Summary of MQ2 Results: Mean =124 (71 %) Hi = 175. Lo = 32. Your scores will be posted on WebCT

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1 Summary of MQ2 Results: Mean =124 (71 %) Hi = 175 Lo = 32 Your scores will be posted on WebCT 16.5 Strong Acids and Bases Strong Acids Strong Bases 16.6 Weak Acids Calculating K a from ph Using K a to Calculate ph Polyprotic Acids 16.7 Weak Bases Types of Weak Bases 16.8 Relationship Between K a and K b 16.9 Acid-Base Properties of Salt Solutions Acid-Base Behavior and Chemical Structure Factors That Affect Acid Strength Binary Acids Oxyacids Carboxylic Acids Lewis Acids and Bases Hydrolysis of Metal Ions Using K a to Calculate ph Percent ionization is another method to assess acid strength. For the reaction HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) + [H ] eqm % ionization = 100 [HA] 0 Percent ionization relates the equilibrium H + concentration, [H + ] eqm, to the initial HA concentration, [HA] 0. Using K a to Calculate ph The higher percent ionization, the stronger the acid. Percent ionization of a weak acid decreases as the molarity of the solution increases. For acetic acid, 0.05 M solution is 2.0 % ionized whereas a 0.15 M solution is 1.0 % ionized. Example 1: In M solution, acetic acid, abbrev HOAc, is 4.2 % ionized. What is ph and K a? ph = 3.4 K a = 1.8 x 10-5 Example 2: The ph of a M solution of chloroacetic acid is What is K a for this acid? K a = 1.4 x 10-3 Example 3: Calculate the concentrations of all species present in a solution that is initially 0.10 M HOCl, K a = 3.5 x 10-8 [H 3 O + ] = [OCl - ] = 5.9 x 10-5 [HOCl] = x 10-2

2 Notice we can summarize and compare, including use of pk CH 3 COOH 1.8 x ClCH 2 COOH 1.4 x HOCl 3.5 x HCN 4.0 x HNO x Likewise for Weak Bases: NH x (CH 3 )NH x (CH 3 ) 2 NH 7.4 x (CH 3 ) 3 N 7.4 x C 5 H 5 N (pyridine) 1.5 x Example 4: The ph of a household ammonia solution is What is the molarity of ammonia in the solution? K b = 1.8 x 10-5 [NH 3 ] = 0.57 M An indicator is just an organic dye with different colors for the acid (HIn) and its conjugate base (In - ). Consider bromothymol blue HIn + H 2 O = H 3 O + + In - (yellow) (blue) An indicator is just an organic dye with different colors for the acid (HIn) and its conjugate base (In - ). Consider bromothymol blue HIn + H 2 O = H 3 O + + In - (yellow) (blue) or + [ H O ][[ In ] 8 3 Ka = = 7.9x10 [ HIn] HIn [ H O ] + [ ] 3 = [ In ] K a When this is >10 we see blue, when <0.1 we see yellow Polyprotic Acids Polyprotic acids have more than one ionizable proton. The protons are removed in steps not all at once: H 2 SO 3 (aq) H + (aq) + HSO - 3 (aq) K a1 = 1.7 x 10-2 HSO 3 - (aq) H + (aq) + SO 3 2- (aq) K a2 = 6.4 x 10-8 It is always easier to remove the first proton in a polyprotic acid than the second. Therefore, K a1 > K a2 > K a3 etc. Most H + (aq) at equilibrium usually comes from the first ionization (i.e. the K a1 equilibrium).

