Acids percent dissociation

Calculating the pH of Weak Acid Solutions The pH of a Mixture of Weak Acids Percent Dissociation ... 

Thioglycolic acid, HSCH2CO2H, a substance used in depilatory agents (hair removers) has pKa = 3.42. What is the percent dissociation of thioglycolic acid in a buffer solution at pH = 3.0 ... 

Words that can be used as topics in essays 5% rale buffer common ion effect equilibrium expression equivalence point Henderson-Hasselbalch equation heterogeneous equilibria homogeneous equilibria indicator ion product, P Ka Kb Kc Keq KP Ksp Kw law of mass action Le Chatelier s principle limiting reactant method of successive approximation net ionic equation percent dissociation pH P Ka P Kb pOH reaction quotient, Q reciprocal rule rule of multiple equilibria solubility spectator ions strong acid strong base van t Hoff equation weak acid weak base... 

IQ, the acid dissociation constant, for an acid is 9 x 10-4 at room temperature. At this temperature, what is the approximate percent dissociation of the acid in a 1.0 M solution ... 

The percent dissociation of a weak acid is the fraction of acid molecules that dissociate compared with the initial concentration of the acid, expressed as a percent. (Some chemists refer to percent dissociation as percent of dissociation.] The percent dissociation depends on the value of Ka for the acid, as well as the initial concentration of the weak acid. The following Sample Problems show how to solve problems that involve percent dissociation. 

You need to find iQ and the percent dissociation for propanoic acid. 

The value of and the percent dissociation are reasonable for a weak acid. 

Calculate the pH of a sample of vinegar that contains 0.83 mol/L acetic acid. What is the percent dissociation of the vinegar ... 

Ka = 6.3 X 10 ). Its structure is shown below. Calculate the pH and the percent dissociation of each of the following solutions of benzoic acid. Then use Le Chatelier s principle to explain the trend in percent dissociation of the acid as the solution becomes more dilute. 

When the pH of the solution and the pKa are equal, 50% of the acid will have dissociated into ions. The percent dissociation of an acid or base can be calculated if the pH of the solution and the pKa of the compound are known (Guswa et al., 1984) ... 

A quantitative measure of the degree of dissociation is given by the equilibrium constant for the acid or base. The higher the equilibrium constant is, the greater the percent dissociation of the acid or base. Therefore, a higher equilibrium constant means a stronger acid or base. Equilibrium constants, K and K, are listed for several com-mon weak acids and bases in Table 13.4. 

The activity coefficients of hydrobromic acid in the mixed solvents are lower, as expected, than those in water (20). Hydrobromic acid completely dissociates in the mixed solvents (e = 49.5 at 298.15° K for the 50 mass percent monoglyme) under investigation. Figure 2 clearly indicates that at a particular molality, the stoichiometric activity coefficient of hydrochloric acid is lower than that of hydrobromic acid in the same mixed solvent, and the heat capacity changes (Cp — Cp) also suggest that there are no ion-pair formations. 

In addition to Ka, another useful measure of the strength of a weak acid is the percent dissociation, defined as the concentration of the acid that dissociates divided by the initial concentration of the acid times 100% ... 

Take, for example, the 1.00 M acetic acid solution in Problem 15.14a. If you solved that problem correctly, you found that 1.00 M CH3CO2H has an H30 + concentration of 4.2 X 10 3 M. Because [H30+] equals the concentration of CH3C02H that dissociates, the percent dissociation in 1.00 M CH3C02H is 0.42% ... 

In general, the percent dissociation depends on the acid and increases with increasing value of Ka. For a given weak acid, the percent dissociation increases with increasing dilution, as shown in Figure 15.8. The 0.0100 M CH3C02H solution in Problem 15.14b, for example, has [H30+] = 4.2 X 10 4 M, and the percent dissociation is 4.2% ... 

