Health Science Chemistry I, CHEM-1180
Chapter 8
Lecture Notes, Bettelheim 10th ed
Acids and Bases
Important Definitions

Arrhenius Acid

Arrhenius Base

Hydrogen Ion

Hydronium Ion

Acid Dissociation
Dissociation of Acids in aqueous solution to produce H3O+
Eg. HCl + H2O → H3O+ + + Cl(Compare to old viewpoint, HCl → H+ + + Cl-)
Dissociation of Bases to release OH- in solution
Eg.
NaOH dissolving and dissociating in water
However, most bases do not have hydroxide ion in their formulas. Instead, many bases produce OH- by
removing a proton from water. These are Brønsted Bases.
B: + H2O → B:H+ + OH- or, more specifically
NH3 + H2O → NH4+ + OHTo be discussed later.
Reminder: Aqueous ammonia vs. ammonium hydroxide name.
Acid and Base Strengths
Define Strong Acid
Reaction of Strong acids with water
List of six strong acids (learn names and formulas)
Define Weak Acid
Reaction of weak acids with water
Examples of common weak acids: Carbonic acid, Acetic Acid, Phosphoric Acid, Hydrofluoric acid
Large range of weak acid strengths
Define Strong Base
Reactions of strong bases with water
List of four strong bases (learn names and formulas)
Define Weak Base
Reaction of weak bases with water
Examples of common weak bases: most heavy metal hydroxides and ammonia
Brønsted-Lowry Acid Base Theory
An acid is a proton donor (as in Arrhenius Theory)
A base is a proton acceptor
When an acid loses a proton it becomes its conjugate base (Examples)
When a base gains a proton it becomes its conjugate acid
Rule: Strong acids have weak conjugate bases. Weak acids have strong conjugate bases. Why?
Consult list of relative acid-base strengths (Table 8.2)
Brønsted-Lowry acids and base reactions become the transfer of a proton from an acid to a base to form a new
acid-base pair.
Examples: Reaction of HCl and water; Reaction of acetic acid and ammonia; ammonium ion and water
Define: monoprotic, diprotic and triprotic acids.
Examples: carbonic acid, phosphoric acid.
Amphiprotic substances such as bicarbonate ion, monohydrogen phosphate ion, etc.
The Position of Equilibrium in Acid-Base Reactions
Rule: A Brønsted-Lowry acid-base reaction always results in a weaker acid-base pair
To determine the equilibrium position of an acid-base system;
1.
Identify the acid on each side of the equation
2.
Identify the base on each side of the equation
3.
Compare the relative strengths of the two acids
4.
compare the relative strengths of the two bases
5.
Draw unequal equilibrium arrows to show direction of equilibrium
Examples: The acetic acid-ammonia system. The ammonium ion-sodium hydroxide system
Acid Ionization Constants
For the general reaction of a weak acid with water, the law of mass action gives
Keq = [A-][H3O+]/[HA][H2O]. Since [H2O] is constant at 55.5 M, it is combined with the constant Keg to give
Keq [H2O] = Ka = [A-][H3O+]/[HA]. Simplify H3O+ to H+
Finally, Ka = [A-][H+]/[HA]
Ka values of most weak acids are well known. Typical values range from 0.1 to 10-15.
Define pKa. why needed?
Simple examples of pKa and Ka interconversions
Reactions of Acids and Bases
Neutralization Reactions
Acid plus Base gives a salt plus water.
This reaction can take place in solution or with solid bases such as lime, Ca(OH)2 (s)
Acid plus Metal Reactions
Iron plus sulfuric acid
Aluminum plus nitric acid
Acid plus metal hydroxides (neutralization)
Hydrochloric acid plus barium hydroxide solution
Acids plus metal oxides
HCl plus Fe2O3 (s)
Acids plus Bicarbonates
HCl plus sodium bicarbonate
Acids plus Carbonates
Sulfuric acid plus calcium carbonate
Acids plus Ammonia
Nitric acid plus aqueous ammonia (also amines)
The Self-Ionization of Water
