CHEM 411 Lecture Notes - Lecture 4: Weak Base, Conjugate Acid, Acid Dissociation Constant

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9 Dec 2017

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6. 114 **acid has high ka, thus it has a negative pka. Reverse reaction, the conjugate base accepts a proton from water to reform an acid. Strongest acid, then it is the weakest base. Added base can be neutralized by reacting with acid. Added acid can be neutralized by reacting with base. Buffer: weak acid, weak base conjugate pair. Buffer works best when concentration of base is equal to concentration of acid. Ka of acid=concentration of ha, pka=ph (choose buffer system based on the ph)--> want it to be within 1 ph unit. Ph= pka +log (concentration of a- (base) and ha (acid))

Related Questions

Experiment 18: Buffers Purpose In this experiment you will prepare several buffers and investigate their properties. Discussion A solution of a weak acid and one of its salts, in approximately equal concentrations, resists changes in pH when acid or base is added to it. (Note: The pH of a solution is a measure of the concentration of free H3Ot ion; it does not indicate concentration of combined acid such as HA.) The same is true for a solution of a weak base and one of its salts. A solution with this property is called a buffer (see illustration next page). Bufter systems are extremely important in many chemical and biological systems. For example, your blood contains a buffer system utilizing carbonic acid and bicarbonate ion. The important equilibrium in an acidic buffer is HA (aq) H20 ) A (a)0 () acid conjugate base in which the concentrations of acid and conjugate base are approximately equal. If a small amount of free acid (H30) is added to this solution, it reacts with the conjugate base A (reaction 1 in reverse). The equilibrium constant for this reaction is very large, so the reaction proceeds to near completion, absorbing the added acid. Similarly, added base (OH-) reacts with the acid HA as shown below: 2. HA (aq) + OH-(aq) ATaq) + H2O (1) This reaction also proceeds to completion, effectively absorbing any added base. Since the H30 and OH- concentrations are kept constant, the solution thus resists pH changes when an acid or base is added to it. A similar argument can be made for a basic buffer The actual pH of a buffer can be determined from the value of the Ka and the concentrations of the acid and conjugate base obtained from an initial-change- equilibrium table. However, a shortcut for buffers is available, and is shown below. We begin by writing and rearranging the equilibrium constant expression for the acid dissociation. If the -log of both sides of equation (3) is taken, a relationship involving pH is then obtained. H30IA] HA] 3. HA HAI -loglH30'1- -log Ka A- 141

For the dissociation reaction of a weak acid in water,

HA(aq)+H2O(l)âH3O+(aq)+Aâ(aq)

the equilibrium constant is the acid-dissociation constant, Ka, and takes the form

Ka=[H3O+][Aâ][HA]

Weak bases accept a proton from water to give the conjugate acid and OHâ ions:

B(aq)+H2O(l)âBH+(aq)+OHâ(aq)

The equilibrium constant Kb is called the base-dissociation constant and can be found by the formula

Kb=[BH+][OHâ][B]

When solving equilibrium-based expression, it is often helpful to keep track of changing concentrations is through what is often called an I.C.E table, where I. stands for Initial Concentration, C. stands for Change, and E. stands for Equilibrium Concentration. To create such a table, write the reaction across the top creating the columns, and the rows I.C.E on the left-hand side.

Initial (M)Change (M)Equilibrium (M)A+ BâAB

Part A

Aspirin (acetylsalicylic acid, C9H8O4) is a weak monoprotic acid. To determine its acid-dissociation constant, a student dissolved 2.00 g of aspirin in 0.600 L of water and measured the pH. What was the Ka value calculated by the student if the pH of the solution was 2.60?

Express your answer numerically using two significant figures.

Hints

Ka =

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You used the initial concentration of aspirin in the Ka expression. Instead, you need to use the equilibrium concentration. Consider that the amount of aspirin that reacted is equal to the amount of H₃O⁺ produced. How much is left over?

Part B

A 0.100 M solution of ethylamine (C2H5NH2) has a pH of 11.87. Calculate the Kb for ethylamine.

Express your answer numerically using two significant figures.

Hints

Kb =

Part A - Calculate the molar solubility in water

Mg(OH)2 is a sparingly soluble salt with a solubility product, Ksp, of 5.61Ã10?11. It is used to control the pH and provide nutrients in the biological (microbial) treatment of municipal wastewater streams.

Based on the given value of the Ksp, what is the molar solubility of Mg(OH)2 in pure H2O?

Express your answer with the appropriate units.

molar solubility =

2.41Ã10^?4 M

Calculate the molar solubility in NaOH

Based on the given value of the Ksp, what is the molar solubility of Mg(OH)2 in 0.120 M NaOH?

Express your answer with the appropriate units.

molar solubility =

3.90Ã10^?9 M

Part C: Calculate how many times more soluble Mg(OH)2 is in pure water

Based on the given value of the Ksp, calculate the ratio of solubility of Mg(OH)2 dissolved in pure H2O to Mg(OH)2 dissolved in a 0.120 M NaOH solution.

A buffer is a mixed solution of a weak acid or base, combined with its conjugate. Note that this can be understood essentially as a common-ion problem: The conjugate is a common ion added to an equilibrium system of a weak acid or base. The addition of the conjugate shifts the equilibrium of the system to relieve the stress of the added concentration of the common ion. In a solution consisting of a weak acid or base, the equilibrium shift also results in a pH shift of the system.

It is the presence of the common ion in the system that results in buffering behavior, because both added H+ or OH? ions can be neutralized.

Part D

What is the pH change of a 0.240 M solution of citric acid (pKa=4.77) if citrate is added to a concentration of 0.175 M with no change in volume?

CHEM 411 Lecture Notes - Lecture 4: Weak Base, Conjugate Acid, Acid Dissociation Constant

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