Weak acids calculating

Given Ka for a weak acid, calculate Kh for its conjugate base (or vice-versa). 

Using the assumptions we ordinarily make in calculating the pH of an aqueous solution of a weak acid, calculate the pH of a 1.0 X 10 6 M solution of hypo-bromous acid (HBrO, Ka = 2 X 10-9). What is wrong with your answer Why is it wrong Without trying to solve the problem, tell what has to be included to solve the problem correctly. 

This is a typical weak acid calculation of the type introduced in Chapter 7. The pH is 2.87. (Check this value yourself.)... 

How is the concentration of counteranion from a weak acid calculated ... 

Weak acid. Calculate the pH of a 0.060 M solution of propionic acid. 

Construct ladder diagrams for the following diprotic weak acids (H2L), and estimate the pH of 0.10 M solutions of H2L, HL , and Using the systematic approach, calculate the pH of each of these solutions. 

The extraction efficiency, therefore, is almost 75%. When the same calculation is carried out at a pH of 5.00, the extraction efficiency is 60%, but the extraction efficiency is only 3% at a pH of 7.00. As expected, extraction efficiency is better at more acidic pHs when HA is the predominate species in the aqueous phase. A graph of extraction efficiency versus pH for this system is shown in Figure 7.23. Note that the extraction efficiency is greatest for pHs more acidic than the weak acid s piQ and decreases substantially at pHs more basic than the pi A- A ladder diagram for HA is superimposed on the graph to help illustrate this effect. 

A sample contains a weak acid analyte, HA, and a weak acid interferent, HB. The acid dissociation constants and partition coefficients for the weak acids are as follows Ra.HA = 1.0 X 10 Ra HB = 1.0 X f0 , RpjHA D,HB 500. (a) Calculate the extraction efficiency for HA and HB when 50.0 mF of sampk buffered to a pH of 7.0, is extracted with 50.0 mF of the organic solvent, (b) Which phase is enriched in the analyte (c) What are the recoveries for the analyte and interferent in this phase (d) What is the separation factor (e) A quantitative analysis is conducted on the contents of the phase enriched in analyte. What is the expected relative erroi if the selectivity coefficient, Rha.hb> is 0.500 and the initial ratio ofHB/HA was lO.O ... 

Titrating a Weak Acid with a Strong Base For this example let s consider the titration of 50.0 mL of 0.100 M acetic acid, CH3COOH, with 0.100 M NaOH. Again, we start by calculating the volume of NaOH needed to reach the equivalence point thus... 

Before adding any NaOH the pH is that for a solution of 0.100 M acetic acid. Since acetic acid is a weak acid, we calculate the pH using the method outlined in Chapter 6. 

Any solution containing comparable amounts of a weak acid, HA, and its conjugate weak base, A-, is a buffer. As we learned in Chapter 6, we can calculate the pH of a buffer using the Henderson-Hasselbalch equation. 

The approach that we have worked out for the titration of a monoprotic weak acid with a strong base can be extended to reactions involving multiprotic acids or bases and mixtures of acids or bases. As the complexity of the titration increases, however, the necessary calculations become more time-consuming. Not surprisingly, a variety of algebraic and computer spreadsheet approaches have been described to aid in constructing titration curves. 

Where Is the Equivalence Point We have already learned how to calculate the equivalence point for the titration of a strong acid with a strong base, and for the titration of a weak acid with a strong base. We also have learned to sketch a titration curve with a minimum of calculations. Can we also locate the equivalence point without performing any calculations The answer, as you may have guessed, is often yes ... 

Tartaric acid, H2C4H4O6, is a diprotic weak acid with a pK i of 3.0 and a pK 2 of 4.4. Suppose you have a sample of impure tartaric acid (%purity > 80) and that you plan to determine its purity by titrating with a solution of 0.1 M NaOH using a visual indicator to signal the end point. Describe how you would carry out the analysis, paying particular attention to how much sample you would use, the desired pH range over which you would like the visual indicator to operate, and how you would calculate the %w/w tartaric acid. 

Calculate or sketch (or both) the titration curves for 50.0 ml of a 0.100 M solution of a monoprotic weak acid (pfQ = 8) with 0.1 M strong base in (a) water and (b) a non-aqueous solvent with ffg = 10 . You may assume that the change in solvent does not affect the weak acid s pfQ. 

The proportion of ionized and unionized forms of a chemical compound can be readily calculated according to the above equation. It can be easily seen that pK is also a pH value at which 50% of the compound exists in ionized form. The ionization of weak acids increases as the pH increases, whereas the ionization of weak bases increases when the pH decreases. As the proportion of an ionized chemical increases, the diffusion of the chemical through the biological membranes is greatly impaired, and this attenuates toxicokinetic processes. For example, the common drug acetosalicylic acid (aspirin), a weak acid, is readily absorbed from the stomach because most of its dose is in an unionized form at the acidic pH of the stomach. 

From Eq. (4) and the data of Tables 1, 2, and 3 it can easily be calculated that the rate of hydrolysis of these enamines should rapidly reach a maximum value in weakly acidic solutions at decreasing pH, assuming that no buffer is used. The raTe should then be constant and pH-independent. For enamine (1) Eq. (4) reduces to k = Kkn,ou and this rate must be the same as that at the intersection of the straight lines in Fig. 1. This appeared to be true for the observed rate at pH 2 15). 

This relationship is known as the Henderson-Hasselbalch equation. Thus, the pH of a solution can be calculated, provided and the concentrations of the weak acid HA and its conjugate base A are known. Note particularly that when [HA] = [A ], pH = pAl,. For example, if equal volumes of 0.1 MHAc and 0.1 M sodium acetate are mixed, then... 

Given the ionization constant of a weak acid and its original concentration, the H+ concentration in solution is readily calculated. The approach used is the inverse of that followed in Example 13.5, where Kz was calculated knowing [H+] here, K is known and [H+] must be calculated. 

In most of the problems you will work, the approximation a — x a is valid, and you can solve for [H+] quite simply, as in Example 13.7, where x = 0.012a. Sometimes, though, you will find that the calculated [H+] is greater than 5% of the original concentration of weak acid. In that case, you can solve for x by using either the quadratic formula or the method of successive approximations. 

We showed in Section 13.4 how Ka of a weak acid can be used to calculate [H+] in a solution of that acid. In a very similar way, Kb can be used to find [OH-] in a solution of a weak base. 

Given the pH and original concentration of a weak acid solution, calculate Ka. 

Given Kz of a weak acid and its original concentration, calculate [H+],... 

Calculate Ka for the weak acids that have the following pK values. 

Penicillin (MM = 356 g/md), an antibiotic often used to treat bacterial infections, is a weak acid. ItsfQis 1.7 X 10-3. Calculate [H+] in solutions prepared by adding enough water to the following to make 725 ml.. 

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