1. Acids And Bases
Chemistry Ms. Piela

2. Key Characteristics of Acids & Bases
Taste sour
Reacts with alkali metals (forms H2 gas)
Forms electrolyte solutions (conducts electricity)
pH paper color: Red
Neutralizes Bases
Bases
Tastes bitter
Slippery feel
pH paper color: Blue
Neutralizes Acids

3. The 3 Main Theories of Acids/Bases
Lewis Acids/Bases
This course will
mainly deal with
BL theory
Bronsted-Lowry Acids/Bases
Arrhenius Acids/Bases

4. Theories of Acids & Bases
Arrhenius Theory of Acids & Bases:
Properties of acids are due to the presence of H+ ions
Example:
HCl  H+ + Cl-
Properties of bases are due to the presence of OH- ions
NaOH  Na OH-

5. HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)
What is an H+?
H+ ions are bare protons
These are so reactive that they do not exist naturally, but will bond with water to form a hydronium ion, or H3O+ ion
Oftentimes H+ and H3O+ are used interchangeably
HCl  H+ + Cl-
HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)

6. Problems with the Arrhenius theory
Only deals with aqueous solutions (solutions in water)
Not all acids and bases contain H+ and OH- ions
Example: NH3 is a base
Considered the most incomplete
theory of acids and bases

7. Theories of Acids & Bases
Brønsted-Lowry Theory of Acids & Bases
Acids are substances that donate H+ ions
Acids are proton (H+) donors
Bases are substances that accept H+ ions
Bases are proton (H+) acceptors
Example:
HBr + H2O  H3O Br-
A B

8. Brønsted-Lowry Theory
The behavior of NH3 can be understood now:
NH3 (aq) + H2O (l) ↔ NH4+ (aq) + OH- (aq)
NH3 becomes NH4+, so NH3 is a proton acceptor (or a Brønsted-Lowry base)
H2O becomes OH-, so H2O is a proton donor (or a Brønsted-Lowry acid)

9. Brønsted-Lowry Theory

10. Brønsted-Lowry Theory
Conjugate Acid-Base Pairs
Definition: An acid and a base that differ only in the presence or absence of H+
Every acid has a conjugate base.
Every base has a conjugate acid.
These pairs only ever differ by exactly one hydrogen ion

11. Brønsted-Lowry Theory
Example Problems
Identify the Brønsted-Lowry acid, base, conjugate acid and conjugate base
NH3 + H2O  NH OH- B A CA CB

12. Brønsted-Lowry Theory
Example
HCl (g) + H2O (l) ↔ H3O+(aq) + Cl- (aq)
HSO HCO3- ↔ SO H2CO3 A B CA CB A B CB CA

13. Theories of Acids & Bases
Lewis Acids & Bases
Acids are electron acceptors
Bases are electron donors
Amphoteric – substances that can act as both an acid and a base
Examples: H2O, HCO3-

14. Summary Of Theories Acids release H+ Bases release OH-
Arrhenius
Acids release H+
Bases release OH-
Brønsted-Lowry
Acids – proton donor
Bases – proton acceptor
Lewis
Acids – electron acceptor
Bases – electron donor

15. The pH scale
Developed by Søren Sørensen in order to determine the acidity of ales
Used in order to simplify the concept of acids and bases for his workers
The pH scale goes from 0 to 14
The acidity/basicity of the solutions depends on the concentration of H+ (or H3O+)

16. The pH scale
Acidic
pH < 7
Neutral
pH = 7
Basic
pH > 7

17. pH scale Low pH values means a high concentration of H+ (acidic)
High pH values means a low concentration of H+ (basic)

18. H2O (l) ↔ H3O+ (aq) + OH- (aq)
Calculations of pH
The Self Ionization of Water
In pure water (pH = 7), the concentrations of the ions (H3O+ and OH-) are equal.
[H3O+]=[OH-]= 1x10-7
This is because water will spontaneously dissociate naturally:
H2O (l) ↔ H3O+ (aq) + OH- (aq)
Writing the equilibrium expression for the self-ionization of water gives:

