Acid-Base Equilibria

Acids Bases Sour taste React with active metals to release hydrogen gas Change the color of indicators Bitter taste Feel slippery Change the color of indicators

Arrhenius Acid/Base Theory According to Arrhenius, acids are substances that produce H + (H 3 O + ) ions in water. Bases are substances that produce OH - ions in water.

Why does water affect acids? A water solution of HCl produces: HCl(g)  H + (aq) + Cl - (aq) Acidic solutions are formed when an acid transfers a proton to water. H + + H 2 O  H 3 O +

Brønsted-Lowry Acids/Bases Acids are substances capable of donating a proton. Bases are capable of accepting a proton. B/L theory applicable to reactions that do not occur in water Can include gas phase reactions NH 3 (g) + H 2 O(l)  NH 4 + (aq) + OH - (aq)

An acid and a base which differ by only one proton are conjugate acid/base pairs. Acid HI HNO ` PH 4 + HSO 4 - Conj. Base HSO 4 – H 2 O NH 2 - F - Conj. Acid

The more readily an acid gives up a proton, the less readily the conjugate base will accept a proton.  The stronger the acid, the weaker its conjugate base and the weaker the acid, the stronger its conjugate base. HCN is a weaker acid than HF. The conjugate base CN - is a stronger base than F -.

Diethylammonium ion (CH 3 ) 2 NH 2 + is a weak acid. Determine the conjugate base: Is the conjugate base stronger or weaker than the conjugate base of HCl?

Strong Acids Ionize completely to form ions HCl HBr HI HNO 3 H 2 SO 4 HClO 4

Autoionization of Water Proton transfer from one water molecule to another H-O-H + H-O-H  H 3 O + + OH - H 2 O  H + + OH - K = [H + ][OH - ] H 2 O K w = [H + ][OH - ] = 1.0x10 -14

[H 2 O] >> 55 M and remains constant  not written in K If solution is: neutral [H + ] = [OH - ] acid [H + ] > [OH - ] base [H + ] < [OH - ] [H + ] = 2x10 -5 M [OH - ] = 3x10 –9 M [OH - ] = 1 x M

pH Scale pH = -log[H + ] Acidic solution pH < 7 Basic solution pH > 7 Neutral solution pH = 7 Number of decimal places in pH = number of significant figures in concentration. 1.0x pH = 12.00

Strong Acids Strong electrolytes React completely with water to form H + pH of a strong acid equals –log of hydrogen concentration

Weak Acids Weak electrolytes [molecules] greatest at equilibrium H atoms bound to C do not ionize while H atoms bound to O will ionize HX(aq)  H + (aq) + X - (aq) K a = [H + ][X - ] [HX]

0.10M HCOOH (formic acid) pH = 2.38 Calculate K a and percent dissociation

In the reaction of a strong acid with metal the conductivity remains constant since all acid molecules ionize. But, for a weak acid, conductivity will fluctuate since degree of ionization increases As [A] decreases. Calculate %HF in 0.10 M HF M HF

Polyprotic Acids Release more than one proton in H 2 O H 2 SO 3  H + (aq) + HSO 3 - (aq) K a1 HSO 3 - (aq)  H + (aq) + SO 3 2- (aq) K a2 K a1 = 1.7 x K a2 = 6.4 x 10 -8

If K a constants differ by a factor of 10 3, consider K a1 only. In other words, treat the acid as if it were monoprotic. The solubility of CO 2 at 25 o C and 0.10 atm equals M. All of the dissolved CO 2 is as H 2 CO 3. What is the pH of a M solution of H 2 CO 3 ?

Strong Bases Most common strong bases are heavy Group IIA and all Group IA. Determine the pOH of a basic solution by: pOH = -log [OH - ] pH + pOH = 14 Determine the pH of a M solution of Ba(OH) 2.

Formation of Strong Bases Stronger bases than water are able to remove an H + ion from water: O 2- + H 2 O(l)  2OH - (aq) H - + H 2 O(l)  H 2 (g) + OH - (aq) N 3- + H 2 O(l)  NH 3 (aq) +3OH - (aq)

Weak Bases Weak base + water  acid + OH - NH 3 + H 2 O  NH OH - K b = [NH 4 + ][OH - ] [NH 3 ]

Calculate the [OH - ] in a 0.15 M solution of NH 3. (K b = 1.8 x )

Amines Weak nitrogen bases N-C bonds Due to lone pair on N, it is able to extract a proton H 2 N-CH 3 + H 2 O  [H 3 NCH 3 ] + + OH -

Anions of Weak Acids When sodium salts dissolve in water, the Na + ion merely acts as a spectator ion. The reaction occurs between the remaining anion and water. C 2 H 3 O H 2 O  HC 2 H 3 O 2 + OH - Calculate the pH of a M solution of NaClO.

Relation Between K a and K b NH 4 + (aq)  NH 3 (aq) + H + (aq) NH 3 (aq) + H 2 O(l)  NH 4 + (aq) + OH - (aq) NH 4 + and NH 3 are conjugate pairs. K a = [NH 3 ][H + ]K b = [NH 4 + ][OH - ] [NH 4 + ] [NH 3 ] Add Reaction 1 and Reaction 2.

H 2 O  H + + OH - Multiply K a x K b Ka Ka x K b = KwKw K a and K b are inversely related, as acid strength increases (K a ), base strength decreases (K b ) since product must equal K w.

Calculate the K b for F - and the K a for NH 4 +. K b = 1.5 x K a = 5.6 x Which of the following has the largest K b ? NO 2 -, PO 4 3-, N 3 - PO 4 3-

Acid-Base Properties of Salts Completely ionized Nearly all salts are strong electrolytes Acid/base properties due to hydrolysis of cations and anions Strong acids and bases produce ions that do not hydrolyze

 Anions of weak acids: NO 2 - (aq) + H 2 O(l)  HN or O 2 (aq) + OH - (aq)  Cations of weak bases: NH 4 + (aq) + H 2 O(l)  NH 3 (aq) + H 3 O + (aq)  Anions of polyprotic acids act as either proton donors or acceptors depending on the value of K a or K b. Predict whether Na 2 HPO 4 will be acidic or basic.

pH of Salts Depends on parent acid and base. If acid and base are both strong, salt is neutral. If acid is strong and base is weak, salt is acidic. If acid is weak and base is strong, salt is basic. If parent acid and base are both weak, salt pH depends on value of K a and K b. List the following in order of increasing pH: Co(ClO 4 ) 2 RbCN Sr(NO 3 ) 2 KC 2 H 3 O 2

Acid/Base Character Any molecule containing H can act as a potential acid, but bond must be polarized. Very strong bonds are less ionized than weak ones.  HF is a weak acid but acidic character increases with increasing atomic number making the remaining elements in the halogen family strong acids.

OxyAcids Acids in which OH groups and possibly additional oxygen atoms are bound to a central atom. Acid strength increases with increasing EN of Y. HIO < HBrO < HClO

In a series of acids with the same Y, acid strength increases with increasing oxidation number. HClO < HClO 2 < HClO 3 < HClO 4

Lewis Acids/Bases According to Lewis, an acid is an electron pair acceptor. Acids will generally have incomplete octets. A base is an electron pair donor. Bases will tend to have lone pairs. BF 3 + NH 3  F 3 BNH 3

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