Ch 9: Acids, Bases and Salts

Suggested Problems:

2, 6, 10, 12, 28-44, 82, 94-100, Bonus: 118

Acids and Bases in Aqueous Solution

• Think back to Chapter 4– Acid

• Definition: A substance that produces H+ when dissolved in H2O

• Examples: HCl, HNO3, H2SO4

– Base• Definition: A substance that produces OH- when

dissolved in H2O – Examples: KOH, NaOH, NH3

Common Acids

Sulfuric Acid H2SO4

Phosphoric acid H3PO4

Acetic acid CH3CO2H

Nitric acid HNO3

Hydrochloric acid HCl

Common Bases

Sodium Hydroxide NaOH

Calcium hydroxide Ca(OH)2

Magnesium hydroxide Mg(OH)2

Ammonia NH3

Common Acids and Bases

• Acids– H2SO4

– HCl– H3PO4

– HNO3

– CH3CO2H

• Chemical Formulas Begin with H

• Formulas Contain CO2H

• Bases– NaOH– Ca(OH)2

– Mg(OH)2

– NH3

• Chemical Formulas Contain OH

Arrhenius Acids and Bases

• 1903 Chemistry Nobel Prize – Barely Awarded Ph.D.

• Technicality issue with Arrhenius acid definition– H+ is very reactive

Updated Definitions

• Arrhenius acids – Substances that produce H3O+ when

dissolved in H2O

• Arrhenius bases– Substances that produce OH- when dissolved

in H2O

• What if the reactions are not in H2O?

Brønsted-Lowry Acids and Bases

Separately

developed the

same theory

pertaining to acids and bases in

1923Johannes Brønsted Thomas Lowry

Brønsted-Lowry Acid

• Definition: any substance that is able to give a hydrogen ion H+, to another molecule or ion– Proton donor

• Not limited to reactions in H2O– Do not have to create appreciable [H3O+]

– NaOH(s) + HCl(aq)

CH3CO2H(aq) + H2O(l) H3O+(aq) + CH3CO2

-(aq)

NaCl(aq) + H2O(l)

Brønsted-Lowry Acids• Different acids can donate different

numbers of H+

Acid # of Acidic H+ Terminology

HCl

H2SO4

H3PO4

CH3CO2H

1

2

3

1

monoprotic

diprotic

triprotic

monoprotic

Brønsted-Lowry Bases• Definition: a substance that accepts H+

from an acid – Proton Acceptor

• Not limited to reactions in H2O– Do not have to create appreciable [OH-]

– NH3(g) + HCl(g)

CH3CO2H(aq) + H2O(l) H3O+(aq) + CH3CO2

-(aq)

B-L Base

NH4Cl(s)

Identify the following as Brønsted-Lowry Acids, Bases or Neither

• HCN• AlCl3• H2CO3

• CH3CO2-

• Mg2+

• CH3NH3+

Acid

Base

Neither

Acid

Acid

Neither

Brønsted-Lowry Acids and Bases

• Summary– Acid-Base reaction is one in which a proton is

transferred• Use B or B- to represent bases• Use HA to represent acids

• Using the symbols B, B- and HA write two general acid-base chemical reactions

B + HA BH+ + A-

B- + HA BH + A-

Brønsted-Lowry Acids and Bases

• Consequence of Brønsted-Lowry Definition– What are species BH+, BH, and A- ?

– BH+

– BH– A-

– Acid-Base reactions are reversible• K is often large, resulting in a single preferred

direction

Acid

Acid

Base

Conjugate Acid – Base Pairs

• Definition: A pair of compounds whose formula differ only by one proton.

– After an acid donates a proton, the remaining species turns into a conjugate base (CB).

– After a base accepts a proton, the resulting species turns into a conjugate acid (CA).

HF(aq) + H2O(l) H3O+(aq)

+ F-(aq)

- H+

+H+

Acid Base ConjugateAcid

ConjugateBase

+ H+

Conjugate Acid – Base Pairs Example

NH3(g) + H2O(l) NH4+

(aq) + OH-

(aq)

+ H+

+H+

Base Acid ConjugateAcid

ConjugateBase

- H+

Conjugate Acid – Base Pairs Example

HF(aq) + H2O(l) H3O+(aq)

+ F-(aq)

Water as Both an Acid and a Base

NH3(g) + H2O(l) NH4+

(aq) + OH-(aq)

• A substance that can react as an acid or a base is

called amphoteric

Acid

Base

Common Acid-Base Reactions

• Neutralization reaction– Acid with a metal hydroxide

• Salt: anion of the acid with the cation of the base

– HCl(aq) + KOH(aq) KCl(aq) + H2O(l)

– Why is this called a neutralization reaction?• Net Ionic Equation

– H+(aq) + OH-

(aq) H2O(l)

Products are always salt and H2O

Acid-Base Reactions

• Write the balanced chemical equation for the reaction of sulfuric acid with magnesium hydroxide.

