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The intent of this lab was to quantify the alkalinity of different water samples based upon the amount of sulfuric acid which was added to lower the pH to certain control points. This is a very important concept with respect to water treatment because the higher the alkalinity of a water sample, the greater ability that sample has to accept the addition of acids. Understanding these principles is crucial to being able to forecast whether natural water sources like rivers and streams will still be safe for use following weather events. By tracking the alkalinity of a water source consistently, environmental engineers are capable of predicting when water samples will no longer be safe for human use due to the addition of acids.


There are various reasons why one may choose to test alkalinity of a sample using either indicators or a pH meter. One advantage of the pH meter is that once calibrated it is very easy to use and can perform multiple trials in a short time. However, issues in the calibration of the machine and a lack of cleansing of the device can result in inaccurate results that may compound and become more inaccurate as trials progress. For this reason, the use of indicators can be beneficial as with adequate time, obvious color change can be quickly observed and accurate measurements may be taken. A limitation of the indicators however is that indicators themselves are weak acids so they may under report the actual acid neutralizing capacity of a water sample. Indicators also depend on visual confirmation that the pH has reached below the indicator’s specific threshold. Requiring visual confirmation to require an accurate pH reading allows for little room for error as the sample continues to mix with the titrant resulting in a possibility the pH reading will be more acidic than the required threshold.

There appears to be roughly two inflection points along the Solution A curve at roughly pHs of 8.2, and 4.4. These values correspond closely to the expected pH values of 8.3 and 4.5 leaning toward a slightly more acidic pH value. This systematic error of a higher acidic concentration could be the result of the titrations consistently using more acid to reach the expected values due the qualitative data, color shift, as opposed to quantitative data, pH value, to determine the threshold of indication. The inflection points determined along the titration curve correspond well to the observation that natural waters have a pH that fall between 6 and 8 when accounting for the systematic error observed. It is expected that natural lakes and streams would have a pH that falls within the range of 6 to 8 because it is the range that requires a higher concentration of acid to affect the pH to greater effect. These streams and lakes would be able to maintain pH values between 6 and 8 due to the large volume of water within the system relative to the volume of acid being introduced from runoff and other sources.

Overall, it was observed that the pH meter allowed for quicker and slightly more accurate testing as less titrant was typically required to produce pH readings closer to desired values. This revealed a systematic error as typically pH readings using indicators resulted in values further from the desired values than pH meter readings. Additionally, the titration curve revealed inflection points near the desired pH values of 8.3 and 4.5 which helps explains why natural river sources maintain a pH between 6 and 8 which is tolerant to the addition of acids. This proves that testing with a pH meter for alkalinity provides more accurate data to better predict the ability of a water source to withstand the addition of acids