then transferred into a 250-ml volumetric flask and was diluted to the 250-mL mark. The flask was inverted several times for uniform mixing. A 50-mL aliquot of the sample solution was measured and placed unto a beaker. 3 drops of the phenolphthalein indicator was added to the solution in the beaker. The electrode of the pH meter was then immersed in the beaker and the solution containing the carbonate-bicarbonate mixture was titrated with the standard HCl solution to the phenolphthalein endpoint. Readings of the pH were taken at an interval of 0.5 mL addition of the titrant. After the first endpoint is obtained, 3 drops of the methyl orange was added to the same solution and was titrated with the standard acid until the formation of an orange-colored solution. Readings of the pH were also taken at 0.5 mL addition of the titrant.