Acid-Base Reactions• Acid with bicarbonate and carbonate ions

– Bicarbonate ion • HCO3

-

– H+(aq) + HCO3

-(aq) [H2CO3(aq)] CO2(g) + H2O(l)

– Carbonate ion• CO3

2-

– 2H+(aq) + CO3

2-(aq) [H2CO3(aq)] CO2(g) + H2O(l)

• Products of these reactions are salt, CO2 and H2O

Acid-Base Reactions

• Write the balanced chemical reaction for nitric acid with baking soda (sodium bicarbonate).

Acid-Base Reactions• Acid with Ammonia

– Products for this general reaction are ammonium salts

• NH3(aq) + HNO3(aq) NH4NO3(aq)

• Write the balanced chemical reaction for ammonia with sulfuric acid.

Challenge Problem

• An over the counter antacid has NaAl(OH)2CO3 as the active ingredient. – How many grams of this antacid are required

to nuetralize 15.0 mL of 0.0955 M HCl?

The Self Ionization of Water

• H2O is amphoteric

• But what if you have just pure water?

H2O(l) + H2O(l) ⇌ H3O+(aq) + OH-

(aq)

• This equilibrium is governed by the equilibrium constant Kw

Equilibrium Constant

• Equilibrium constant (K) is equal to the concentration of the products divided by the reactants

• aA + bB cC + dD

ba

cd

B][[A]

C][[D]K

[x] = concentration of species X in molarity

KW

H2O(l) + H2O(l) ⇌ H3O+(aq) + OH-

(aq)

L

mol 10 x 1.0 ]OH][OH[K 14--

3w

Strong Acids

• Strong acids give away all of their hydrogen ions

• For example, HCl is a strong acid, and when HCl dissolves in water:

HCl H+ + Cl-

Weak Acids

• Weak acids do not give away their H+ ions, and are in equilibrium with their ionized form

• Most acids are weak acids• For example, acetic acid is a weak acid,

and when HC2H3O2 dissolves in water:

HC2H3O2 ⇌ H+ + C2H3O2-

Strong Bases

• A strong base will give away all of its hydroxide ions (OH-)

• For example, NaOH is a strong base, and when NaOH dissolves in water:

NaOH Na+ + OH-

Weak Bases

• To think about weak bases you must think in terms of a proton acceptor not in terms of OH-. (Brønsted-Lowry Base)

• Weak bases accept some H+.• Again as with weak acids there is an

equilibrium present.

Strong/Weak Acids and Bases

• The description of strong/weak has nothing to do with concentration

• Concentration is independent of it being strong or weak.

• Concentration is a measure of the amount of moles per liter

• You can have low concentrations of Strong Acid and Bases

pH and pOH • pH informs a person about whehter or not a

solution is acid or basic

• pH = -log[H+]• pOH = -log[OH-]

• pH of 7 is nuetral• pH less than 7 is acidic• pH greater than 7 is basic

pH + pOH = 14

pH and pOH calculation examples

• Determine the concentration of H3O+ and OH- from the following pH values

• pH = 9.0• pH = 3.0• pH = 11.0

pH and pOH Calculations

• This course will deal only with non-equilibrium acids and bases when calculating pH or pOH

• Therefore the concentration of H+ and OH- will be able to be determined from the stoichiometry of the formula.

• For example– What is the pH of a solution of 0.10 M HCl?– What is the pH of a solution of 0.20 M NaOH?

Salts• Definition: a substance composed of the cation

of a base with the anion of the acid• Need to discuss Equivalent units

– This term is terribly misused by the medical and biological profession

• A equivalent is the quantity of material necessary to deliver one unit of chemical reactivity– It makes no sense outside of the context of a

chemical reaction!• However, in blood analysis,

Equivalents = moles x charge on ion

Equivalents Example

• Determine the number of equivalents in the following:

• 0.10 mol of NaCl• 0.10 mol of CaCl2

– Only consider either the positive or negative charges not both and the origination of the species is also important

Equivalent Example

• A sample of blood serum contains 0.139 eq/L of Na+ ion. Assume the Na+ comes from dissolved NaCl, and calculate the number of equivalents, number of moles and number of grams of NaCl in 250 mL of the serum.

Titration Calculation• A 25 mL sample of vinegar (which

contains acetic acid) is titrated with 0.100 M NaOH. If 6.75 mL of NaOH are required, what is the molarity of the acetic acid in vinegar?

25 mL of vinegar

0.100 M NaOH

Titration Example

• A 25.0 mL sample of H2SO4 solution requires the addition of 16.3 mL of 0.200 M NaOH solution to reach the equivalence point. What is the concentration of the acid?

Buffers

• A pH buffer is a solution that resists changes in pH

• A pH buffer must contain a weak acid (HA) and its conjugate base (A-)

HA + OH- H2O + A-

A- + H+ HA

Buffer Example

• Carbonic acid and bicarbonate are important blood buffers