19. The Self-ionization of Water
Plugging in the concentrations in pure water, this gives an equilibrium constant of 1x10-14
This is referred to as the ion product constant of water
The ion product constant of water has its own symbol: Kw
Unlike other equilibrium constants, the Kw will always be the same value

20. Calculations of H3O+/OH-
Example #1
What is the H3O+ concentration in a solution with [OH-] = 3.0 x 10-4 M?
Kw = [H3O+][OH-]
1 x = [H3O+][3.0 x 10-4]
______
___________________
3.0 x 10-4
3.0 x 10-4

21. Calculations of H3O+/OH-
If the hydronium-ion concentration of an aqueous solution is 1.0 x 10-3 M, what is the hydroxide ion concentration in the solution?
  Kw = [H3O+][OH-]
1 x = [1 x 10-3][OH-]
______
___________________
1.0 x 10-3
1.0 x 10-3

22. pH = -log [H3O+] or [H3O+] = 10-pH
Calculations of pH
pH can be expressed using the following equation:
pH = -log [H3O+] or [H3O+] = 10-pH
Example #1
What is the pH of a solution with M H3O+? Is this solution an acid or a base?
Acid!

23. Calculating pH of a solution
Example #2
What is the pH of a solution where the concentration of hydroxide ions is M? Is this an acid or a base?
Kw = [H3O+][OH-] pH = -log [H3O+]
Base!

24. Calculating pH of a solution
Practice #1
Practice #2

25. Calculating H3O+/OH- from pH
Example #1
What is the hydronium ion concentration in fruit juice that has a pH of 3.3?
[H3O+] = 10-pH

26. Calculating H3O+/OH- from pH
What are the concentrations of the hydronium and hydroxide ions in a sample of rain that has a pH of 5.05?
[H3O+] = 10-pH Kw = [H3O+][OH-]

27. Calculating H3O+/OH- from pH
Practice #1
Practice #2

28. Strength of Acids & Bases
When a solution is considered strong, it will completely ionize in a solution
Nitric acid is an example of strong acid:
HNO3 (l) + H2O (l) ⇋ NO3- (aq) + H3O+ (aq)
In a solution of nitric acid, no HNO3 molecules are present!
Strength is NOT equivalent to concentration!

29. Strength of Acids & Bases
Knowing the strength of an acid is important for calculating pH
If given concentration of strong acid (such as HNO3) assume it is the same as the concentration of hydronium, H3O+, ions
Given concentration of a strong base, assume it has the same concentration as the hydroxide, OH-, ions

30. Strong Acids & Bases Ionize 100%
Example
NaOH  Na OH-
1 M
1 M
1 M
OH-
Na+
Na+
Na+
OH-
OH-

31. Weak Acids & Bases Ionize X%
Example
HF  H F-
1 M
? M
? M F- HF H+ H+ HF H+ F- F-

32. Naming Bases Bases are soluble metal hydroxides Examples
Follow identical naming rules for ionic compounds
Examples
NaOH
Ba(OH)2
NH3
NH4+
Sodium hydroxide
Barium hydroxide
Ammonia
Ammonium

33. Naming Acids Binary Acids (HX)
If the acid has an anion that ends in “-ide” use the following basic format to name the acid:
“Hydro – root – ic acid”
Example
HCl
Hydrochloric acid

34. Naming Acids Example Practice HBr HI H2S Hydrobromic acid
Hydroiodic acid
Hydrosulfuric acid

35. Naming Acids Polyatomic acids (aka oxoacids, HxAyOz)
Name depends on the polyatomic used:
If polyatomic ends in “-ite”, replace with “ous acid”
If polyatomic ends in “-ate”, replace with “ic acid”
Trick: “I ate something icky”

36. Naming Acids Examples HClO4 HClO2 Sulfuric acid Perchloric acid
Chlorous acid
H2SO4