1. ACIDS and BASES

2. Properties of ACIDS Sour to taste
React with some metals to form Hydrogen gas
Turn Litmus RED
Phenolphthalein stays colorless
Electrolytes (conduct)
Form H+ (H3O+ when attached to water molecules)

3. Properties of BASES Bitter to taste Slippery to touch Turn Litmus BLUE
Phenolphthalein turns MAGENTA
Electrolytes
Many form OH- in water

4. Definitions of Acids and Bases
Arrhenius (traditional)
Acids: produce H3O+ in water
Bases: produce OH- in water
Examples
HCl (g) → H+(aq) + Cl-(aq)
NaOH (s) → OH-(aq) + Na+(aq)
Most common definition

5. Definitions of Acids and Bases
Bronsted-Lowry
Acids: H+ donor (proton donor)
Bases: H+ acceptor (proton acceptor)
Examples
HCl → Cl- (donates H+)
NH3 → NH4+ (accepts H+)

6. Bronsted-Lowry cont.
When acids and bases donate or accept hydrogen ions, conjugate acids and bases are formed.
Conjugate Acid: particle formed when a Base gains a H+
Conjugate Base: particle formed when an Acid donates a H+

7. Bronsted-Lowry cont.
Conjugate acid-base pair: two substances related by the loss or gain of a single H+
Always paired with an acid and base
Examples- Label acid/base and conjugate acid/base
NH3 + H2O → NH4+ + OH-
HCl + H2O → H3O+ + Cl-
Hint: Acid/Base is always Reactant; Conjugate Acid/Base is always Product

8. Terms Amphoteric substances can behave as acids or bases (water)
Monoprotic acids donate only one H+
Ex. HCl, HNO3
Polyprotic acids donate more than one H+
H2SO4, H2CO3

9. Relative Strengths of Acids and Bases
The stronger the acid, the weaker the conjugate base.
H+ is the strongest acid that can exist in equilibrium in aqueous solution.
OH- is the strongest base that can exist in equilibrium in aqueous solution.

10. Relative Strengths of Acids and Bases

11. Conjugate Acids and Bases
Any acid or base that is stronger than H+ or OH- simply reacts stoichiometrically to produce H+ and OH-.
The conjugate base of a strong acid (e.g. Cl-) has negligible acid-base properties.
Similarly, the conjugate acid of a strong base has negligible acid-base properties.

12. Acid-Base Equilibria
In pure water, the following equilibrium is established:
This is referred to as the autoionization of water.

13. Types of Acids and Bases
Weak acids and bases
Weak electrolytes
Partially ionize in water (much less than 100%)
Establish equilibria
Weak Acids
H3PO4, HC2H3O2, H2CO3
Weak Bases
NH3, low [OH-]

14. pH Pouvoir hydrogene: “Hydrogen Power” Uses [H3O+] or [H+]
Measure of the acidity of a solution
Uses [H3O+] or [H+]
Concentrations usually expressed as powers of 10
pH = -log [H+]
pH scale 0-7 acid, 7 neutral, 7-14 base
Can be lower than 0 or higher than 14

15. pH Scale

16. pOH “Hydroxide Power” Uses [OH-] pOH = -log [OH-]
Measures alkalinity of a solution
Uses [OH-]
pOH = -log [OH-]
pOH scale 0-7 base, 7 neutral, 7-14 acid
Can be lower than 0 or higher than 14
pH + pOH = 14
From the autoionization of water

17. Calculating Ka from pH Weak acids are simply equilibrium calculations.
The pH gives the equilibrium concentration of H+.
Using Ka, the concentration of H+ (and hence the pH) can be calculated.
Write the balanced chemical equation clearly showing the equilibrium.
Write the equilibrium expression. Find the value for Ka.
Write down the initial and equilibrium concentrations for everything except pure water. We usually assume that the change in concentration of H+ is x.
Substitute into the equilibrium constant expression and solve. Remember to turn x into pH if necessary

18. Percent ionization is another method to assess acid strength.
For the reaction

19. Percent ionization relates the equilibrium H+ concentration, [H+]eq, to the initial HA concentration, [HA]0.
The higher percent ionization, the stronger the acid.
Percent ionization of a weak acid decreases as the molarity of the solution increases.

21. Polyprotic acids Polyprotic acids have more than one ionizable proton.
The protons are removed in steps not all at once:
It is always easier to remove the first proton in a polyprotic acid than the second.
Therefore, Ka1 > Ka2 > Ka3 etc.

22. Relationship between Ka and Kb
For a conjugate acid-base pair
Therefore, the larger the Ka, the smaller the Kb. That is, the stronger the acid, the weaker the conjugate base.
Taking negative logarithms:

23. Strong Acids and Bases Strong electrolytes
Ionize (separate) 100% in water
Strong Acids: HCl, HNO3, H2SO4, HClO3, HClO4, HI, HBr
Usually only source of H+, so pH can be calculated from the molarity of the acid (unless < 10-6)
Strong Bases: NaOH, KOH, Ca(OH)2
Most Group 1 and 2 metal hydroxides are strong bases
Ionic metal oxides, hydrides, and nitrides

24. Salt Solutions Most salts are strong electrolytes (ionize in solution)
Acid-Base properties result from their ions.
Many ions react with water to form H+ and OH- (hydrolysis)
Anions from weak acids are basic.
Anions from strong acids are neutral.
All cations (except alkali/alkaline earth metals) are weak acids

25. Salt Solutions SA + SB = Neutral Salt SA + WB = Acidic Salt
WA + SB = Basic Salt
WA + WB = Acidic or Basic Salt
Based on relative strength of Ka and Kb
Ka > Kb = acidic
Ka < Kb = basic

26. Acid-Base Behavior and Chemical Structure
For compound H-X,
If H is partially positive, then it is an acid
If H is partially negative, then it is a base
Bond Strength and Polarity affects acid/base strength
Bond strength used to determine strength in a group; Bond polarity used to determine strength in a period
Acid strength tends to increase down a group; Base strength tends to decrease down a group
Acid strength tends to increase L to R; Base strength tends to decrease L to R

27. Acid-Base Behavior and Chemical Structure

28. Oxyacids
Acids that contain OH groups bound to the central atom ( Y – O – H)
Strength depends on Y and the atoms attached to Y
Increasing electronegativity of Y = increasing acidity
Increasing the number of O atoms attached to Y increases polarity, which increases strength
Ex. HClO < HClO2 < HClO3 < HClO4

29. Carboxylic Acids Contain –COOH
Additional oxygen atom on the carboxyl group increase the polarity of the O-H bond and stabilizes the conjugate base

30. Lewis Acids and Bases Lewis Acid – electron pair acceptor
Lewis Base – electron pair donor
Do not need to contain protons – most general definition of acids/bases
Many Lewis acids have an incomplete octet (BF3)
Transition metal ions are usually Lewis acids
Lewis acids must have an empty orbital
Compounds with multiple bonds can be Lewis acids

31. Amino Acids Amphoteric
Contains carboxyl group (acid) and ammine group (basic)
Proton of the carboxyl group is transferred to the basic nitrogen of the ammine
Results in a zwitterion or dipolar ion