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Acid Base Equilibria

Acid Base Equilibria. Dr. Harris Ch 20 Suggested HW : Ch 20: 5, 9, 11*, 18*, 19*, 21 , 29 **, 35, 56**, 59, 66 * Use rule of logs on slide 10 ** Use K a and K b tables in ch 20 text. . Introduction.

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Acid Base Equilibria

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  1. Acid Base Equilibria Dr. Harris Ch 20 Suggested HW: Ch 20: 5, 9, 11*, 18*, 19*, 21, 29**, 35, 56**, 59, 66 * Use rule of logs on slide 10 ** Use Ka and Kb tables in ch 20 text.

  2. Introduction • In chapter 10, we defined an acidas a substance capable of producing H+(aq) ions in solution. • For strong acids, this equilibrium lies so far to the rightthat you can make the mathematical assumption that [HX] ≈ 0. • These “simple” acids are known as Arrhenius acids. The Arrhenius definition of an acid is restricted to aqueous acid solutions. HXaq H+aq + X-aq MOHaq M+aq + OH-aq • An Arrhenius base is a substance that produces OH-(aq) ions in aqueous solution.

  3. Bronsted-Lowry Acids and Bases • Until now, H+(aq) has been used to represent the species formed when an acid is dissolved in water. • In the Bronsted-Lowry definition of acids and bases, an acid is a H+ (proton) donor.A base is a proton acceptor. • Here, we define the hydronium ion, H3O+(aq). This species is formed when a water molecule accepts a proton. In a chemical equation, H+ and H3O+ are interchangeable. • For example: • In the example above, HA is a Bronsted acid, and water acts as a Bronsted base.

  4. H2O(l) + H2O(l) H3O+(aq) + OH−(aq) Autoionization of Water • Pure water contains very low concentrations of hydronium and hydroxide ions that arise from the equilibrium • In the reaction above, one water molecule transfers a proton to another. Thus, water can act as BOTH a Bronsted acid and Bronsted base • The equilibrium constant expression for the autoionization of water is: • At 25oC, the value of Kw is 1 x 10-14 M2. • In pure water, the concentrations of hydronium and hydroxide are equal.

  5. Strong Acids and Bases (RECAP) • By now, you should know the strong acids and bases. If not, make sure that you learn them. They are listed in the table to the right.

  6. Conjugate Acid-Base Pairs • In any acid-base equilibrium, the forward and reverse reactions involve proton transfers. • When an acid, HA, loses its H+(aq), the A-(aq) left behind can behave as a base. Likewise, when a base accepts the H+(aq), it can now act as an acid. • An acid and a base such as HA and A- that differ only in the presence of a proton are called a conjugate acid-base pair. Every acid has a conjugate base, and visa-versa.

  7. Relative Strengths of Acids and Bases • The stronger an acid, the weaker the conjugate base. Same for bases and conjugate acids. • A strong acid completely transfers its protons into water. Its conjugate base has a negligible tendency to be protonated. • A weak acid partially dissociates, and so it exists as a mixture of molecules and ions. The conjugate of a weak acid/base is a weak base/acid HCl(aq) + H2O(l) H3O+(aq) + Cl−(aq) x The position of equilibrium is always favored in the direction that forms the weaker acid/base.

  8. pH • The rates of many chemical reactions depend on the concentration of [H3O+]. However, [H3O+] can range vastly. • It is very difficult to plot such a large range of values, so we use the pH scale. In terms of [H3O+]: • Recall that for pure water, we have the equilibrium: • And in pure water: • So in pure water, = 10-7 M, and pH = 7.00

  9. pOH and the PH scale • In terms of basicity, we can introduce a quantity pOH. • For a basic solution, you can not directly calculate pH.You must calculate pOH first. Referring back to the following relationship: • Therefore: * at 25oC

  10. pH scale and Important Notes • For this class, we will assume the temperature of all aqueous solutions to be exactly 25oC so that the previous expression holds true. • It is important to know the following: An acidicsolutionis defined as one in which [H3O+] > [OH-] (pH < 7) A neutral solution is defined as one in which [H3O+] = [OH-] (pH = 7) A basicsolution is defined as one in which [OH-] > [H3O+] (pH > 7) Also, please note that H+ and H3O+ are the same, and may be used interchangeably in the book. this will come in handy

  11. Example • What is the pH of a 0.040M solution of HClO4at 25oC? • Perchloric acid is a strong acid, so we can assume complete dissociation. I 0.040M 0 0 C -(~0.040) +(~0.040 M) +(~0.040 M) E ~ 0 ~0.040M ~ 0.040M • The dissociation of the acid may also be written as: 0.040M 0.040M • It is a matter of preference.

