1. Chapter 16: Acids and bases
Chemistry 1062: Principles of Chemistry II
Andy Aspaas, Instructor

2. Acids and bases
Three concepts to describe behaviors of acids and bases
Arrhenius concept: ionization of water molecules
Bronsted-Lowry concept: donation and acceptance of protons
Lewis concept: donation and acceptance of electron pairs

3. Arrhenius concept of acids and bases
Svante Arrhenius, 1884
Arrhenius concept:
Acid: increases concentration of hydronium ion, H3O+(aq), when dissolved in water
Base: increases concentration of hydroxide ion, OH-(aq), when dissolved in water
Hydronium and hydroxide ions are in equilibrium with water molecules, and addition of acids or bases alters this equilibrium

4. Arrhnenius concept
Strong acid: completely ionizes in water to give a hydronium ion and an anion
HClO4(aq) + H2O(l)  H3O+(aq) + ClO4-(aq)
Other strong acids: H2SO4, HI, HBr, HCl, and HNO3
Strong base: completely ionizes in water to give a hydroxide ion and a cation
NaOH(aq)  Na+ + OH-
Group IA and IIA hydroxides are strong bases (except beryllium hydroxide)
Singles out OH- as source of base character, even though other ions can give the effects of bases

5. Bronsted-Lowry concept
1923, Johannes Bronsted and Thomas Lowry
Bronsted-Lowry concept involves proton-transfer reactions
Acid: species which donates a proton
Base: specis which accepts a proton
H3O+(aq) + NH3(aq)  H2O(aq) + NH4+(aq)

6. Reversible acid-base reactions
In reversible reactions, both forward and reverse reactions involve a proton transfer
NH3(aq) + H2O(aq) ↔ NH4+(aq) + OH-(aq)
Conjugate acid-base pair: compounds that differ only by the loss or gain of a proton
NH3 (base) and NH4+ (acid) are a conjugate acid/base pair
H2O (acid) and OH- (base) are the other pair

7. Amphiprotic species
Amphiprotic species: species which can act as either an acid or a base, depending on the nature of the other reactants
Water is amphiprotic:
NH3(aq) + H2O(aq) ↔ NH4+(aq) + OH-(aq)
Base Acid Acid Base
HC2H3O2(aq) + H2O(aq) ↔ C2H3O2-(aq) + H3O+(aq)
Acid Base Base Acid

8. Lewis concept
Lewis concept:
Lewis acid: species which can form a covalent bond by accepting an electron pair from another species
Lewis base: species which can form a covalent bond by donating an electron pair to another species
BF3 + :NH3  BF3—NH3

9. Relative strengths of acids and bases
Strong acids ionize completely in water
HCl(aq) + H2O(aq)  Cl-(aq) + H3O+(aq)
Water acts as a base, accepting a proton from HCl
The forward reaction is predominant, but in the reverse reaction, H3O+ would be the acid
Of the two, HCl is the stronger acid, since it more readily donates its proton
The arrow in a reaction containing a strong acid will point to the side which contains the weaker acid
The same applies to the bases, water is a stronger base than Cl- since it more successfully attracts the proton
The arrow also points towards the weaker base

10. Relative strengths of acids and bases
But, in a weak acid like acetic acid, a small amount of its molecules are ionized
HC2H3O2(aq) + H2O(aq) = C2H3O2-(aq) + H3O+(aq)
The reverse direction predominates, and therefore H3O+ is a stronger acid than acetic acid, and acetate is a stronger base than water
Stronger acids have weaker conjugate bases
Weaker acids have stronger conjugate bases

11. Molecular structure and acid strength
Electronegativity and bond strength determine the acidity of a proton
If a proton is bonded to a more electronegative element, the bond is more polar and the proton is more easily lost
Prevalent in comparisons across rows
Long covalent bonds are weaker than short, so as the size of the atom H is bonded with increases, its acidity also increases
Prevalent in comparisons down columns
Oxoacids (H-O-Y-) increase in acidity with increasing electronegativity of Y, and with increasing number of oxygens attached to Y (both factors increase the partial negative charge of Y)

12. Self-ionization of water
Pure water is ionized to a small extent
A proton from one water molecule is transferred to another
H2O(l) + H2O(l) = H3O+(aq) + OH-(aq)
But since this only occurs to a small extent, the equilibrium constant for this process is small, and the concentration of water remains essentially unchanged
Kw = [H3O+][OH-] = 1.0 x at 25 °C
Known as the ion-product constant
The product of hydronium and hydroxide concentration for any aqueous solution is always Kw at 25 °C

13. Calculating ion concentrations
In a strong acid solution, H3O+ will come completely from the acid itself
The autoionization equilibrium is reversed due to Le Chatelier’s principle
[H3O+] will equal the acid concentration
In a solution of 0.1 M HCl, [H3O+] = 0.1 M
Similarly for strong bases, [OH-] = base conc.
Substitute into the Kw equation to find the other concentration

14. pH of a solution
In acidic solutions, [H3O+] > 1.0 x 10-7 M
In neutral solutions, [H3O+] = 1.0 x 10-7 M
In basic solutions, [H3O+] > 1.0 x 10-7 M
It is more convenient to give these values as pH
pH = -log [H3O+]
So, if [H3O+] = 1.0 x 10-3 M, pH = 3.00
Acidic solutions, pH < 7
Neutral solutions, pH = 7
Basic solutions, pH > 7

15. pH, pOH, and Kw
pOH = -log [OH-]
Kw = [H3O+][OH-] = 1.0 x 10-14
pH + pOH = 14
So, to find the pH of a basic solution, first find pOH, then subtract it from 14

16. Acid-base indicators
Indicators change color to indicate the pH of a solution
Phenolphthalein is colorless in acidic solutions, and pink in basic solutions
Protonated phenolphthalein is a colorless acid. When deprotonated, it becomes a pink-colored base.