Chapter 16 Acid-Base Equilibria

Dissociation of water Autoionization or autoprotolysis Ion-product constant Autoprotolysis constant constant

K w = [H + ][OH - ] = 1.0x When [H + ] = [OH - ] neutral. Doesn’t usually happen. As one increases, the other decreases; the product must equal 1.0x When [H + ] > [OH - ]acidic [OH - ] > [H + ]basic

H + is a proton with no electrons. In water:        H HOH Hydronium ion

Bronstead-Lowry Acid-Base Acid - Can donate a proton Base - Can accept a proton *Doesn’t have to be in H 2 O. Can be in other solvents.

Conjugate Acid-Base Pairs conj base conj acid

The stronger an acid, the weaker its conjugate base. The weaker an acid, the stronger its conjugate base.

pH scale pH = -log [H + ] Remember K w = (1x10 -7 )(1x10 -7 ) = 1.0x pH = -log [H + ] = -log (1x10 -7 ) pH = 7 (neutral) [H + ]pH acidic> 1.0x10 -7 < 7.00 basic 7.00

You can also speak in terms of [OH - ] pOH = -log [OH - ] = 14 - pH Because pH + pOH = -log K w = 14

Measure pH by pH meter Acid-base indicators Litmus red = pH < 5 blue = pH > 8 Figure 16.7 shows several acid-base indicators and their ranges

Strong Acids and Bases Strong electrolytes Completely ionize HA + H 2 O  A - + H 3 O + Bases form hydroxides in solvent

In H 2 O, Alkali metal hydroxides Alkaline earth metal Hydroxides (except Be) Many are insoluble Also, substances that will abstract a H + from H 2 O. O 2- + H 2 O  2OH - Na 2 O or CaO would do this. O 2-, H -, N 3- bases that would do this.

Weak acids Only partially ionize Acid dissociation constant

Larger K a means stronger acid. ex. N O C - O - H O = 0.020M solution pH = 3.26 ? K a pH = -log [H + ] = 3.26 [H + ] = 5.50x10 -4

N O C - OH O = N O C - O O = + H +  HA A-A- H+H+ 1:1

Can calculate pH in same manner if you have K a and concentration of solution. Let’s use niacin again. N O C - OH O = N O C - O O = + H +  HA A-A- H+H+

** Simplifying Assumption ** x is very very small compared to 0.010M sooooooooo, ignore x in denominator

pH = -log [H + ] x = [H + ] = 3.9x10 -4 pH = 3.41 What percent of niacin molecules ionized?

Polyprotic Acids ex. H 2 SO 4 H 3 PO 4 H 2 SeO 4 H 2 SO 4  H + + HSO 4 - K a1 = 1.7x10 -2 HSO 4 -  H + + SO 4 2- K a2 = 6.4x10 -8 K a1 always larger than K a2 If K a1 / K a2  10 3, can estimate pH by K a1 only.

Weak Bases ex. Amines “an organic substituted ammonia” ammonia NH 3 N H HHN H CH 3 H methyl amine

N H CH 3 + H 2 O  HN H CH 3 + OH - H H ClO - + H 2 O  HClO + OH - K b = 3.3x10 -7 Can use this in the same manner in which you used K a. Anions of weak acids

K a and K b How are they related?

1) 2) 3) When two reactions are added together, the equilibrium constant for the third reaction is given by the product of equilibrium constants of equations 1 and 2.

K 1 x K 2 = K 3 rxn 1 rxn 2 rxn 3

Special Case K a x K b = K w For conjugate acid-base pairs.

Bond polarity and Bond strength effect on Acid-base behavior: In binary acids  polarity(across a row)  acidity  bond strength(in a group)  acidity  stability of conj. base  acidity

Metal hydrides are basic or show no acid/base properties in H 2 O. Nonmetal hydrides are acidic or show no acid/base properties in H 2 O (except NH 3 ) Acidity increases moving down a group.

Oxyacids HOS O O O H Have unprotonated and protonated oxygens. YOH H 3 PO 4 As electronegativity of Y increases, acidity increases. As number of unprotonated oxygens increases, acidity increases (effect of formal charge and oxidation number) Ex. HClO, HClO 2, HClO 3, HClO 4

Carboxylic Acids R C OH OCOOH = Carboxyl group R = H or an organic group. The more electron withdrawing R is, the greater the acidity (this stabilizes anion and weakens O-H bond) ex. CHC H H O O H Acetic acid K a = 1.8x10 -5 CFC F F O O H Trifluoroacetic acid K a = 5.0x10 -1

Lewis Acids and Bases This is a completely different definition for acid/base chemistry than what you have seen thus far!!! Lewis acid = electron pair acceptor Lewis base = electron pair ‘donor’ Not giving them away, just has them available to ‘share’.

H + Bronstead-Lowry acid also a Lewis acid H + electron pair acceptor OH - Electron pair donor Lewis base also Bronstead-Lowry base

B H H H BH 3 not a Bronstead- Lowry acid, but it’s a Lewis acid Incomplete Octet N H H H Lewis Base has an electron pair available to attack an area that is e - deficient

Transition metal ions are often Lewis Acids. They have vacant d orbitals. (s and p also) H H O O = C = OCan be a Lewis Acid because e - density around the C is bound in just 2 directions.

H H O = = O O C H H O O O C H H O O O C Carbonic acid Hydrolysis of metal ions Metal ions have positive charge so they attract the lone e - pair on H 2 O molecules

6 of these H O H Fe 3+ H O H H O H H O H H O H H O H H O H Fe 3+ Because the metal is (+), e - density of H 2 O moves toward the metal. When this happens, there is less e - density in water’s O-H bonds, so H + can come off easier…  pH will drop.

The higher the charge density of the metal ion, the greater the acidity of its aqua complex.