The Determination of Acid Constant Ka

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The Determination of Acid Constant Ka

When an acid or base dissolves in water, a certain percentage of the acid or base particles will break up, or dissociate, into oppositely charged ions. The Arrhenius theory defines an acid as a compound that can dissociate in water to yield hydrogen ions, H+, and a base as a compound that can dissociate in water to yield hydroxide ions, OH. The base sodium hydroxide, NaOH, dissociates in water to yield the required hydroxide ions, OH-, and also sodium ions.The objectives of this experiment were: a) to review the concept of simple acid-base reactions; b) to review the basic lab procedure of a titration and introduce the student to the concept of a primary standard and the process of standardization. From my graph my equivalence point was PH=8.80, volume is 15 ml , pka=4.80, ka=1.58 × 10-5 Introduction

Acids dissolve in water with dissociation and the formation of hydrogen ions. According to the equation HA(aq)+H2O(l)-------H3O+(aq)+A-(aq). They are classed as strong or weak depending on the extent to which this dissociation occurs. Bases are substances that liberate hydroxide ions in water. They too are classed as weak or strong. The Henderson –Hasselbach equation is derived from pH=pka+log base/acid. In the experiment titration was used to study acid-base neutralization reaction quantitatively. In acid-base titration experiment, a solution of accurately KH concentration was added gradually to another solution of NaOH concentration until the chemical reaction between the two solutions were completed. The equivalence point was the point at which the acid was completely reacted with or neutralized by the base, and no HA remains; all is in the form of the conjugate base, A- and water. When these concentrations are equal log[A-]/[HA] is zero and pH=pka. The point was signaled by a changing of color of an indicator that had been added to the acid solution. The common indicator was Phenolphthalein which was colorless in acidic and neutral solutions, but a pink was the result in basic solutions. This experiment has four objectives: (1) to standardize a NaOH solution , (2) to determine the molarity and pH of the acetic acid solution, (3) develop a titration curve to determine the molarity and ka, (4) and perform a half-titration to determine the Ka of Acetic Acid. It was expected that the NaOH solution would have a 0.1 molarity, the experimental ionization constant of acetic acid would be close to ka= 1.76 X10^-5. Experimental

The very first thing in beginning the experiment is to set up the apparatus, rinse out the apparatus using its chemicals that’s going to be put in there, for example rinse out the burette using some HCl to get out any unwanted particles which may be present. Fill up the burette with HCl using a funnel. At this make sure there is protection in the eyes by wearing goggles so that nothing can into the eyes. Whilst putting the HCl in the burette, make sure the burette is clamped into the stand firmly so to avoid any breakages. When the burette is filled to a reasonable level, read how much HCl is in there, first make sure the funnel is out as this can alter the reading. When reading of the burette make sure to read below the meniscus at eye level. In the first part of the experiment standardization of the NaOH solution was done, in order to know the total concentration. A 500 ml of 0.1 M NaOH was diluted to an amount of 6.0 M NaOH. On an analytical balance 0.5 grams of KHP were measured using the bottle technique , and the total weight was recorded. To standardize the supposed 0.1 M NaOH solution prepared in part 1, NaOH was titrated with potassium hydrogen phthalate of known amount. Complete neutralization, the reaction of one mole of NaOH with one mole of KHP, had occurred with the indicator phenolphthalein turning to a light pink color indicating that the titration has come to an endpoint. The volume was recorded. Three trials were...
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