3 Ascorbic acid, vitamin C, can be abbreviated as H 2 Asc. It has K 1a = 7.9 x 10-5 and K 2a = 1.6 x What is the ph of a 0.10 M solution? 1 st H 2 Asc + H 2 0 = H 3 O + + HAsc - K 1a = 7.95 x 10-5 = x 2 /(0.100-x) x2/0.100 x = check assumption?? 2 nd HAsc - + H 2 O = H 3 O + + Asc -2 i δ -y + y + y K 2a = 1.6 x = y( y)y / ( y) y y = [Asc -2 ] = 1.6 x check [H 3 O + ] Weak bases remove protons from substances. There is an equilibrium between the base and the resulting ions: Weak base + H 2 O conjugate acid + OH - Example: NH 3 (aq) + H 2 O(l) NH + 4 (aq) + OH - (aq) The base dissociation constant, K b, is defined as + - [NH4 ][OH ] K b = [NH ] 3 Types of Weak Bases Bases generally have lone pairs or negative charges in order to attack protons. Most neutral weak bases contain nitrogen. Amines are related to ammonia and have one or more N-H bonds replaced with N-C bonds (e.g., CH 3 NH 2 is methylamine). Anions of weak acids are also weak bases. Example: OCl - is the conjugate base of HOCl (weak acid): ClO - (aq) + H 2 O(l) HClO(aq) + OH - (aq) K b = 3.3 x 10-7 Relationship Between K a and K b We need to quantify the relationship between strength of acid and conjugate base. When two reactions are added to give a third, the equilibrium constant for the third reaction is the product of the equilibrium constants for the first two: Reaction 1 + reaction 2 = reaction 3 has K 1 K 2 = K 3. Relationship Between K a and K b For a conjugate acid-base pair K a K b = K w Therefore, the larger the K a, the smaller the K b. That is, the stronger the acid, the weaker the conjugate base. Taking negative logarithms: pk a + pk b = pk w

4 Consider the following interesting reaction: HF + CN - HCN + F - Two important acid equilibria: HF + H 2 O H 3 O + + F - K a = 6.6 x 10-4 and subtract HCN + H 2 O H 3 O + + CN - K a = 6.17 x HF + H 2 O H 3 O + + F - K a = 6.6 x H 3 O + + CN - HCN + H 2 O K = 1 / K a = 1.6 x giving our original reaction with K = K a / K a = 1.1 x 10 6!!!! Consider the following interesting reaction: HF + CN - = HCN + F - K = K a (HF)/K a (HCN) = 1.1 x 10 6 Which is the stronger acid? the stronger base? Now consider CN - in water: CN - + H 2 O = HCN + OH - K b (CN - ) =? K b = K w /K a = (1.0 x ) / 4.0 x ) = 2.5 x 10-5 = K b (CN - ) Summary of MQ2 Results: Mean =122 (69 %) Hi = 175 (4) Lo = 21 Your scores are posted on WebCT 16.5 Strong Acids and Bases Strong Acids Strong Bases 16.6 Weak Acids Calculating K a from ph Using K a to Calculate ph Polyprotic Acids 16.7 Weak Bases Types of Weak Bases 16.8 Relationship Between K a and K b 16.9 Acid-Base Properties of Salt Solutions Acid-Base Behavior and Chemical Structure Factors That Affect Acid Strength Binary Acids Oxyacids Carboxylic Acids Lewis Acids and Bases Hydrolysis of Metal Ions Acid-Base Properties of Salt Solutions To determine whether a salt has acid-base properties we use: Salts derived from a strong acid and strong base are neutral (e.g. NaCl, Ca(NO 3 ) 2 ). Salts derived from a strong base and weak acid are basic (e.g. NaOCl, Ba(C 2 H 3 O 2 ) 2 ). Salts derived from a weak base and strong acid are acidic (e.g. NH 4 Cl, Al(NO 3 ) 3 ). Salts derived from a weak acid and weak base can be either acidic or basic. Equilibrium rules apply! Acid-Base Properties of Salt Solutions Nearly all salts are strong electrolytes. Therefore, salts exist entirely of ions in solution. Acid-base properties of salts are a consequence of the reaction of their ions in solution. The reaction in which ions produce H + or OH - in water is called hydrolysis. Anions from weak acids are basic. Anions from strong acids are neutral. Anions with ionizable protons (e.g. HSO 4- ) are amphoteric.

