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CHEM 1315 Study Guide - Final Guide: Sodium Acetate, Buffer Solution, Quadratic Equation

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10 May 2017

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University of Oklahoma

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If we have two solutes each containing the same ion, the equilibrium is affected. Consider adding acetic acid (ch3cooh) and sodium acetate (nach3coo) to water. The acetate anion from sodium acetate will inhibit acetic acid from dissociating (le chatelier). First consider the situation with only acetic acid, at 0. 10 m. X = 1. 34 10 3 which means the ph = 2. 87 and the % ionization = 1. 34% Next, consider adding 0. 050 moles of sodium acetate to 1. 00 l of our acetic acid solution (assume volume does not change). Equilibrium (note: the equilibrium must shift to the left from le chatelier) 1. 34 x 10 3 - x 5. 134 x 10 2 -x. This gives [h+] = 3. 5 x 10 5 (using the quadratic equation) which means the ph = 4. 46 and the % ionization = 0. 035% Comment: the last calculation was hard because we had to use the quadratic equation.

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Related Questions

what are the buffer solutions for these 3 slide questions?

Prelab questions and AP Chemistry Review questions, thanks

Buffers are also nature, alkaline. The addi- important in certain commercial are, by ciuric acid and of citric acid this alialinity and prevents possible beams to the skin and scalp, Baby often other pathogens Even sodium to buffer the to a slightly growth of bacteria and PH alcohol production acidic pH six, which inhibits the pH range. If the can rely on buffers. Yeasts that fement the sugars only within narrow work side the range of 4.0-6.0. yeast growth may be inhibited or even destroyed. Experiment overview The purpose of this advanced inquiry lab is to investigate the capacity and buffer componen benzoate products. Many household products contain buffering chemicals such as citric acid, sodium carbonate sodium citric acid curve phosphates or phosphoric acid. The lab begins with an introductory activity generating the for and identify the buffering regions n neutralization of a weak acid. provide a beverages design of a procedure uo the agents in eight foods and buffer capacity drugs. Procedures may include creating titration different household products, the eurves, calculating pK, values, ana ng and composition. Students may recommend additional consumer products for further inquiry Pre-Lab Questions Figure 2 shows dae pH curve for the titration of 25.0 mL of a o 10 M solution with 0.10 M sodium of acetic acid. ,cool, hydroxide solution. K. of acetic acid is 1.8 x 10 1. the concentration of equal the concentration of the acetate ion, At what point in the titration does acetic acid, CH,COOH, CH Coor? What is the pH of the equimolar, ideal buffer solution at this point? 2 If weak acid base pair of 10:1 and 1:10. this is effective as a buffer between the concentration ratios for the conjugate ac what pH range does this cover? Volume NaOH Added, mL Figure 2. Titration of Acetic Acid with NaOH 3. What measurements are needed in the titration of a weak acid? Explain in detail the technique or procedure for adding the titrant to accurately determine the concentration and pkaof weak acid. TN7665 may be reproduced or Iranmitad any info

Exp 25. pH Measurements- Buffers and Their Properties

I need help with finding kPa in the calculations part C for 1,2, and 3 Thanks!

Experiment 25 pH Measurements- Buffers and Their Properties ne of the more important propertics of an aqucous solution is its concentration of hydrogen ion. The H ion (or, more precisely, the HO ion) has a great effect on the solubility of many inorganic and organic species, on the nature of complex mctallic cations found in solutions, and on the rales of many chemical reac- tions. It is important that we know how to measure the concentration of hydrogen ion and understand its effect on solution properties. For convenience, the concentration of H' ion is frequendly expressed as the pH of the solution rather than as molarity. The pH of a solution is defined by the following equation: pH =-log where the logarithm is taken to the base 10. If [H] is I x 10-*moles per liter, the pH of the solution is, by the equation, 4. If the [H'1 is 5 x 10 M, the pH is 1.3, Basic solutions can also be described in terms of pH. In aqucous solutions the following equilibrium rela- tion will always be obeyed: In distilled water url equals [ OHÂ·], s, by Equation 2, [HÂ·l must be l x 10" M. Therefore, the pH of distilled water is 7. Solutions in which [H.] > 10é¨are sad to be acidic and will have a pH 7. A solution with a pH of 10 will have a [H"] of 1 x 10 0 M and a (OH] of I x 10M We measure the pH of a solution experimentally in two ways. In the first of these we use a chemical called an indicator, which is sensitive to pH. These substances have colors that change over a relatively short pH range (about two pH units) and can, when properly chosen, be used to delermine roughly the pH of a solu tion. Two very common indicators are litmus, usually used on paper, and phenolphthalein, the most common indicator in acid-base titrations. Litmus changes from red to blue as the pH of a solution goes from about 6 to about 8. Phenolphthalein changes from colorless to red as the pH goes from 8 to 10. A given indicator is useful for determining pH only in the region in which it changes color. Indicators are available for measure- ment of pH in all the important ranges of acidity and basicity. By matching the color of a suitable indicator in a solution of known pH with that in an unknown solution, one can determine the pH of the unknown to within about 0.3 pH units. The other method for finding pH is with a device called a pH meter. In this device two electrodes, one of which is sensitive to [H']. are immersed in a solution. The potential between the two electrodes is related to the pH. The pH meter is designed so that the scale will directly furnish the pH of the solution. A pH meter gives much more precise measurement of pH than does a typical indicator and is ordinarily used when an accurate determination of pH is needed Some acids and bases undergo substantial ionization in water, and are called strong because of their essentially complete ionization in reasonably dilute solutions. Other acids and bases, because of incomplete ionization (often far less than even 1% in 0.1 M solution), are called weak. Hydrochloric acid, HC, and sodium hydroxide, NaOH, are typical examples of a strong acid and a strong base. Acetic acid, HC,H,0, and ammonia, NH, are classic examples of a weak acid and a weak base. A weak acid will ionize according to the Law of Chemical Equilibrium: HBlaq)--H'(aq) + B-(aq)