Unit 9

Acids, Bases, & Salts

Acid/Base Equilibrium

Properties of Acids

• sour or tart taste

• strong acids burn; weak acids feel similar

to H2O

• acid solutions are electrolytes

• acids react with most metals to release H2

• acids cause indicators to change color

– Acids turn litmus red

Properties of Bases

• bases taste bitter

• basic solutions do not burn

• basic solutions feel smooth and slippery

• basic solutions are also electrolytes

• bases usually do not react with metals

• bases also cause indicators to change

color

– Bases turn litmus blue

Three definitions

• Arrhenius

– Deals with H+ and OH-

• Bronsted Lowry

– Deals with protons

• Lewis

– Deals with pairs of electrons

Arrhenius Definition

• Arrhenius defined an acid as a substance

that ionizes in water to produce H+ ions

– Ex. HCl

– Any of our acids with H+ in the beginning

• Arrhenius defined a base as a substance

that ionizes in water to produce OH- ions

– Ex. NaOH

– Any of the metal hydroxides

Bronsted Lowry Definition

• NH3 is known to turn litmus blue, but it

does not have hydroxide ion

– Needed a new definition

– Substances do not need to be in water

• A B/L acid is any substance that can

donate a proton (H+)

• A B/L base is any substance that can

accept a proton (H+)

Conjugate pairs

HC2H3O2 + H2O → C2H3O2- + H3O

+

B/L acid B/L base conjugate base conjugate acid

NH3 + H2O → NH4+ + OH-

B/L base B/L acid conjugate acid conjugate base

Amphoteric Substances

• Notice in the examples, that H2O can act

as a B/L acid and as a B/L base.

• It is know as an amphoteric substance.

• Additional examples include

– HSO3-

– HCO3-

– H2PO4-

– HPO42-

These are the ions that result

when polyprotic acids release

H+ ions one at a time. Notice

that they can accept an H+ or

they can release an H+.

Lewis definition

• This is the most broad of the definitions

• A Lewis acid is an electron pair acceptor.

• A Lewis base is an electron pair donor.

• Example Lewis acid/base reaction….

Ion product constant for H2O

• Remember [H+][OH-] = 1.0 x 10-14 = Kw

– So if you know [H+] you can find [OH-] and

vice versa

Ex. If [H+] = 9.3 x 10-4 M, what is [OH-](9.3 x 10-4 M)[OH-] = 1.0 x 10-14

[OH-] = 1.1 x 10-11 M

pH or pOH

• pH = -log[H+] or pOH = -log[OH-]

• pH + pOH = 14

Ex. What is the pH of a solution that contains

4.9 x 10-9 M OH-? Find pOH first

pOH = -log (4.9 x 10-9)

= pOH 8.31

Since pH + pOH = 14, pH = 5.69

What is the [H+] of a solution that is pH4.82?

pH = -log [H+], rearranging for [H+]

[H+] = 10-pH

= 10-4.82

= 1.5 x 10-5 M H+

What is the [OH-] of a solution that is pH8.56?

If pH is 8.58, pOH is 5.42

[OH-] = 10-pOH

= 10-5.42

= 3.8 x 10-6 M OH-

You must understand that…

• When [H+] = [OH-], the solution is neutral, and the pH = 7

• When [H+] > [OH-], the solution is acidic, and the pH <7

• When [H+] < [OH-], the solution is basic, and the pH > 7.

• Increasing pH means decreasing [H+]…fewer H+

ions floating in solution….less acidic.

• Decreasing pH means increasing [H+]…more H+

ions floating in solution….more acidic.

Calculate the number of moles of LiOH in 8.75 mL of .150

M LiOH solution. What is the pH? pOH?

Moles = M x L

= (.150 M)(.00875 L) = .00131 moles LiOH

pOH = -log [OH-]

= - log .150 M

= pOH .824

pH = 14 - .824

= pH13.2

1 Li 6.94

1 O 16.00

1 H 1.01

23.95 g

1 H 1.01

1 Cl 35.45

36.46 g

Weak Acids

Weak Bases

Acid Equilibrium problems

• Things to know

– All of our acids will be monoprotic, that is, give

off only one hydrogen ion.

