1. Substances that affect the pH of solutions.
What’s wrong in this picture?
Acids & Bases
Substances that affect the pH of solutions.

2. Acids & Bases typically are, or behave as, IONIC compounds.
Are slippery
Taste bitter
React with indicators
Neutralize acids
Ex.
NaOH (sodium hydroxide),
NH4OH (ammonium hydroxide)
Baking soda (NaHCO3)
Acids:
Are corrosive
Taste sour
React with indicators
Neutralize bases
Ex.
HCl (hydrochloric acid), H2SO4 (sulfuric acid)
Litmus is a vegetable dye obtained from certain lichens found principally in the Netherlands.

3. Typical with Acids
Typical with Bases
The difference between the aqueous solution processes of ionization and dissociation.

4. Nomenclature Acids Most are “hydrogen” bonded with an anion Examples:
HNO3 (nitric acid)
HC2H3O2 (acetic acid)
HCl (hydrochloric acid)
Bases
Most are metal hydroxides
Examples:
NaOH (sodium hydroxide)
KOH (potassium hydroxide)
Ba(OH)2 (barium hydroxide)
NH4OH* (ammonium hydroxide)

5. Definitions Acids Arrhenius - acids donate H+ (in soln)
Bronsted-Lowery -acids donate H+ (in soln)
Bases
Arrhenius - bases donate OH- (in soln)
Bronsted-Lowery - bases accept H+ (in soln)
Coordinate covalent bond

6. Conjugate Acid-Base Pairs
The transfer of protons illustrates the characteristics of conjugate pairs
HNO2 + H2O <==> H3O+ + NO2-
NO2- is the conjugate base of HNO2
H3O+ is the conjugate acid of H2O

7. Protocity Monoprotic Diprotic Triprotic
HCl, HNO3
Diprotic
H2CO3
Triprotic
H3PO4
Acids can be classified according to the number of hydrogen ions (protons) they can transfer per molecule during an acid-base reaction.

8. Acid-Base Strength (You can dilute an acid or a base but you can’t change its strength)
Strong
“ions” completely dissociate in water
ACIDS:
HCl, HBr, HI,
HClO4, H2SO4, HNO3
BASES:
LiOH, NaOH, KOH,
Ca(OH)2, Sr(OH)2, Ba(OH)2
Weak
“ions” partially dissociate in water
All non-strong acids & bases
These exist as equilibrium systems in solution, thus, their “weakness” exists within a range defined by their Keq values.

9. A comparison of the number of acidic species present in strong acid and weak acid solutions of the same concentration.

10. Weak A/B equilibrium
Two reactions (forward & reverse) occur at the same rate
HA <==>H+ + A-
BOH <=> B+ + OH-
Equilibrium expressions are ways to show the mathematical relationships
Keq = [Products]n
[Reactants]m
n & m are the coefficients of each substance

11. Ionization Constants for Acids & Bases
HA(aq) + H2O(l) <==> H3O+(aq) + A-(aq)
Ka =
B(aq) + H2O(l) <==> BH+(aq) + OH-(aq)
Kb =

12. Neutralization reactions - a special type of DR rxn
AX + BY --> AY + BX
HCl + KOH --> HOH + KCl
Acid + Base --> Water + Salt
To balance these rxns. Balance the H in the acid with the OH in the base :)!
For a complete reaction, stoichiometric equivalents of the acid and base must be used.

13. Neutralization equations
HCl(aq) + NaOH(aq) --> H2O(l) + NaCl(aq)
H2SO4 + Ba(OH)2 -->
H3PO4 + KOH -->
HNO Al(OH)3 -->
The acid-base reaction between sulfuric acid and barium hydroxide produces the insoluble salt barium sulfate.

14. Calculations
A sample of mol HCl is dissolved in water to make 1500 mL solution. Calculate the molarity of the HCl solution and the [H3O+].

15. Self-Ionization of Water (pH is a derivative of this concept)
Water molecules can break apart when they collide
H2O(l) <==> H+(aq) + OH-(aq)
Kw =
Kw = 1.0 x M2
Adding an acid or a base changes the relative amounts of [H+] and [OH-] but not the value of Kw.

