Acids, Bases and Equilibria

Overview

• Definitions• Strong acids• pH• Water equilibrium• Weak acids• Buffers• Other equilibria• LeChatlier’s Principle

Defining Acids and Bases

• Arrhenius model– Acid – Proton donor – e.g. HCl– Base – Hydroxide donor – e.g. NaOH

• But how about Sodium Carbonate?

Defining Acids and Bases - 2

• Brønsted-Lowery model– Acid – Proton donor – same as Arrhenius– Base – Proton ACCEPTOR

• Aha – so Na2CO3 IS basic!

Na2CO3 + 2HCl 2NaCl + H2CO3

Strong Acids and Bases

• Ionic solids like NaOH; completely form ions in water:

NaOH + H2O Na+ + OH- + H2O

• Covalent molecules like HCl completely IONIZE in water:

HCl + H2O H3O+ + Cl-

• H3O+ is “hydronium” ion – no bare protons

Defining pH

• Remember pH? – Less than 7 = acid– More than 7 = base

• But what does it mean?

• pH is a measure of the concentration of hydronium ion in water pH = - log [H3O+]

Translation: - log

• Suppose we have 0.1M HCl solution

• Since it is fully ionized, we have 0.1M H3O+

• 0.1 = 10-1

• -log (10-1) = 1!

• Therefore pH of this acid solution is 1

Getting the pH of a base

• Even in base, pH measures hydronium ion

• H3O+ and OH- are related by the equilibrium of water

See p. 611

So, what’s equilibrium?

• Second grade analogy – see-saw

• In an equilibrium situation, reactions or changes go both ways

• Hold ice and water at 0o

– Water melts and ice freezes at the same time– “Dynamic” equilibrium

Equilibrium 2

• Form a saturated solution of NaCl– NaCl dissolves;– Same time, NaCl forms new crystals

Water is amphoteric

• H2O + H2O H3O+ + OH-

• Reaction moves to right at same rate as to the left

• Water is being both an acid and a base

• On the other side, “conjugates” are formed– H3O+ is the conjugate acid of H2O

– OH- is the conjugate base of H2O

Water’s “Equilibrium Constant”

• K = [H3O+][OH-]

• K = 10-14

• Square root of 10-14 = 10-7

• [H3O+] = [OH-] = 10-7

• Therefore pH of pure water = 7!

So now to pH of bases:

• Find the pH of 0.01M NaOH• Fully ionized; therefore 0.01M OH-

• [OH-] = 10-2

• K = [H3O+][OH-]• 10-14 = [H3O+] * 10-2

• 10-12 = [H3O+]; pH = 12• OR pK = pH + POH• 14 = pH + 2• 12 = pH

And Weak Acids (or Bases)

• A weak acid is one which is NOT fully ionized• Acetic Acid == HAc (or CH3COOH)• HAc + H2O H3O+ + Ac-

– Acetate ion is the conjugate base of Acetic acid– At equilibrium, HAc is largely NOT ionized

• Because the reaction goes both ways, Acetate can accept a proton: from H3O+ OR from H2O

Ac- + H2O HAc + OH-• Yes, a salt made from a weak acid and a strong base is

basic!

Typical weak acids:

• Acetic acid CH3COOH

• Carbonic acid H2CO3

• Second or third H+ of phosphoric: H2PO4-1,

HPO4-2

So let’s make a “Buffer”

• A buffer is a solution of a weak acid and the strong base salt of its conjugate base:

• Acetic acid and sodium acetate

0.1M 0.1M

CH3COOH + H2O H3O+ + CH3COO-

And let’s add some acid

• First to water:– Add 0.01M HCl to water

– pH becomes 2, right? ([H3O+] = 10-2)

• But add the same acid to the buffer:

0.1M 0.1M

CH3COOH + H2O H3O+ + CH3COO-

0.11M 0.09M

• [H3O+] is almost unaffected! pH stays “same”

Buffer: definition

• A buffer is a solution of a weak acid and its conjugate base OR

• A buffer is a solution of a weak base and its conjugate acid

• Which resists changes in pH when small amounts of strong acid or base are added

• Blood is (or contains) a buffer!

LeChatlier’s Principle

• Notice that a buffer takes advantage of a reversible reaction which shifts away from the species we add: H3O+ or OH-

• LeChatlier said ANY system in equilibrium will shift in such a way as to minimize the effect of a stress applied

Illustration of Principle

Wasn’t that fun?????

• Definitions• Strong acids• pH• Water equilibrium• Weak acids• Buffers• Other equilibria• LeChatlier’s Principle