– x will represent the amount of acid that

dissociates.

• Therefore x also represents the [H+]

– We will use ICE to determine concentrations

for the Ka expression

– We can use HA to represent the acid and A-

to represent the conjugate base.

Ka example

pH = -log(.0019)

= pH 2.72

Polyprotic acids

• Polyprotic acids are weak acids (except for H2SO4 which is a strong acid).

– They release their H+ ions one at a time.

– Each has its own Ka

– Ex. H3PO4

H3PO4 → H+ + H2PO4- Ka = 7.5 x 10-3

H2PO4- → H+ + HPO4

2- Ka = 6.2 x 10-8

HPO42- → H+ + PO4

3- Ka = 4.8 x 10-13

Kw revisited

• Water comes to equilibrium with its ions

according to the following reaction…

H2O(l) ↔ H+(aq) + OH-

(aq)

Kw = [H+][OH-] = 1.0 x 10-14

Kw = KaKb pKa + pKb = 14

Strong Acids/Bases

Strong Acids Strong Bases

HCl LiOH

HBr NaOH

HI KOH

HNO3 RbOH

HClO4 CsOH

H2SO4 Ba(OH)2

Sr(OH)2

Strong Acids/Base description

• Strong acids and bases completely

dissociate in water, therefore no Ka or Kb

– The dissociations do not reverse.

• Oxoacids are acids that contain oxygen.

– The greater the number of oxygen atoms

attached to the central atom in an oxoacid,

the stronger the acid.

• That’s because increasing the number of oxygen

atoms that are attached to the central atom

weakens the attraction that the central atom has

for the H+ ion.

Strong Acid/Base calculations

• Since these acids and bases completely

dissociate in water, the final concentration

of H+ ions is the same as the original

concentration.

• So you can always find the pH of a strong

acid solution directly from its

concentration.

– What is the pH of .20 M HCl?

pH = -log(.20) =

Titration

• When an acid and a base are mixed, a neutralization reaction occurs.– Acid + base → salt + water

• Neutralizations reactions are generally performed by titration, where a base of known concentration is slowly added to an acid (or vice versa)

• The progress of a neutralization reaction can be shown in a titration curve.

• The equivalence point is the point in the titration when exactly enough base has been added to neutralize all the acid that was initially present.

• An indicator will be used to mark the equiv. pt.

• Strong acid/strong base titration.

– Equivalence point at pH7

• Notice the shape of the curve if a strong

acid is added to a strong base

– The equivalence point is also pH7

• Weak acid vs strong base

– The equivalence point is > pH7

• Weak base vs strong acid

– Notice the equivalence point is < pH7

Which indicator to use?• The indicator needs to change color close

to the equivalence point.

pH of dissolved salts• If a salt is composed of the conjugates of a

strong base and a strong acid, its solution will be neutral. (NaCl)

• If a salt is composed of the conjugates of a weak base and a strong acid, its solutions will be acidic. (NH4Cl)

• If a salt is composed of the conjugates of a strong base and a weak acid, its solution will be basic. (NaC2H3O2)

• If a salt is composed of the conjugates of a weak base and a weak acid, the pH of its solution will depend on the relative strengths of the conjugate acid and base of the specific ions in the salt. (NH4C2H3O2)

Anhydrides

• An acid anhydride is a substance that

combines with water to form an acid.

– Generally nonmetal oxides.

• Ex. CO2 + H2O → H2CO3

• A basic anhydride is a substance that

combines with water to form a base.

– Generally metal oxides.

• Ex. Na2O + H2O → 2NaOH

What to know for the test?

• Find pH from either [H+] or [OH-]

• Find [H+] from pH or pOH

• Find the # of grams of a base to make a

solution of a certain pH

• Use titration info to find molar mass acid

• Ka problems

– Find pH - find Ka from pH

– Find concentrations

– Find % ionization