1. A.P. Chemistry Chapter 14 Acid- Base Chemistry

2. 14.1 Arrhenius Acid- an acid is any substance that dissolves in water to produce H + (H 3 O + ) ions Base- a base is a substance that dissolves in water to produce OH - ions Bronsted- Lowry Acid- an acid is a proton (H + ) donor Base- a base is a proton acceptor Hydronium Ion- the ion that foms when water accepts a proton (H + )H 3 O + General reaction that occurs when an acid is dissolved in water HA (aq) + H 2 O(l)  H 3 O + (aq) + A - (aq) Conjugate base- everything that remains of an acid molecule after a proton is lost Conjugate acid- the molecule which is formed when the proton is transferred to the base Conjugate acid- base pair consists of two substances related to each other by the donating and accepting of a single proton.

3. Equilibrium expression for a general reaction when an acid is dissolved in water would be K a = [H 3 O + ][A - ]=[H + ][A - ] [HA][HA] (water not included; pure liquid) K a is called the Acid Dissociation Constant

4. Example: Write the formula of the conjugate base of H 2 SO 4. (a conjugate base differs from its acid by the lack of a proton: HSO 4 - ) Example: For the reaction HSO 4 - (aq) + HCO 3 - (aq)  SO 4 2- (aq) + H 2 CO 3 (aq) identify the acids and bases for the forward and reverse reactions (forward: HSO 4 - is proton donor, acid; HCO 3 - is proton acceptor, a base.) identify the conjugate acid-base pairs

5. 14.2 Acid Strength Strong Acid- one in which the equilibrium lies far to the right; it has a large value of K a. Strong acids also yield weak conjugate bases. Weak Acids- one in which the equilibrium lies far to the left; it has a small value of K a. Weak acids yield relatively strong conjugate bases. See Table 14.1, p. 662 Example: The strengths of the following acids increase in the order HCN < HF < HNO 3. Arrange the conjugate bases of these acids in order of increasing base strength. (CN - > F - > NO 3 - )

6. Diprotic Acid- an acid having two acidic protons; Examples: H 2 SO 4 Oxyacids- the acidic proton is attached to an oxygen atom, Examples: Organic Acids- those acids with a carbon atom backbone; commonly contain the carboxylic group (-COOH, or ) Examples: Monoprotic acid- acids having one acidic proton. Examples:

7. Amphoteric- can behave either as an acid or a base; ex., water. Autoionization- involves the transfer of a proton from one water molecule to another to produce a hydroxide ion and a hydronium ion. 2H 2 O(l)  H 3 O + (aq) + OH - (aq) Ion-product Constant K w (or dissociation constant for water)- always refers to the autoionization of water. Using the above reaction K w = [H 3 O + ][OH - ] = [H + ][OH - ] At 25 o C, [H + ] = [OH - ] = 1.0 x 10 -7 M therefore, at 25 o C, K w = (1.0 x 10 -7 M) 2 = 1.0 x 10 -14 mol 2 L -2

8. It is important to recognize the meaning of K w : The product of [H + ] and [OH - ] must always equal 1.0 x 10 -14. This lends itself to three possibilities: A neutral solution, where [H + ] = [OH - ] An acidic solution, where [H + ] > [OH - ] A basic solution, where [H + ] < [OH - ] K a x K b = K w = 1.0 x 10 -14 Example:The H + concentration in a solution is 5.0 x 10 -5 M. What is the OH - ion concentration? (2.0 x 10 -10 M; don’t forget unit when asked for a concentration!)

9. 14.3 pH The pH scale is used to determine the strength of an acid or a base. Traditionally it ranges from 0-14. Acids have pH 7 and 7 on the scale is considered to be neutral. pH is defined as pH = -log[H 3 O + (aq)] or pH = -log[H + (aq)] In the case of strong acids and bases dissociation is complete and therefore the concentration of the H+ ion or OH-ion can be determined directly from the stoichiometric ratio in the balanced equation and the concentration of the acid or the base. Therefore we can simply plug in the numbers to the equation above, or, use the fact that, @ 25 o C logK w = log[H+] + log[OH-] = 7 + 7 = 14 = pH + pOH And then use pOH = -log[OH - (aq)]

10. ***Significant Figures for logarithms: the number of decimal places in the log is equal to the number of significant figures in the original number.***

11. Example: The OH - ion concentration in a certain ammonia solution is 7.2 x 10 -4 M. What is the pOH and pH? (Answer: 3.14; 10.86) Problem: What is the pH and pOH of a 0.001 M solution of HNO 3 ? Problem: What is the pH of a 0.045 M solution of HCl? Problem: What is the pH and pOH of a 0.02 M solution of H 2 SO 4 ? Example: The pH in many cola-type soft drinks is about 3.0. How many times greater is the H + concentration in these drinks than in neutral water?(10 4 or 10,000X)

12. (Working backwards: If you know pH and want to find concentration, use the antilog.) [H+] = 10 -pH [OH-] = 10 -pOH Another useful relationship in calculations is pK a = -logK a and pK b = -logK b and pK w = pH + pOH Problem: What is the concentration of a HCl solution for which the pH is 4.32?

13. 14.4 Calculating the pH of Strong Acid Solutions See Table for List of strong and weak acids and bases. Know the List! Example: What is the pH of a 0.1 M HCl solution? Of a 0.0001 M HCl solution? Example: What is the pOH of a 0.1 M NaOH solution? The pH? Example: What is the pOH of a 0.00001 M NaOH solution? The pH?

