Investigating The Ph Of Water Sources Essay

1. To titrate a hydrochloric acid solution of “unknown” concentration with standardized 0.5M sodium hydroxide.

H. How would you prepare 10 mL of a 0.25M HCl solution if 1M HCl was available? How much

The Distilled water pH average of HCl for zero drops was 6.67, and the pH for the final thirtieth drop

In 2 and 7 I added 50 mL of .1 M NaCl. I added sodium acetate to the rest of the beakers: 1 gram to 3 and 8, 5 grams to 4 and 9, and 10 grams to 5 and 10. I then filled the beakers that contained the solid sodium acetate with 50 ml of .10 M acetic acid. Specifics can be found on page 84 of the lab manual. Though the lab manual instructed to use a pipet, we did not have an accurate 1 mL pipet or a graduated pipet, so we instead prepared two graduated burets with 1 M Sodium Hydroxide and 1 M hydrochloric acid. Using a standardized pH probe with a Lab Pro to measure changes in pH, we added 1 mL of HCl at a time and recorded the changes. The same was done for the NaOH.

To improve the results from the experiment buffer solutions that were not whole pHs could have been used e.g. pH 4.5, 5.5 etc. This would have provided more reliable results as a wider range of results would have been produced. Using pHs with decimals would also help to more accurately determine the optimum pH as the optimum may have been above or below the pH stated in the hypothesis; 8. In this experiment however the optimum is taken at 8 because the graph does not rise again.

To make the buffer solution you need 0.2 mol dm-3 of Na2HPO4 and 0.1mol dm-3 of citric acid this will give 100cm3 of buffer. Here is how to get the different pH in the buffer solutions:

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

All around the world, countries are fighting to keep their drinking water clean. Whether it’s streams, rivers, or lakes, countries have taken great measures to maintain high quality drinking water for both human consumption and animal consumption. Countries must first understand the sources of the polltion, then determine the best methods to eliminate the pollution. Clean drinking water is a valuable resource and a the key to human survival. Plants and animals also depend on water for their growth, so all water must be kept clean. The major contributors to water pollution can be classified in three categories, industrial, agricultural and municipal.

In this demonstration, water was mixed with phenolphthalein, a pH indicator, in a test tube and topped with gasoline; the gasoline did not mix with the water due to opposite polarities. The layer of gasoline stayed above the water because the density gasoline is that of water’s. A small piece of metal sodium was then placed in the test tube and the reaction began to occur once the sodium came into contact with the layer of water. The reaction was visible through the pH indicator which changed color from clear to a vibrant pink. This change in color proved that the sodium altered the pH balance of the water and phenolphthalein mixture. The sodium and water also reacted to create hydrogen gas which produced bubbles in the test tube. These hydrogen

pH is usually to measure acidic or basic water is. pH is a measure of the relative amount of free hydrogen and hydroxyl ions in water. There are 0-14 range to be measure which is 7 is for neutral while pH which less that 7 is acidity and the pH which is greater than 7 is alkali. Water that more free hydrogen ions is acidic while water has more free hydroxyl ions in water is basic. Chemicals in the water can affected the pH in water. Ph show an important indicator for water if it chemically changes. pH is showed in logarithmic units. The change acidity or alkali in water represent in 10-fold. pH of 5 is ten times more acidic in water than pH of 6.

Sample one pH reading was 8.14, sample two pH reading was 7.72, sample 3 pH reading was 8.23, sample 4 was 8.30, sample 5 was 8.22, sample 6 was 7.76, sample 7 was 8.52, sample 8 was 8.22 and sample 9 pH reading was 7.73. Possible reasons may be that the additional rain and runoff had an effect at certain locations to slighting lower the pH. Looking at the aerial view of where samples were taken sample two and six appear to possible get a concentration of rainwater from the corners of the near by buildings roofs, this could account for a lower pH. Speculation as too sample nine’s lower pH, I recall seeing many trees having new mulch around them, if they were also fertilized this could account for a bit lower pH. However this is only speculation and assumptions on my part.

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

2. Following solutions are added to the tubes and the pH of each tube is determined:

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.

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.