FIGURE 15.8 The percent dissociation of acetic acid increases as the concentration of the acid decreases. A 100-fold decrease in [CH3C02H] results in a 10-fold increase in the percent dissociation. 

Phenol (C5H5OH) is a weak acid used as a general disinfectant and in the manufacture of plastics. Calculate the pH and the concentrations of all species present (H30+,C6H50, C6H50H, and OH-) in a 0.10 M solution of phenol (Ka = 1.3 X 10-10). Also calculate the percent dissociation. 

A typical aspirin tablet contains 324 mg of aspirin (acetylsalicylic acid, C9H8O4), a monoprotic acid having Ka = 3.0 X 10-4. If you dissolve two aspirin tablets in a 300 mL glass of water, what is the pH of the solution and the percent dissociation ... 

Calculate the percent dissociation in each of the following solutions. What is the quantitative relationship between the percent dissociation and the concentration of the acid What is the quantitative relationship between the percent dissociation and the value of Ka ... 

Beginning with the equilibrium equation for the dissociation of a weak acid HA, show that the percent dissociation varies directly as the square root of Ka and inversely as the square root of the initial concentration of HA when the concentration of HA that dissociates is negligible compared with its initial concentration. 

Use data from the Testing pH activity (eChapter 15.7) to determine the of methylamine, CH3NH2. Determine the percent dissociation of methylamine at 0.10 M, 0.0010 M, and 1.0 X 10 5 M concentrations. What is the relationship between the percent dissociation and the pH for a weak base Would you expect the same relationship for a weak acid Explain. 

The following pictures represent solutions of a weak acid HA that may also contain the sodium salt NaA. Which solution has the highest pH, and which has the largest percent dissociation of HA (Na+ ions and solvent water molecules have been omitted for clarity.)... 

The real importance of the Henderson-Hasselbalch equation, particularly in biochemistry, is that it tells us how the pH affects the percent dissociation of a weak acid. Suppose, for example, that you have a solution containing the amino acid glycine, one of the molecules from which proteins are made, and that the pH of the solution is 2.00 pH units greater than the pKa of glycine ... 

The common-ion effect is the shift in the position of an equilibrium that occurs when a substance is added that provides more of an ion already involved in the equilibrium. An example is the decrease in percent dissociation of a weak acid on addition of its conjugate base. 

It is often useful to specify the amount of weak acid that has dissociated in achieving equilibrium in an aqueous solution. The percent dissociation is defined as follows ... 

For a given weak acid, the percent dissociation increases as the acid becomes more dilute. For example, the percent dissociation of acetic acid (HC2H302, Ka = 1.8 X 10-5) is significantly greater in a 0.10 M solution than in a 1.0 M solution. 

The effect of dilution on the percent dissociation and [H+] of a weak acid solution. 

The more dilute the weak acid solution, the greater the percent dissociation. 

Since Q is less than Ka, the system must adjust to the right to reach the new equilibrium position. Thus the percent dissociation increases as the acid becomes more dilute. This behavior is summarized in Fig. 7.5. In Example 7.4 we see how the percent dissociation can be used to calculate the Ka value for a weak acid. 

The small value for the percent dissociation clearly indicates that HC3H5O3 is a weak acid. Thus the major species in the solution are the undissociated acid and water ... 

The change needed to reach equilibrium can be obtained from the percent dissociation and Equation (7.3). For this acid... 

Sketch two graphs (a) percent dissociation of weak acid HA versus initial concentration of HA ([HA]o), and (b) H+ concentration versus [HA]0. Explain both. 

Calculate the percent dissociation of the acid in each of the following solutions. 

Compare these values for [H+] and percent dissociation of HF with those for a 1.0 M HF solution, where [H+] = 2.7 X 10 2 M and the percent dissociation is 2.7%. The large difference clearly shows that the presence of the F- ions from the dissolved NaF greatly inhibits the dissociation of HF. The position of the acid dissociation equilibrium has been shifted to the left by the presence of F ions from NaF. 

---