In pure water or in any aqueous solution:
H2O + H2O  H3O+ + OH- or more simply, H2O  H+ + OHKeq = [H3O+][OH- ]/[H2O]
Keq x [H2O] = Kw = [H3O+][OH- ] or Kw = [H+][OH- ]
Kw is called the ion product of water and has a value of 1.0 x 10-14 at 25 ºC
The equation Kw = [H+][OH- ] can be used to calculate [H+] or [OH- ] whenever the other is known.
A solution is neutral when [H+] = [OH-] = 1.0 x 10-7 M.
A solution is acidic when [H+] > [OH-] or [H+] >1.0 x 10-7 M.
A solution is basic when [H+] < [OH-] or [OH-] >1.0 x 10-7 M.
 Calculate [H+] if [OH-] = 2.5 x 10-10 M. Is the solution acidic or basic?
 Calculate [OH-] if [H+] = 3.91 x 10-8 M. Is the solution acidic or basic?
The pH Concept
Why is pH needed to express solution acidity? Wide range of possible [H+] values.
Define pH = -log [H+]. Significance?
Simple pH calculations show the logarithmic definition is more convenient.
Acidic solutions have pH between 0 and 7; basic solutions have pH between 7 and 14.
Examples of [H+] to pH and [H+] to pH without a calculator. Then use log and antilog keys.
What about fractional pH values?
How to calculate log of 4.5 x 10-4.
How to calculate [H+] from fractional pHs.
Define pOH = -log[OH-]. Why needed?
Whenever just one quantity is known, can now calculate pH, pOH, [H+] and [OH-] and identify whether a
solution is acidic or basic.
 Examples of pH, pOH, [H+] and [OH-] interconversions
Back to pKa values: Convenience of pKa vs. Ka values
 Examples of pKa to Ka interconversions.
Acid and Base Properties of Salt Solutions
Some cations and anions react with water (undergo hydrolysis) to produce [H+] or [OH-], affecting solution pH.
Rules:
1.
Anions of strong acids do not react with water to produce OH- in solution. Ex. Cl-.
2.
Anions of weak acids do react with water to produce OH- in solution. Ex. F-.
3.
Cations of strong bases do not react with water to produce H+ in solution. Ex. Na+ and Ca+2.
4.
Cations of weak bases do react with water to produce H+ in solution.
Ex. Ammonium ion is a weak acid that dissociates to produce H+.
Ex. Aluminum ion creates H+ out of its coordination sphere to produce H+.
Two ways to envision the aluminum hydrolysis reaction.
Ex: Predict whether a solution of each salt will be acidic, basic or neutral.
NaCl, K2S, (NH4)2SO4, Al(NO3)3, Ca(CH3COO)2, FeCl2, etc.
Titration
Definition of titration, specifically acid-base titration.
Define equivalence point and endpoint of a titration.
Titration often involves the use of high-precision glassware. Use one molarity to calculate another.
Must know the pertinent acid-base neutralization equation.
Ex: 3KOH (aq) + H3PO4 (aq) → K3PO4 (aq) + 3H2O (l)
Calculate the concentration of a KOH solution if 23.27 mL is needed to titrate 25.00 mL of 0.138 M H3PO4.
Titration tools and techniques:

Buret

Volumetic pipet

Erlenmeyer flask

Pipetting techniques

Use of soluble indicator dyes to track the titration. How to choose the proper indicaor?

Use of pH meter to track the titration.
Buffers
Define buffer
Why are buffers needed in chemistry and biochemistry?
How do buffers work?
Chemically, what is a buffer?
Define Buffer Capacity
Derive the Henderson-Hasselbalch Equation from the equation for the Ka of a weak acid.
Blood pH is maintained by three separate buffer systems, the carbonate buffer, the phosphate buffer and
amphiprotic proteins.
Ka = [H+][A-]/[HA]
What happens to pH when pKa varies?
What happens to pH when [A-] and [HA] vary?
What is the “characteristic pH” of a buffer? How to choose the correct buffer?
Phosphate buffer calculations (dihydrogen phosphate ion pKa = 7.21)
Carbonate buffer calculations (carbonic acid pKa = 6.37)