Solution 4 like Solution 2 is a buffer solution comprised of a weak base and

Solution 4 like solution 2 is a buffer solution

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colored solution. Solution 4, like Solution 2, is a buffer solution comprised of a weak base and its soluble ionic salt. Solutions containing weak bases and their ionic soluble salts are always less basic than solutions with weak bases alone. By that fact, a pH value lower than that of solution 3 is expected. A modified version of the Henderson- Hasselbalch equation (3) suitable for a solution comprised of a weak base is to be used to find the pH of this solution, pOH = p K b + log ( cojugateacid base ) (6) Using both equation (6) and equation (5), the calculated ideal pH value of the solution is 9.25. The solution is expected to have a light violet color. Experimentally, the solution also had a light violet color which verifies with theoretical expectations. Table 2. Effect of Strong Acid and Strong Base on Buffers. Solution Estimated pH range pH Reading + Methyl Orange + Phenolphthalein pH meter Calculated 1 a pH < 3.1 - 2.35 2.30 b pH > 4.5 - 3.87 3.47 c 3.1 < pH < 4.5 - 3.46 2.87 a pH < 4.5 - 4.40 4.70
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2 b pH > 4.5 - 4.70 4.80 c pH > 4.5 - 4.58 4.75 3 a - pH < 10.00 9.89 10.53 b - pH > 10.00 11.56 11.70 c - pH > 10.00 10.46 11.13 4 a - 8.3 < pH < 10.00 8.97 9.23 b - 8.3 < pH < 10.00 9.28 9.30 c - 8.3 < pH < 10.00 9.07 9.25 c is the control In the second part of the experiment, for solutions 1 and 3, it is expected that the pH would drastically change (rise or fall) from the addition of drops of strong acids and bases since they are not buffer solutions. They are pure solutions of weak acids and bases without much capacity to counteract additional stresses, unlike solutions 2 and 4 which are buffer solutions that are expected to negate or minimize pH changes. In solution 1a, the addition of HCl follows a special case of Le Chatelier’s principle, the Common-ion effect. This is when a weak electrolyte’s dissociation is suppressed by the addition of one of its ions from another source [5]. In this solution, the common ion is H 3 O + . Suppression would cause the formation of hydronium ions from the acid source, CH 3 COOH, to be negligible. Because of this, majority of the hydronium ions would come from the addition of 1.0 M HCl, CH 3 COOH + HCl → CH 3 COO - + H 3 O + (7) For solution 1b, a neutralization reaction occurred given by the following chemical equation, CH 3 COOH + NaOH → NaCH 3 COO + H 2 O (8) The neutralization would cause some CH 3 COOH to be consumed in the production of a salt, NaCH 3 COO. Ideally, this would cause a lowered hydronium ion concentration due to the presence of the common ion, CH 3 COO - , from NaCH 3 COO, causing a leftward shift in the equilibrium. Equation (3) can then be used to calculate for the new pH. The changes in pH for solutions 1a and 1b are -0.57 and +0.60 respectively. Solutions 2a and 2b follow a similar course as that of the divisions of solution 1. But because this is a buffer solution, the change in pH level is different, or lowered to some extent. NaCH 3 COO completely dissociates into Na + and CH 3 COO - but CH 3 COOH undergoes partial ionization only. Because of the common ion CH 3 COO - , there will be minimal change of pH in the solution compared to solutions 1a and 1b. CH 3 COO - will react with HCl, a strong acid, for solution 2a while CH 3 COOH will react with NaOH, a strong base, for solution 2b.
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  • Fall '17
  • Sir Jaden Smith
  • pH, Buffer solutions, NH3

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