Equilibrium expression weak acid

Br0nsted-Lowery acids are H+ donors and bases are H+ acceptors. Strong acids dissociate completely in water. Weak acids only partially dissociate, establishing an equilibrium system. Weak acid and base dissociation constants (Ka and Kb) describe these equilibrium systems. Water is amphoteric, acting as both an acid or a base. We describe water s equilibrium by the Kw expression. A pH value is a way of representing a solution s acidity. Some salts and oxides have acid-base properties. A Lewis acid is an electron pair acceptor while a Lewis base is an electron pair donor. 

The equilibrium constant for equation 6.13 is K. Since equation 6.13 is obtained by adding together reactions 6.11 and 6.12, may also be expressed as the product of Ka for CH3COOH and Kb for CH3COO-. Thus, for a weak acid, HA, and its conjugate weak base, A-,... 

In the discussion of the relative acidity of carboxylic acids in Chapter 1, the thermodynamic acidity, expressed as the acid dissociation constant, was taken as the measure of acidity. It is straightforward to determine dissociation constants of such adds in aqueous solution by measurement of the titration curve with a pH-sensitive electrode (pH meter). Determination of the acidity of carbon acids is more difficult. Because most are very weak acids, very strong bases are required to cause deprotonation. Water and alcohols are far more acidic than most hydrocarbons and are unsuitable solvents for generation of hydrocarbon anions. Any strong base will deprotonate the solvent rather than the hydrocarbon. For synthetic purposes, aprotic solvents such as ether, tetrahydrofuran (THF), and dimethoxyethane (DME) are used, but for equilibrium measurements solvents that promote dissociation of ion pairs and ion clusters are preferred. Weakly acidic solvents such as DMSO and cyclohexylamine are used in the preparation of strongly basic carbanions. The high polarity and cation-solvating ability of DMSO facilitate dissociation... 

Case 3. Salt of a weak acid and a weak base. The hydrolytic equilibrium is expressed by the equation ... 

We calculate the pH of solutions of weak bases in the same way as we calculate the pH of solutions of weak acids—by using an equilibrium table. The protonation equilibrium is given in Eq. 9. To calculate the pH of the solution, we first calculate the concentration of OH ions at equilibrium, express that concentration as pOH, and then calculate the pH at 25°C from the relation pH + pOH = 14.00. For very weak or very dilute bases, the autoprotolysis of water must be taken into consideration. 

STRATEGY Because NH4+ is a weak acid and Cl- is neutral, we expect pH < 7. We treat the solution as that of a weak acid, using an equilibrium table as in Toolbox 10.1 to calculate the composition and hence the pH. First, write the chemical equation for proton transfer to water and the expression for Ca. Obtain the value of Ka from Kh for the conjugate base by using K, = KxJKh (Eq. 11a). The initial concentration of the acidic cation is equal to the concentration of the cation that the salt would produce if the salt were fully dissociated and the cation retained all its acidic protons. The initial concentrations of its conjugate base and H30+ are assumed to be zero. 

For each of the following weak acids, write the proton transfer equilibrium equation and the expression for the equilibrium constant Kv Identify the conjugate base, write the appropriate proton transfer equation, and write the expression for the basicity constant Kb. (a) HC102 (b) HCN ... 

STRATEGY First, identify the weak acid and its conjugate base. Then, write the proton transfer equilibrium between them, rearrange the expression for Ka to give [H30+], and find the pH, by using the approximation that the equilibrium concentrations of the acid and its conjugate base are essentially the same as the initial concentrations. 

The values of [HA] and [A ] in this expression are the equilibrium concentrations of acid and base in the solution, not the concentrations added initially. However, a weak acid HA typically loses only a tiny fraction of its protons, and so [HA] is negligibly different from the concentration of the acid used to prepare the buffer, [HA]initia. Likewise, only a tiny fraction of the weakly basic anions A- accept protons, and so [A-] is negligibly different from the initial concentration of the base used to prepare the buffer. With the approximations A ] [base]initia and [HA] [acid]initia, we obtain the Henderson-Hasselbalch equation ... 

Step 5 Use an equilibrium table to find the H.O concentration in a weak acid or the OH concentration in a weak base. Alternatively, if the concentrations of conjugate acid and base calculated in step 4 are both large relative to the concentration of hydronium ions, use them in the expression for /<, or the Henderson—Hasselbalch equation to determine the pH. In each case, if the pH is less than 6 or greater than 8, assume that the autoprotolysis of water does not significantly affect the pH. If necessary, convert between Ka and Kh by using Kw = KA X Kb. 

C16-0105. Write the equilibrium reaction and equilibrium constant expression for each of the following processes (a) Trimethylamine, (CH3)3 N, a weak base, is added to water, (b) Hydrofluoric acid, HF, a weak acid, is added to water, (c) Solid calcium sulfate, CaSOq, a sparingly soluble salt, is added to water. 

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... 

Since this is the equilibrium constant associated with a weak acid dissociation, this particular Kc is the weak add dissociation constant, Ka. The IQ expression is ... 

The [HA] is the equilibrium molar concentration of the undissociated weak acid, not its initial concentration. The exact expression would then be [HA] = Minitia y — [H+], where Minitia y is the initial concentration of the weak acid. This is true because for every H+ that is formed an HA must have dissociated. However, many times if the Ka is small, you can approximate the equilibrium concentration of the weak acid by its initial concentration, [HA] Minitia y. 