5 Consider the effect of dissolving some NaOAc in water: Examples NaOAc(s) Na + + OAc - but OAc - + H 3 O + = HOAc + H 2 0 K = 1/K a H H 2 O = H 3 O + + OH - K w so overall OAc - + H 2 O = HOAc + OH - K b = usual expression but we can rearrange to find that K b = K w /K a = (1.0 x ) / (1.8 x 10-5 ) = 5.6 x 10-5 Calculate [OH - ], ph, and % hydrolysis of a 0.10 M soln of NaCN Contrast: OAc - + H 3 O + = HOAc + H 2 O 1/K a [HCN] = [OH - ] = 1.6 x 10-3 M [CN - ] = ( x 10-3) 0.10 M ph = poh = = % hydrolysis = 1.6 % OAc - + H 2 O = HOAc + OH - K b Similar calculations for a solution of 0.10 M soln of NaOAc would give [OH - ] = 7.5 x 10-6 result of values of K a for HOAc (1.8 x 10-5 ) and HCN (4.0 x ) Consider a 0.10 M solution of NaOAc, K a = 1.8 x 10-5 for HOAc Consider a solution prepared by adding enough NaOCl to water to make 2.00 L of solution with a ph = How many moles of NaOCl were added? K a (HOCl) = 3.5 x 10-8 NaOCl OCL - + H 2 O = HOCl + OH K b = K w / K a = (3.5 x 10-8 ) / (1.0 x 10-14) = 2.86 x 10-7 [OCl - ] i = [NaOCl] i = 0.35 M therefore 0.62 moles of NaOCl were added

6 Acid-Base Behavior and Chemical Structure Factors That Affect Acid Strength Consider H-X. For this substance to be an acid we need: H-X bond to be polar with Hδ+ and Xδ- (if X is a metal then the bond polarity is Hδ-, Xδ+ and the substance is a base), the H-X bond must be weak enough to be broken, the conjugate base, X -, must be stable. Acid-Base Behavior and Chemical Structure Binary Acids Acid strength increases across a period and down a group. Conversely, base strength decreases across a period and down a group. HF is a weak acid because the bond energy is high. The electronegativity difference between C and H is so small that the C-H bond is nonpolar and CH 4 is neither an acid nor a base. Acid-Base Behavior and Chemical Structure Oxyacids Oxyacids contain O-H bonds. All oxyacids have the general structure Y-O-H. The strength of the acid depends on Y and the atoms attached to Y. If Y is a metal (low electronegativity), then the substances are bases. If Y has intermediate electronegativity (e.g. I, EN = 2.5), the electrons are between Y and O and the substance is a weak oxyacid. Acid-Base Behavior and Chemical Structure Oxyacids If Y has a large electronegativity (e.g. Cl, EN = 3.0), the electrons are located closer to Y than O and the O-H bond is polarized to lose H +. The number of O atoms attached to Y increase the O-H bond polarity and the strength of the acid increases (e.g. HOCl is a weaker acid than HClO 2 which is weaker than HClO 3 which is weaker than HClO 4 which is a strong acid). Strengths of H-O-Y Acids (Oxyacids) Acid EN of Y K a pk a HClO x HBrO x HIO x

7 Acid-Base Behavior and Chemical Structure Carboxylic Acids These are organic acids which contain a COOH group (R is some carbon containing unit): O Acid-Base Behavior and Chemical Structure Carboxylic Acids When the proton is removed, the negative charge is delocalized over the carboxylate anion: O O R C OH R C O R C O The acid strength increases as the number of electronegative groups on R increases. Lewis Acids and Bases Brønsted-Lowry acid is a proton donor. Focusing on electrons: a Brønsted-Lowry acid can be considered as an electron pair acceptor. Lewis acid: electron pair acceptor. Lewis base: electron pair donor. Note: Lewis acids and bases do not need to contain protons. Therefore, the Lewis definition is the most general definition of acids and bases. Lewis Acids and Bases Lewis acids generally have an incomplete octet (e.g. BF 3 ). Transition metal ions are generally Lewis acids. Lewis acids must have a vacant orbital (into which the electron pairs can be donated). Compounds with π-bonds can act as Lewis acids: H 2 O(l) + CO 2 (g) H 2 CO 3 (aq)

8 Lewis Acids and Bases Hydrolysis of Metal Ions Metal ions are positively charged and attract water molecules (via the lone pairs on O). The higher the charge, the smaller the metal ion and the stronger the M-OH 2 interaction. Hydrated metal ions act as acids: Fe(H 2 O) 3+ 6 (aq) Fe(H 2 O) 5 (OH) 2+ (aq) + H + (aq) K a = 2 x 10-3 The ph increases as the size of the ion increases (e.g. Ca 2+ vs. Zn 2+ ) and as the charge increases (Na + vs. Ca 2+ and Zn 2+ vs. Al 3+ ).