1. Acids and Bases

2. Properties of Acids and Bases
Acids – H+ ion, sour taste (lemons, limes), react with metals very readily.
Bases – OH- ion, bitter taste (baking soda), feel slippery to the touch (soap).

3. Arrhenius Acid
An Arrhenius acid is a substance that contains hydrogen and ionizes in water to produce H+ ions.
HCl(g)  H+(aq) + Cl- (aq)

4. Arrhenius Base
An Arrhenius base is anything that will dissociate in water to produce OH- ions.
NaOH(s)  Na+ (aq) + OH- (aq)

5. BrØnsted-Lowry model An acid is a H+ (proton) donor.
HCl(g)  H+(aq) + Cl- (aq)
Notice that HCl donated an H+ ion.
A base is a H+ (proton) acceptor.
NaOH(s)  Na+ (aq) + OH- (aq)
The OH- ion will accept a H+ ion.

6. Naming Acids
There are two main types of acids Binary acids and Oxyacids.
Binary acids contain only two elements. One of them is hydrogen and the other is a anion.
Ex. HCl, HBr, HI, HF

7. Binary Acids
When naming a binary acid you name the hydrogen by putting hydro-
You name the second atom as normal but drop –ide and add –ic.
Lastly add acid

8. Oxyacids To name an oxyacid you must identify the polyatomic ion.
Put the name of the polyatomic ion, drop –ide and add –ic.
Lastly add the word acid.
DO NOT name the hydrogen in oxyacids.

9. Conjugate Acid Base Pairs
A conjugate acid is the product when a base accepts the H+ ion from an acid. Conversely a conjugate base is what is left behind once an acid donates to a base.

10. Conjugate Acid Base Pairs

11. Conjugate Acid Base Pairs

12. Conjugate Acid Base Pairs

13. Conjugate Acid Base Pairs

14. Strong and Weak Acids and Bases
Ionization
Strong and Weak Acids and Bases

15. Strong acids
We know that if we dissolve an ionic substance into water it will ionize or dissociate.
HCl + H2O ↔ H3O+ + Cl-

16. Strong acids are called such because they ionize completely.
That means that there will be no HCl left only H3O+ ions and Cl– ions.

17. Acid Dissociation Constant
The acid dissociation constant (Ka) is a measure of the strength of an acid.
The greater the Ka the stronger the acid. For strong acids the Ka is 1.

18. Weak Acids
Weak acids do not ionize completely. Only a small portion of the acid will ionize, leaving some reactant “un-ionized”.

19. HA + H20 ↔ H3O+ + A-
Ka = [H3O+] [A-]
[HA]

20. Calculating Ka
A 0.1 M solution of HF a weak acid has a H3O+ concentration at equilibrium of M. Calculate Ka for the acid.
HF + H2O ↔ H3O+ + F-

21. Ka = [H3O+] [F-]
[HF]
Enter in the concentrations for each species and the solve.

22. pH and pOH

25. pH and pOH scales

26. pH scale
The [H3O+] concentrations are often very small numbers and can be cumbersome. Chemists adopted an easier way to express concentrations using a pH scale.
pH is the negative log of the [H3O+]

27. pH = -log[H+] A pH below 7 is acidic, and a pH above 7 is basic.
Conversely there is the pOH scale. This is a measure of the concentration of the [OH-].
pH + pOH = 14
14-pH = pOH

28. pOH = -log[OH-]
If you are asked to find the pH of a solution that has an [OH] concentration then you will find the pOH then subtract from 14 to find the pH.
Find the pH of a solution with a [OH-] concentration of 8.2X10-6M.

29. -log[8.2X10-6] = pOH = 5.08
We want to find the pH not the pOH.
14 – pOH = pH
14 – 5.08 = pH = 8.92

30. Calculate concentration from pH
Now that we are pH calculating machines we can also calculate concentration from pH.
If we were told that a solution of weak acid had a pH of 5.2 we could then find the [H3O+].
Antilog (-5.2) = [H3O+] = ??

31. [H+] Vs. [OH-] If water is neutral what is the [H+]?
What is the [OH-] of water?
2H2O ↔ H3O+ + OH-
This is what water does.

32. [H+] Vs. [OH-] Keq = [H3O+][OH-] [H2O] Kw = [H3O+][OH-]
Kw is the ion product constant for water.

33. [H+] Vs. [OH-] Kw = [H3O+][OH-] Kw = 1.0X10-14 = [H+][OH-]
For pure water the [H+] and [OH-] are equal so they each equal 1.0X This is not true for acidic or basic solutions.

34. [H3O+] Vs. [OH-] PROBLEM TIME!
The [H+] ion in a solution is 1.0X10-5M. What is the [OH-] ion concentration in the solution? Is the solution acidic, basic, or neutral?

35. Kw = 1.0X10-14 = [H+][OH-] Given- [H3O+] = 1.0X10-5 Kw = 1.0X10-14
Unknown = [OH-]
Kw = [H3O+][OH-]
How do we manipulate our equation to solve for [OH-]

36. [OH-] = Kw
[H3O+]
Now plug in your numbers to solve
[1.0X10-14] = [OH-] = 1.0X10-9
[1.0X10-5]
Which concentration is greater the [H3O+] or the [OH-]?

37. Acidic or Basic [H+] = 1.0X10-5 [OH-] = 1.0X10-9
1.0X10-5 > 1.0X10-9
[H+] > [OH-], since [H+] is greater the solution is acidic.

38. Acid Base Neutralization
When an acid and a base are mixed together the products are water and a salt.
HCl + NaOH  H2O + NaCl
HNO3 + KOH  H2O + KNO3
H2SO4 + Mg(OH)2  2H2O + MgSO4

39. Neutralization
For an acid to neutralize a base you must have equal moles of each present.
We use the following equation.
MacidVacid = MbaseVbase
Where M is molarity and V is volume.

40. If 25 mL of 0.3 M HCl is added to 57 mL of an unknown concentration of NaOH and completely neutralizes it what is the concentration of NaOH?
Macid = 0.3 M
Vacid = 25 mL
Mbase = ?
Vbase = 57 mL

41. MacidVacid = MbaseVbase
Rearrange to find Mbase
MacidVacid = Mbase
Vbase
25 ml X 0.3 M = 0.13 M
57 mL