1. Chapter 16 : Acid-Base Equilibria Created by Lauren Querido

2. Table of Contents 16.1 Review 16.2 Brønsted-Lowry Acids and Bases 16.3 Autoionization of Water 16.4 pH Scale 16.5 Strong Acids and Bases 16.6 Weak Acids 16.7 Weak Bases 16.8 Relationship Between K a and K b 16.9 Acid-Base Properties of Salt Solutions 16.10 Acid-Base Behavior and Chemical Structure 16.11 Lewis Acids and Bases

3. 16.1 Review Acids –Sour in taste –Litmus paper turns red Bases –Bitter, slippery –Litmus paper turns blue When acids and bases mix, their properties disappear!

4. Arrhenius Acids and Bases Svante Arrhenius (1880) –In aqueous solutions: Acids will increase the concentration of H + ions when dissolved in water. Bases will increase the concentration of OH - ions when dissolved in water.

5. 16.2 Brønsted-Lowry Acids and Bases 1923 Brønsted and Lowry made a more general definition –Brønsted-Lowry Acid is a substance that can transfer a proton. It must have a hydrogen atom that can be lost as H +. –Brønsted-Lowry Base is asubstance that can accept a proton. Must have a nonbonding pair of electrons to gain a H + ion.

6. Conjugate Acid-Base Pairs Conjugate base- Removal of proton from the acid Conjugate acid- Addition of proton to the base

7. Relative Strengths of Acids and Bases The stronger the acid, the weaker its conjugate base. The stronger the base, the weaker its conjugate acid. 1. Strong acids completely transfer protons to water. 2. Weak acids partly dissociate in aqueous solutions and exist as a mixture of acid molecules and component ions. 3. Negligible acidity contain Hydrogen but do not demonstrate acidic behavior. Ex: CH 4 Position of equilibrium favors transfer of proton from stronger acid to stronger base.

8. 16.3 Autoionization of Water Ion product of water – –1.0 x 10 –14 = [H + ] [OH - ] –This is used to calculate concentrations of H + and OH -. If [H + ] = [OH - ], than neutral equation If [H + ] > [OH - ], than acidic equation If [H + ] < [OH - ], than basic equation =

9. 16.4 The pH Scale pH = -log [H + ] pH of 7 is neutral Acidic solution 0 < pH < 7 Basic solution 14 > pH > 7 Other p scales are –pOH = -log [OH - ] –pOH + pH = -log K w = 14.0

10. Examples on the pH Scale

11. Measuring pH A pH meter consists of a pair of electrodes connected to a meter which pH is generated when placed in the solution. An acid-base indicator turns a color if placed in acid or base. Ex: litmus paper

12. 16.5 Strong Acids and Bases Strong Acids –7 most common strong acids are HCl, HBr, HI, HNO 3, HClO 3, HClO 4, and H 2 SO 4 –In acidic reactions, equilibrium lies entirely to the right side. –Completely dissociates –Example: HNO 3 => H + + NO 3 -

13. Strong Bases Most common strong bases are ionic hydroxides of alkali metals (1A) and heavier alkaline earth metals (2A). Examples: LiOH, RbOH, CsOH, NaOH, KOH, and Ca(OH) 2, Sr(OH) 2, and Ba(OH) 2. Other strong bases react with water to form OH- such as Na 2 O, CaO. Also, anions O 2-, H -, and N 3- are stronger bases than OH- and therefore remove a proton from H 2 O. –Example: N 3- + H 2 O => NH 3 + 3OH -

14. 16.6 Weak Acids A weak acid only partially ionizes in aqueous solutions. General weak acid equation –HX  H + + X - where H is Hydrogen –Many weak acids contain some Hydrogen atoms bonded to carbon atoms and oxygen atoms (organic compounds). –K a is the acid dissociation constant. –The larger the value of K a, the stronger the acid.

15. Calculating K a from pH Use and ICE box! Sample exercise –A student prepared a.10 M solution of formic acid and measures its pH which was 2.38. A) calculate Ka for formic acid B) what percentage of the acid is ionized in the.10M solution?

16. Answer a) HCHO 2  H + + CHO 2 - Ka = [H + ][CHO 2 - ] [HCHO 2 ] pH= -log[H + ] =10 –2.38 = 4.2 X 10 -3 M Ka = [ 4.2 X 10 –3 ][ 4.2 X 10 –3 ] [.10 ] 1.8 X 10-4 = [ 4.2 X 10 –3 ][ 4.2 X 10 –3 ] [.10 ] b) Percent Ionization = Concentration of H+ Initial concentration of component = 4.2% HCHO 2  H+H+ CHO - I.10 M0 M C-4.2 X 10 –3 +4.2 X 10 –3 E.10 - 4.2 X 10 –3 +4.2 X 10 –3

17. Using K a to Calculate pH The best way to explain this is by an example. Calculate the pH of a.30 M solution of acetic acid at 25 o C. (K a = 1.8 X 10 -5 ) So… HC 2 H 3 O 2  H + + C 2 H 3 O 2 - Ka = [H + ][C 2 H 3 O 2 - ] = 1.8 X 10 -5 [HC 2 H 3 O 2 ] What now?

18. HC 2 H 3 O 2  H+H+ C2H3O2-C2H3O2- I.30 M0 M C-x+x E.30-xxx Ka = (x)(x) = 1.8 X 10 -5 (.30 – x) Either do the quadratic equation or in this case you can take out x in the denominator. [H + ] = x = 2.3 X 10 -3 pH = -log 2.3 X 10 -3 = 2.64

19. Polyprotic Acids Polyprotic acids have more than one ionizable Hydrogen atom. Example: H 2 SO 3  H + + HSO 4 - HSO 4 -  H + + SO 3 2- The second K a (K a 2 ) is much smaller than K a1 because it is easier to remove the first proton.

