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A. Acids and bases Historically there have been several influential definitions of acids and bases. The main definition in current use and the one we'll be using is Brønsted-Lowry definition, which defines acids as compounds capable of releasing a proton (H +) while bases are compounds that can accept a proton. Acid-base reactions are therefore proton-transfer reaction, where an acid releases a proton to form its conjugated base and a base accepts the proton to form its conjugated acid. The more general Lewis definition defines acids as compounds capable of receiving an electron pair while bases are electron pair donors. An older definition by Svante Arrhe-nius defined acids as compounds releasing a proton and bases as releasing a hydroxide anion. The process of releasing a proton (or any other ion) is called dissociation. As any other chemical reaction, dissociation will reach equilibrium, which can be characterised by an equilibrium constant, K. For a generic acid HA dissociating in water Protons in aqueous solutions become hydrated to H 3 O + called the hy-dronium or hydroxonium ion—the conjugated acid of water. HA + H 2 O −→ A − + H 3 O + (1) the equilibrium constant is K = [H 3 O + ][A − ] [HA][H 2 O] (2) Since the concentration of water in a dilute aqueous solution is pretty much constant we can rewrite the expression as K A = [H 3 O + ][A − ] [HA] (3) where K A is the dissociation constant of the acid (note that it is not the same as the original equilibrium constant). The value of K A is a good measure of the strength of an acid, i.e its ability to dissociate in solution: acids with K A > 10 −2 are usually called strong acids and the rest are weak acids. Please note that strength of an acid is not the same as its corrosive power or dangerousness! K A is often expressed as its negative decadic logarithm pK A = − log 10 K A. For example, HF is a relatively weak acid (K A = 6.8×10 −4) but can cause serious burns or dissolve glass. As the majority of acid-base reactions in biochemistry occur in aqueous solutions it is important to understand the role of water in these processes. Water is an amphoteric compound which means that it can both accept and donate a proton, i.e. it can act as an acid or as a base. In reaction (1) it accepts a proton from HA and thus behaves as a base. Similarly, water could behave as an acid and donate a proton to ammonia NH 3 + H 2 O − − − − NH + 4 + OH − (4) Moreover water can even donate protons to itself in a process called self-dissociation H 2 O + H 2 O − − − − H 3 O + + OH − (5)
Experimental Report 13: " pH Buffer Solutions "
Preparing different pH buffer solutions and find by comparison which buffer has the higher buffer capacity were the main objectives in this experiment. In order to accomplish the objectives, a solution of hydrochloric acid (HCl) and sodium hydroxide was prepared, as well as another solution with acetic acid (HC 2 H 3 O 2) and sodium acetate (NaC 2 H 3 O 2). Consequently the solutions were titrated in order to obtain a pH more than , as needed. The general results obtained in the experiment were that the best buffer ± 1 capacity using hydrochloric acid were in the next order: 4.27 with a 0.39 ml/pH unit, 4.03 with a 0.15 ml/pH unit and 3.74 with a 0.08 ml/pH unit. In the other hand, using s odium hydroxide , the best buffer capacity resulted as following: 3.74 in first place, with a 0.98 ml/pH unit, 4.27 with a 0.77 ml/pH unit and finally 4.03 with a 0.71 ml/pH unit. GRAPHICAL ABSTRACT OBJECTIVES AND HYPOTHESIS The main objective of this experiment is to prepare different pH buffer solutions (4.76, 4.5, and 4.2) with a weak conjugate acidbase pair, and titrate them with hydrochloric acid and sodium hydroxide in order to compare which of the three buffers has the highest capacity. We hypothesized that the greatest buffering capacity would be obtained if the buffer solution contains equimolar concentrations of the weak acid and its conjugate base. INTRODUCTION A buffer is a solution that can resist pH change upon the addition of an acidic or basic components. It is able to neutralized small amounts of added acid or base, maintaining the pH relatively
The purpose of this experiment is to identify the best buffer by knowing the effect of concentration of buffer and the ratio of the conjugate base to the weak acid. For the Effect of Concentration of Buffer, both the 0.05 M and 0.1 M concentrations of Phosphate and Acetate buffer solutions were effective buffers but 0.1M is the best buffer because the greater the concentration the more the buffering species to resist the drastic change in pH upon the addition of NaOH. It can be inferred that the greater the concentration of the buffer, the greater the buffering capacity. For the second part, the closer the pH of the buffer solution to the pKa the greater is the buffering capacity. IThe best buffers for the second part should be the 0.5M phosphate buffer at pH=7.2 and 0.5M Acetate buffer at pH=4.7. However, errors in the conduct of experiment might have caused the error in 0.5M phosphate buffer results. In the titration of Glycine, the pKa for each of the three trials were identified and the pKa for the first and second trial is 10.819 while for the third trial it is 10.754.
TITRATION OF ACID AND BASES LAB REPORT_CHE142
2022
This experiment relies on the titration technique to determine the unknown concentration of monoprotic acid in the solution. There are two types of acids that are used in this experiment which is (1) Potassium Hydrogen Phthalate (KHP) and (2) Unknown Acid. The base that was used in this experiment is Sodium Hydroxide (NaOH) solution. This experiment involves two sections or parts: (i) Standardization of Sodium Hydroxide (NaOH) solution using Potassium Hydrogen Phthalate (KHP) solution, and (ii) Analysis of Unknown acidic solution using standardize NaOH solution.
Buffer preparation and pH measurements
TITLE: Buffer preparation and pH measurements RESULTS Weak Acid (W.A.), Strong Acid (S.A), Weak Base (W.B.) and Strong Base (S.B.) are paired weak acid + strong base, weak base + strong acid, and weak acid + weak base to create a three different buffer solutions with a pH target of 7.05. Potassium phosphate, dibasic (K2HPO4 174 g/mol) served as the W.B., and potassium phosphate, monobasic (KH2PO4 136 g/mol), served as the W.A. NaOH was the S.B. and HCl was the S.A. Table 1 reviews results of the pH for each of the three buffer types and their average pH. The ingredients were measured out with a target pH of 7.05, a volume of 50 mL, and a concentration of 0.1M using the Henderson-Hasselbach equation. The pKa of the phosphate buffer is 7.2.
The Conditions Needed for a Buffer to Set the pH in a System
InTech eBooks, 2017
It is a known fact that buffer systems are widely used in industry and diverse laboratories to maintain the pH of a system within desired limits, occasionally narrow. Hence, the aim of the present work is to study the buffer capacity and buffer efficacy in order to determine the useful conditions to impose the pH on a given system. This study is based on the electroneutrality and component balance equations for a mixture of protons polyreceptors. The added volume equations are established, V, for strong acids or bases, as well as the buffer capacity equations with dilution effect, β dil , and the buffer efficacy, ε, considering that the analyte contains a mixture of the species of the same polyacid system or various polyacid systems. The ε index is introduced to define the performance of a buffer solution and find out for certain, whether the buffer is adequate to set the pH of a system, given the proper conditions and characteristics.