6.03 Calorimetry Lab Report

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B. Buffers.
1. Two solutions, A. being 1.0 x HCl, and B. being 0.10 M will be given, as well as buffer solution 1 M NaOH.
2. Have 25 mL of solution A. and B. in two separate beakers.
3. Check the pH of both solutions without any NaOH first. Then begin checking pH after 1 drop, then after 10 drops, and lastly 25 drops. After each set of drops be sure to mix solutions.
4. Compare both sets of pH to determine which had an easier time resisting change.

C. Titration curves.
1. Insert NaOH into burette, proceed to clamp burette to stand and insert into beaker containing 10 mL of HCl.
2. Add drops of NaOH into the beaker until a change in pH by 0.50 units is seen or you have added 1.0 mL of base. Be sure to go at a pace which is good for catching correct equivalence point.
3. Record both the new pH each time as well as the
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pH of acids, bases, and salts As predicted, going down the list of solutions the pH of each got higher and higher, so it went from acidic to basic. There is even a difference in pH between the HCl solutions due to the difference in molarity—even being the same solution, the less concentrated it is, the more basic it becomes. Comparing between how the pH paper did versus how the pH meter did, they were both similar at times with each other, and others having more than .50 difference; overall ranging in difference between .04 to 1.74. From the beginning there were some technical difficulties with the pH meter, going through some that did not work, to some that were just completely off, so the other meters being precise was not definite considering the one that ended up being used was not very accurate for even deionized water. All pH meters may not have those kind of issues if based on age, quality, etc., but concerning the ones that were being used, the pH paper may have been the more accurate route if not contaminated. So between the two, the pH paper technique in this situation seemed to have less chance of error than the pH meter.


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