20. 16.7 Weak Bases Weak base + water => conjugate acid + hydroxide ion K b is the base-dissociation constant (equilibrium in which base reacts when H 2 O to form conjugate acid and OH - ion). Types of weak bases: 1.Neutral substances that have atoms with a non-bonding pair of electrons that can serve as a proton acceptor. –Most of these contain amines, N-H which is sometimes replaced with a bond between C or N Ex: NH 2 CH 3 2.Anions of weak acids –Ex: ClO - + H 2 O  HClO + H + –ClO - is the weak base

21. 16.8 Relationship Between K a and K b Reaction 1 + reaction 2 = reaction 3 Which leads to K 1 x K 2 = K 3 Which leads to K a x K b = K w K w is the ion-product constant for water –K w = 1 x 10 -14 As the strength of the acid increases, the strength of the base decreases and visa-versa. pK a + pK b = pK w = 14.00

22. 16.9 Acid-Base Properties of Salt Solutions Hydrolysis is the process at which ions react with water and produce H + or OH - X- + H 2 O  HX + OH - Anions of strong acids do not influence pH –Ex: NO 3 - Anions that still have ionizable protons are amphoteric –Ex: HSO 3 - from H 2 SO 4 Most cations (except 1A elements and Ca +2, Sr +2. Ba +2 ) act as weak acids in solution.

23. Predicting the pH of a Solution 1. Salts derived from a strong acid and a strong base makes a neutral pH (pH of 7). NaOH + HCl => NaCl + H 2 O 2. Salts derived from a strong base and a weak acid will yield a pH of above 7 because the anion hydrolyzes to produce OH- ions and the cation does not hydrolyze. NaOH + HClO => NaClO + H 2 3. Salts derived from a weak base and a strong acid will result in a pH that is below 7 because the cation hydrolyzes to produce H+ ions and the anion does not hydrolyze. Al(OH) 3 + 3HNO 3 => Al(NO 3 ) 3 + 3H 2 O

24. 4. Salts derived from a weak base and a weak acid will yield a pH that is dependant on the constant value of the constant dissociations (K a and K b ). if the base is more basic than the acid is acidic, then the solution will have a pH that is greater than 7. If the acid is more acidic, than the pH will be less than 7. NH 4 + + CN -  NH 4 CN NH 4 + K a = 5.6 X 10 -10 CN - K b = 2.0 X 10 -5 Therefore, the pH of NH 4 CN is greater than 7

25. 16.10 Acid-Base Behavior and Chemical Structure Factors that effect acid strength –If H-X bond is polarized (X is more electronegative) the H acts as a proton acceptor. –Non-polar bonds (CH 4 ) produce neutral solutions. –Weaker bonds dissociate more easily than very strong bonds. –HF is a weak acid because of this. –The greater the stability of the conjugate base, the weaker the acid. –Ultimately, there are three factors effecting acid strength: Polarity of H-X bond Strength of H-X bond Stability of conjugate base, X -

26. Binary Acids Binary acids are composed of Hydrogen and a non-metal. –Ex: HCl, HF, H 2 S, etc. The more polar the bond,the stronger it is The weaker the bond, the stronger the acid. Strength of the bond decreases (acidity increases) as the element increases in size or moves down a group. Acid strength increases (acidity decreases) moving from left to right

27. Group 4A5A6A7A Period 2 CH 4 No acid or base properties NH 3 Weak base H 2 O ------- HF Weak acid Period 3 SiH 4 No acid or base properties PH 3 Weak base H 2 S Weak acid HCl Strong acid Increasing acid strength Increasing base strength

28. Oxyacids Oxyacids are acids with an OH group is bound to a central atom. –Example: H 2 SO 4 –

29. OH - Bonding To determine if an OH group acts as an acid or base, consider this: If Y is a metal than sources of OH - behave as bases. If Y is a non-metal than the compound will not readily lose the OH - ion. –The electronegativity will increase and so will the acidity. The increasing number of Oxygen atoms stabilizes the conjugate base and thus increases the strength of the acid.

30. Oxyacid Rules of Thumb 1.Oxyacids that have the same number of OH groups and the same number of Oxygen atoms, acid strength increases with increasing elecronegativity of the central atom Example: HClO > HBrO > HIO (> = more acidic) 2. For oxyacids with the same central atom, acid strength increases with increasing number of Oxygen atoms that are attached. Example: HClO < HClO 2 < HClO 3 < HClO 4 ( < less acidic)

31. Carboxylic Acid Carboxylic acids are organic compounds. -COOH is the functional group -R is either a Hydrogen or Carbon based group –If an extra Oxygen is added than it stabilizes the conjugate base and increases the acidity. –If conjugate base has resonance structures, it spreads the negative charge evenly over the compound. –Acid strength of carboxylic acid increases as the number of electronegative atoms increase.

32. 6.11 Lewis Acids and Bases G.N. Lewis proposed this: Lewis Acids have an incomplete octet of electrons. Function as electron pair acceptors Lewis Bases act as electron pair donators

33. Hydrolysis of Metal Ions Hydration is a process when when metals attract unshared electron pairs of water molecules. –The metal acts as Lewis acid –The water acts as Lewis base –Ex: Fe(H 2 O) 6 +3  Fe(H 2 O) 5 (OH) 2+ + H + –So, general equation M(H 2 O) n c  M(H 2 O) n-1 (OH) c-1 + H +