16. Ionic Concentration If [H+] = [OH-] the solution is neutral
If [H+] > [OH-] the solution is acidic
If [H+] < [OH-] the solution is basic
[H+] x [OH-] = 1.0 x 10-14M2
The relationship between H3O+ and OH- in aqueous solution is an inverse proportion.

17. Calculations
If the [OH-] = 3.5 x 10-3 M, what is [H+]?

18. pH: a logarithmic scale of a solution’s hydrogen (hydronium) ion concentration (molarity)
This is a way to express the relative acidity/basicity of a solution.
pH = -log[H+]
Therefore, each difference in pH of 1.0 is equivalent to a concentration change by a factor of 10
High [H+] causes low pH
Low [H+] causes high pH
Therefore, strong acids have lower pH!
pOH = -log[OH-]

19. pH scale 0 - 14 is the usual range pH + pOH = 14 pH < 7 = acid
pH > 7 = base
pH = 7 = neutral
pH + pOH = 14

20. Calculations
If the [H+] = 3.35 x 10-5 M, what is the pH of the solution?
On your calculator:
- log (3.35 x 10-5) =

21. Calculations
If the [OH-] = 2.8 x 10-4M, what is the pH of the solution?

22. pH --> [H+] calculations
What is the [H+] for a solution with a pH = 3.92?
pH = -log[H+]
3.92 = -log[H+]
-3.92 =log[H+]
= [H+]
[H+] = 1.20 x 10-4M

23. Practice: determine the [H+] for the solutions with the following values.
[H+] = 2.82 x 10-8 M
[H+] = 3.98 x M
[H+] = 7.86 x 10-3 M
[H+] = 3.16 x M
pH = 7.55
pH = 10.4
pH = 2.12
pOH = 4.5

24. Salt Hydrolysis
Some aqueous salt solutions have the ability to split (hydrolyze) water and form compounds which result in larger [H+] or [OH-] in the solution.
Example: Aluminum chloride
AlCl3(aq) --> Al+3(aq) + 3Cl-(aq)
Cation of WB Anion of WA
Aluminum ion will react with OH- in solution:
Remember: H2O <==> H+ + OH-
Al+3(aq) + H2O(l) <==> Al(OH)3(aq) + 3H+(aq)
Chloride ion will NOT react with H+ in solution!

25. Rules for Determining pH Strength wins!
Strong Acid + Strong Base --> Neutral sol’n
HCl + NaOH --> NaCl + H2O
Strong Acid + Weak Base --> Acidic sol’n
HCl + Al(OH)3 --> AlCl3 + H2O
Weak Acid + Strong Base --> Basic sol’n
H2S + NaOH --> Na2S + H2O
Weak Acid + Weak Base --> depends on the salt
HNO2 + NH4OH --> NH4NO2 + H2O

26. Buffers
Buffers are solutions in which the pH remains relatively constant when small amounts of acid or base are added
Two active chemical species:
A substance to react with & remove added base
A substance to react with & remove added acid.
Buffers are solutions of a weak acid and one of its conjugate base OR a weak base and one of its conjugate base.

27. Carbonic acid and Sodium bicarbonate
H2CO3 <==> H+ + HCO Ka = 1.7 x 10-3
NaHCO3 --> Na+ + HCO3-

28. Buffering Action in Human Blood
H2CO3 <==> H+ + HCO3-
High concentration High concentration
Ratio: 1 :
Add a base [OH-] and the equilibrium position shifts ; pH doesn’t change much
Add an acid [H+] and the equilibrium position shifts ; pH doesn’t change much
Reason: high [ ] of acid and anion can accommodate large shifts of EQ position.
Lots of acid is
produced in the body
daily.

29. Buffer Systems

30. Titration
At the completion of a neutralization reaction (equivalence point) the
# moles acid = # moles base
So,
MaVa = MbVb
but, keep the reaction stoichiometry in mind.
Diagram showing setup for titration procedures.

31. Chemical Titration
This process can be done for any reaction in which a stoichiometric equivalence is reached and can be identified by an indicator
At the equivalence point an indicator will change color permanently.

32. Calculations
How many mL of 0.10M NaOH solution are needed to neutralize 15 mL of 0.20M H3PO4 solution?