The common-ion effect is an application of Le Chatelicr s principle to equilibrium systems of slightly soluble salts. A buffer is a solution that resists a change in pH if we add an acid or base. We can calculate the pH of a buffer using the Henderson-Hasselbalch equation. We use titrations to determine the concentration of an acid or base solution. We can represent solubility equilibria by the solubility product constant expression, Ksp. We can use the concepts associated with weak acids and bases to calculate the pH at any point during a titration. 

The weak acid curves can also be calculated. This involves the use of the equilibrium constant expression for a weak monoprotic acid ionization ... 

Weak acids are not completely ionised in aqueous solution and are in equilibrium with the undissociated acid, as is the case for water, which is a very weak acid. The dissociation constant Ka is given by the expression below ... 

The way indicators work can be understood if we consider indicators to be weak acids. If we let Hln represent the general formula of an indicator, then we can write the following equilibrium expression ... 

Independent equations cannot be derived from one another. As a trivial example, the equations a = b + c and 2a = 2b + 2c are not independent. The three equilibrium expressions for Ka, K and K for a weak acid and its conjugate base provide only two independent equations because we can derive Kb from /(, and Kw Kt, = K ll. ... 

In Chapter 16, we apply the fundamental general equilibrium expression to gaseous equilibrium reactions. In this chapter, we apply the same expression to the equilibria that involve weak acids and bases in aqueous solution, the principal difference being that all concentrations are expressed in moles/liter (rather than in atmospheres as for gases). All the general conclusions given in Chapter 16, and summarized in the principle of Le Chatelier, apply to equilibria in solutions as well as to those in gases. 

PK. A measurement of the complete ness of an incomplete chemical reaction. It is defined as the negative logarithm ito the base 101 of the equilibrium constant K for the reaction in question. The pA is most frequently used to express the extent of dissociation or the strength of weak acids, particularly fatty adds, amino adds, and also complex ions, or similar substances. The weaker an electrolyte, the larger its pA. Thus, at 25°C for sulfuric add (strong acid), pK is about -3,0 acetic acid (weak acid), pK = 4.76 bone acid (very weak acid), pA = 9.24. In a solution of a weak acid, if the concentration of undissociated acid is equal to the concentration of the anion of the acid, the pAr will be equal to the pH. 

Strategy Acetic acid is a weak acid consequently, we expect the molarity of H30+ ions to be less than 0.10 moI-L-1 and, therefore, its pH to be greater than 1.0. To find the actual value, we set up an equilibrium table S with the initial molarity of acid equal to 0.10 mol-L 1 and allow the molarity of acid to decrease by x mol-L1 to reach equilibrium. Assume that the presence of acid dominates the pH and therefore that the autoprotolysis of 5 water need not be considered. We assume x is less than about 5% of the ini-j rial molarity of acid and simplify the expression for the equilibrium constant f by ignoring x relative to the initial molarity of the acid. This assumption i must be verified at the end of the calculation. 

It is helpful to rewrite the expression for fCa to show how the pH of a solution relates to the pfCa of the acid present. Let s consider a solution containing a weak acid HA (such as CH3COOH or NH4+) and the conjugate base anion A- (such as CH3C02 or NH3, respectively). The proton transfer equilibrium is... 

With a weak acid, the calculations are a bit more complex. We begin by computing the initial pH, using the method of Example 1, pH = 2.79. At the equivalence point, we have 50.0 mL of 0.060 M sodium acetate. Since acetate is the base conjugate to a weak acid, the solution is basic and we must compute the pH. From eq (13-12) for acetate ion, Kb = KyKa =5.62 x 10 10, and the equilibrium constant expression is... 

In considering reactions in biochemical systems it is convenient to move the activity coefficients into the equilibrium constants. For example, the equilibrium constant expression for the dissociation of a weak acid can be written as follows ... 

Calculating the pH of weak acids and bases is more challenging, because they do not ionize completely. So, in a 0.10 M solution of acetic acid, the [H+] concentration is not 0.10 M. We must use the equilibrium constant expression to find the pH of weak acid solutions. So, what is the pH of a 0.10 M solution of acetic acid The K of acetic acid is 1.8 x 10"5. 

Just as with the weak acid example, we can summarize the equilibrium concentration in a table. We need to use a single variable, so let s express the equilibrium concentrations of each species in terms of the hydroxide ion. 

To graph the curves representing [HA] and [A-], a mathematical expression of each as a function of [H+] (a function of the master variable) is needed. The appropriate equation for [HA] is derived by combining the equilibrium constant for dissociation of a weak acid [equation (2.10)] with the mass balance equation [equation (2.13)] to yield... 

For multistep complexation reactions and for ligands that are themselves weak acids, extremely involved calculations are necessary for the evaluation of the equilibrium expression from the individual species involved in the competing equilibria. These normally have to be solved by a graphical method or by computer techniques.26,27 Discussion of these calculations at this point is beyond the scope of this book. However, those who are interested will find adequate discussions in the many books on coordination chemistry, chelate chemistry, and the study and evaluation of the stability constants of complex ions.20,21,28-30 The general approach is the same as outlined here namely, that a titration curve is performed in which the concentration or activity of the substituent species is monitored by potentiometric measurement. 

In this situation, if the hydrogen ion concentration increases (pH becomes lower), the reaction will be driven to the left by mass action (to the original condition), and the proportion of the drug in the nonionized form will increase and, hence, the number of lipid-soluble molecules. For example, if the pICa of a weak acid is 5.0 and it is placed in a medium of pH 4.0, 90 percent will be in the unionized form. Therefore, weak acids are preferentially absorbed in a relatively acidic environment. For a weak base, the equilibrium dissociation constant can be expressed as follows ... 

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