1. Acid/Base Theories Acids:
are electrolytes, sour tasting, dry to the touch and turn litmus paper red.
Bases:
are electrolytes, bitter tasting, slippery to the touch and turn litmus paper blue.

2. Arrhenius’ Theory HCl(l) H+(aq) + Cl−(aq)
Acids – are compounds that ionize in H2O to form H+ ions
Ionization – any process by which a neutral atom or molecule is converted into an ion
HCl(l)
H+(aq) Cl−(aq)

3. Bases – are compounds that dissociate in H2O to form OH- ions
Dissociation – the separation of ions that occurs when an ionic compound dissolves in H2O.
NaOH(s)
Na+(aq) OH−(aq)

4. • this accounts for the fact that both acids and bases are electrolytes
• cannot account for acidic solutions that are not in water
• does not explain acids that don’t have H’s or bases that don’t contain OHs.

5. The Brønsted-Lowry Theory
Acid – is a proton (H+) donor
Base – is a proton acceptor
• Acids and Bases have to react together
• there has to be an acid present to donate to a base that accepts it.
• now H2O is a reactant not just the solvent

6. eg. HCl(aq) + H2O(l)  H3O+(aq) + Cl−(aq)
transfer H+
eg. HCl(aq) H2O(l)  H3O+(aq) Cl−(aq)
acid base1
base2 acid2
conjugate acid/base pairs
• H3O+ is responsible for the acidic properties of H2O

7. eg. NH3(aq) + H2O(l)  NH4+(aq) + OH−(aq)
transfer H+
eg. NH3(aq) + H2O(l)  NH4+(aq) + OH−(aq)
base acid1
acid base2
conjugate acid/base pairs
• OH- is responsible for the basic properties of H2O

8. Water is amphoteric, it can act as an acid or a base.
• some others include HCO3−, HSO4−, HS−, NH3
eg. NH3(aq) + H2O(l)  NH4+(aq) + OH−(aq)
eg. NH3(aq) + H2O(l)  NH2−(aq) + H3O+(aq)

9. Strong Acids/Bases
 • strong or weak refers to the electrolyte’s conductivity.
• the stronger the acid/base, the weaker the conjugate
• Percent ionization gives a measure of the acid/base strength

10. Strong: 100% ionization/dissociation
Acids: only HClO4, HI, HBr, HCl, HNO3, H2SO4
Bases: only the highly soluble OH− compounds
 group 1, NH4+, Ba2+, Sr2+ and Tl+
Weak: <50% ionization/dissociation
Acids: include H3PO4, HF, CH3COOH and H2CO3
Bases: include NH3, ammonia-carbon compounds and slightly soluble OH−s such as Mg(OH)2, Ca(OH)2

11. Relative Acid/Base Strength
1) Acids vs Bases
 • both NaOH and HClO have an O-H bond
• why is one a base and the other an acid?
Na+…. O ─ H− H ─ O ─ Cl    
ionic
polar covalent
very polar
less polar
• NaOH - the ionic bond creates a stronger ion-dipole force than the polar covalent bond does with H2O
• HClO - the very polar HO bond has stronger hydrogen bonding interactions with H2O than the dipole – dipole forces of the O - Cl bond does.

12. 2) Binary Acids – HCl vs HI
• even though HCl is more polar than HI, I− ion is larger than Cl− ion which results in a weaker ionic bond, so HI takes less energy to ionize.
For water to ionize acids, they must be in the ionic form at the time of separation:
H – Cl is H+ Cl- and H – I is H+ I-

13. even though more energy is required to separate more H2O molecules from each other, the larger I− ion is surrounded by more H2O molecules where each ion-dipole interaction releases energy (solvation or hydration energy)
the greater net energy release occurs with the ionization of HI and thus its acid strength is greater.

14. 3) Polyprotic Acids
• have more than 1 acidic H’s, eg. H2SO4.
• the first ionization step is always more acidic.
• once the acid anion is formed (HSO4−), the other O ─ H bond is not as polar due to the charge on the ion
• as a result, the second or third H is harder to remove

15. 4) OxyAcids
• as the # of O’s bonded to the central atom increases, the degree of electron flow (polarization) away from the O ─ H bond increases.
• this makes the O ─ H bond more polar and allows for a greater transfer of the H.
• HClO4 is the only strong acid of the oxychloro group

16. 5) Competition for H
• Weak Acids and Bases are equilibrium systems
• the amount of ionization depends on which attraction for the H is stronger, the original acid/base or H2O