  12. Example • What is the pH of a 0.028M solution of NaOH at 25oC? • NaOH is a strong base. We can therefore assume that it completely dissociates, so [OH-] = 0.028 M. We must calculate pOH first. E ~0M ~.028M ~.028M

  13. Acid Dissociation Constant, Ka • Let’s consider acetic acid, CH3COOH, a weak acid. The equilibrium expression of this dissociation is defined by Ka: CH3COOH (aq)+ H2O(L) H3O+(aq) + CH3COO-(aq) • The size of Ka describes mathematically how readily an acid dissociates. The stronger the acid, the higher the value of Ka.

  14. Example • The pH of a 0.40 M formic acid solution (HCOOH) is 2.08. What is the concentration of H3O+? What is Ka? What % of the acid dissociates? • Formic acid is a weak acid. Weak acids do not fully dissociate HCOOH (aq)+ H2O(L) H3O+(aq) + HCOO-(aq) I 0.40 M 0 0 C -x +x +x E 0.40-x x x

  15. Group Example Calculate the pH of a 0.050 M acetic acid solution given that Ka = 1.8 x 10-5M 1. We set up the ICE table to determine the concentration of H3O+ CH3COOH (aq) + H2O(L) H3O+(aq) + CH3COO-(aq) 2. Solve for x. We see that Ka is very small because CH3COOH is a weak acid. 3. Calculate pH

  16. Method of Successive Approximations • Weak acids tend to dissociate quite poorly, so the percentage of dissociation is very small. Hence, Ka is very small as well. Let’s revisit the previous example. • Since x is so small: • Therefore: Use this method when K < 10-4

  17. pKa • Ka values range over many orders of magnitude. Therefore, you may find it more convenient to use the quantity pKa • pKa represents the pH at which 50% of the acid is deprotonated. • As Ka increases, pKa decreases. Thus, a low pKa indicates a stronger acid.

  18. Amines • Similar to acids, we can describe the protonationof bases. For example, lets look at the equilibrium involving ammonia in water. • The lone electron pair of electrons on the Nitrogen allows it to accept an H+ ion (proton). • Ammonia falls under a class of bases known as amines, which are composed of a central nitrogen atom with 3 bonding domains. • Amines react with acids to form ammonium salts(positively charged central N atom with 4 bonding domains).

  19. Kb and pKb • We can use the base protonation constant Kb, or pKb to describe the degree of the dissociation of a base. • Higher Kb means stronger base Example: Determine the pH of a 0.75 M hydroxylamine, HONH2(aq) solution given that Kb = 8.7 x 10-9 M HONH2 (aq) + H2O(L) HONH3+ (aq) + OH-(aq) We can write out the equilibrium expression showing the base protonation.

  20. Example, contd. Set up ICE table Solve x THIS IS A BASE. That means that you must calculate pOH FIRST. Then, determine pH.

  21. Acidic and Basic Salts • Suppose we dissolve NaF(s) in water. The result would be an aqueous solution of Na+(aq) and F-(aq). You would expect this solution to be neutral. • However, F-(aq) is the conjugate base of a weak acid, HF acid conjugate acid base conjugate base • Therefore, a solution of F-(aq) ions will be weakly basic. strong base negligible acid weak acid weak base • Na+(aq) is a cation of a strong base, NaOH. Thus, it has negligible acidity and does not affect the pH. So it is considered a neutral cation.

  22. Acidic and Basic Salts Neutral cations are the cations of strong bases. Neutral anions are the anions of strong acids. Basic anions are the conjugate bases of weak acids. Acidic cations are the conjugate acids of weak bases, and hydrated metal ions, as shown on pg 760.

  23. Polyprotic Acids • A polyprotic acidis an acid than can donate more than one H+. • For example, sulfuric acid, H2SO4, is a strong, diprotic acid. Thus, it can undergo two deprotonations: • We can assume 100% dissociation of the H2SO4 in step 1). However, the remaining HSO4- is a much weaker acid. EACH PROTON IS SUBSTANTIALLY HARDER TO REMOVE THAN THE PREVIOUS. • Therefore, the degree of dissociation of the weaker acid is determined from its Ka

  24. Polyprotic Acids • Ka values of polyprotic acids are labeled Ka1, Ka2, and so forth to denote each deprotonation. • The table to the left shows pKa values of common polyprotic acids. • Remember, pKa= -log Ka

  25. Example 1.) We first have the initial deprotonation. Because sulfuric acid is a strong acid, you can assume that all of it is dissociated at equilibrium. • Calculate the pH of a 0.10 M H2SO4(aq) solution. E 2.) Now, we have the second deprotonation. Use the concentrations of H3O+ and HSO4- from above as your initial values. Use pKa values from the table to determine the final concentration of [H3O+]

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