14. 14.5 Calculating the pH of a Weak Acid Solution VERY IMPORTANT!!! Step 1- List the major species in the solution Step 2- Choose the species that can produce H+, and write a balanced equation for the reaction producing H+ Step 3- Using the values of the equilibrium constant for the reactions you have written, decide which equilibrium will dominate in producing H+. Step 4- write the equilibrium expression for the dominant equilibrium. Step 5- List the initial concentrations of the species participating in the dominant equilibrium Step 6- Define the change needed to achieve equilibrium; that is, define x. Step 7- Write the equilibrium concentrations in terms of x. Step 8- Substitute the equilibrium concentrations into the equilibrium expression Step 9- Solve for x the “easy” way; that is, assume [HA] o - x = [HA] o (assume x is negligible) Step 10- Using the 5% rule, verify is that approximation is valid Step 11- calculate [H + ] and pH.

15. 5% Rule:x/[HA] x 100% < 5% then the approximation is acceptable Percent Dissociation: % dissociation = amount dissociated x 100% Initial concentration For a given weak acid, the percent dissociation increases as the acid becomes more dilute. For solutions of any weak acid HA, [H + ] decreases as [HA] o decreases, but the percent dissociation increases as [HA] o decreases.

16. Example: Determine the pH of a of a 0.10 M solution of CH 3 COOH. K a for acetic acid is 1.8 x 10 -5 Determine the percent ionization of acetic acid in the example.

17. Problem: Calculate the concentrations of H +, F -, and HF in a 0.31 M HF solution. (K a for HF is 7.1 x 10 -4 ) Calculate the % ionization. (Answer: [H + ] = [F - ] = 1.5 x 10 -2 M [HF] = 0.31 M - 0.015 M = 0.30 M; 4.8%)

18. 14.6 Bases Strong Base – one which dissociates completely in water and has a large Kb value. General reaction for a base B and water is B(aq) + H 2 O(l)  BH + (aq) + OH - (aq) Base acid conjugate conjugate acid base The equilibrium constant, K b, for this equation is K b = [BH + ][OH - ] [B] Weak bases have small values of K b. (see Table 14.3, p. 685)

19. Example: What is the H + concentration in 0.0025 M NaOH? (4.0 x 10 -12 M) Example: What is the pH of a 0.010 M C 5 H 5 N (pyridine) solution?(8.62)

20. 14.7 Polyprotic Acids Polyprotic acids- acids which can furnish more than 1 proton. (See Table 14.4, p. 689) Characteristics of Weak Polyprotic Acids (p. 694) Typically, successive K a values are so much smaller than the first value, that only the first dissociation step makes a significant contribution to the equilibrium concentration of H +. This means that the calculation of the pH for a solution of a typical weak polyprotic acid is identical to that for a solution of a weak monoprotic acid. Sulfuric acid is unique in being a strong acid in its first dissociation step and a weak acid in its second dissociation step. For relatively concentrated solutions of sulfuric acid (1.0 M or higher) the large concentration of H + ions from the first dissociation step represses the second dissociation step (which can then be neglected as a contributor of H + ions). For dilute solutions of sulfuric acid, the second step does make a significant contribution, and the quadratic equation must be used to obtain the total H + concentration.

21. Example: Calculate the concentrations of H 2 A, HA -, A 2-, and H + in a 1.0 M H 2 A solution.

22. 14.8 Acid-Base Properties of Salts Salt- another name for an ionic compound Salts that contain highly charged metal ions produce an acidic solution. The metal ion becomes hydrated which then causes the solution to become acidic. See Table 14.6, p. 700.

23. Example: Write net ionic equations to show which of the following ions hydrolyze in aqueous solution? NO 3 -, NO 2 -, NH 4 + Example: calculate K b for NO 3 - (2.2 x 10 -11 ) Example: calculate K a for NH 4 + (5.6 x 10 -10 )

24. Example: Determine the pH of a 0.10 M solution of NH 4 Cl. (K b for NH 3 is 1.8 x 10 -5 ) Example: Predict whether the following aqueous solutions will be acidic, basic, or neutral. KI (neutral) NH 4 I (acidic) CH 3 COOK (basic)

25. 14.9 The Effect of Structure on Acid-Base Properties There are two main factors that determine whether a molecule containing an X-H bond will behave as a Bronsted- Lowry acid: the strength of the bond and the polarity of the bond. Increased polarity and high electron density typically lends to large K a values (strong acids). (homework problem; has been on AP exam!!) HF? Example: Which is the stronger acid? HCl or HBr? (HBr) HCl or H 2 S (HCl) Example: Which is the stronger acid? HClO 3 or HBrO 3 ? (HClO 3 ) H 3 PO 3 or H 3 PO 4 (H 3 PO 4 )

26. 14.10 Acid-Base Properties of Oxides (Important in Equation Writing!!) A compound containing the H-O-X group will produce an acidic solution in water if the O-X bond is strong and covalent. If the O-X bond is ionic, the compound will produce a basic solution in water. (homework problem; has been on AP exam!!) Acidic oxides- a covalent oxide dissolves in water, and an acidic solution forms. Basic oxides- an ionic oxide dissolves in water, and oxide has a great affinity for H +, causing basic solutions.

27. 14.11 Lewis Acid-Base Model (encompasses the B-L model, but the reverse is not true!) (Be able to draw the Lewis structure of this reaction) Acid- an acid acts as an electron pair acceptor Base- a base acts as an electron pair donor Example: identify the Lewis acids and bases in each of the following reactions: Ag + (aq) + Cl - (aq)  AgCl(s)(acid: Ag + ) BF 3 (g) + NF 3 (g)  F 3 N-BF 3 (s) (acid: BF 3 ) SO 2 (g) + H 2 O(l)  H 2 SO 3 (aq) (acid: SO 2 )