1. Acids and Bases
Chapter 15

2. Acids and Bases
The concepts acids and bases were loosely defined as substances that change some properties of water.
One of the criteria that was often used was taste.
Substances were classified
salty-tasting, sour-tasting, sweet-tasting, bitter-tasting
Sour-tasting substances would give rise to the word 'acid', which is derived from the Greek word oxein, which mutated into the Latin verb acere, which means 'to make sour'
Vinegar is a solution of acetic acid. Citrus fruits contain citric acid.

3. Bases Acids React with certain metals to produce hydrogen gas.
React with carbonates and bicarbonates to produce carbon dioxide gas
Bases
Have a bitter taste
Feel slippery.
Many soaps contain bases.

4. Properties of Acids
Produce H+ (or H3O+) ions in water (the hydronium ion is a hydrogen ion attached to a water molecule)
Taste sour
Corrode metals
Good Electrolytes
React with bases to form a salt and water
pH is less than 7
Turns blue litmus paper to red

5. Properties of Bases Generally produce OH- ions in water
Taste bitter, chalky
Are electrolytes
Feel soapy, slippery
React with acids to form salts and water
pH greater than 7
Turns red litmus paper to blue

6. Arrhenius Definition Arrhenius
Acid - Substances in water that increase the concentration of hydrogen ions (H+).
Base - Substances in water that increase concentration of hydroxide ions (OH-).

7. Bronsted-Lowry Definition
Acid - neutral molecule, anion, or cation that donates a proton.
Base - neutral molecule, anion, or cation that accepts a proton.
HA + :B  HB :A-
HCl + H2O  H3O Cl-
Acid Base Conj Acid Conj Base

8. Conjugate Acid Base Pairs
Conjugate Base - The species remaining after an acid has transferred its proton.
Conjugate Acid - The species produced after base has accepted a proton.
HA & :A- - conjugate acid/base pair
:A- - conjugate base of acid HA
:B & HB+ - conjugate acid/base pair
HB+ - conjugate acid of base :B

9. Examples of Bronsted-Lowry Acid Base Systems
Note: Water can act as acid or base
Acid Base Conjugate Acid Conjugate Base
HCl + H2O  H3O Cl-
H2PO4- + H2O  H3O HPO42-
NH4+ + H2O  H3O NH3
Base Acid Conjugate Acid Conjugate Base : NH H2O  NH OH-
PO H2O  HPO OH-

10. G.N. Lewis Definition Lewis Acid - an electron pair acceptor
Base - an electron pair donor

11. pH and acidity HC2H3O2 + H2O  H3O+ + C2H3O2- NH3 + HCl  NH4+ + Cl-
HSO4- + HPO4-2  H2SO4 + PO4-3
HSO4- + HPO4-2  SO4-2 + H2PO4-1

12. pH and acidity
The pH values of several common substances are shown at the right.
Many common foods are weak acids
Some medicines and many household cleaners are bases.

13. The pH Scale
pH [H3O+ ] [OH- ] pOH

14. pH and acidity
Acidity or Acid Strength depends on Hydronium Ion Concentration [H3O+]
The pH system is a logarithmic representation of the Hydrogen Ion concentration (or OH-) as a means of avoiding using large numbers and powers.
pH = - log [H3O+]
pOH = - log [OH-]
In pure water [H3O+] = 1 x 10-7 M
 pH = - log(1 x 10-7) = 7
pH range of solutions:
pH < 7 (Acidic) [H3O+] > 1 x 10-7 M
pH > 7 (Basic) [H3O+] < 1 x 10-7 M

15. pH and acidity [H3O+] = 10-pH pOH = - log [OH-] [OH-] = 10-pOH
Important Equations
pH = - log [H3O+]
[H3O+] = 10-pH
pOH = - log [OH-]
[OH-] = 10-pOH
pH + pOH = 14
Kw = [H3O+] [OH-] = 1.0 x10-14
In pure water: [H3O+] = [OH-] = 1.0 x10-7

16. Calculating the pH pH = - log [H3O+]
Example 1: If [H3O+] = 1 X pH = - log 1 X 10-10
pH = - (- 10)
pH = 10
Example 2: If [H3O+] = 1.8 X 10-5 pH = - log 1.8 X 10-5
pH = - (- 4.74)
pH = 4.74

17. pH and acidity
If a substance has a pH of 3.5, determine: [H3O+], [OH-], pOH and whether it is acidic, basic, or neutral.
If a substance has an [OH-] of 8.7 x 10-5, determine: pH, pOH, [H3O+] and whether it is acidic, basic, or neutral.

18. Indicators

19. Acid Strength
Strong Acid - Transfers all of its protons to water; Completely ionized or dissociated; Strong electrolyte; The conjugate base is very weak
Weak Acid - Transfers only a fraction of its protons to water;
- Only partly ionizes or dissociates; Weak electrolyte; The conjugate base is stronger
As acid strength decreases, base strength increases.
The stronger the acid, the weaker its conjugate base
The weaker the acid, the stronger its conjugate base

20. Base Strength
Strong Base - all molecules accept a proton; completely ionizes or dissociates; strong electrolyte; conjugate acid is very weak
Weak Base - fraction of molecules accept proton; partly ionizes or dissociates weak electrolyte; the conjugate acid is stronger.
As base strength decreases, acid strength increases.
The stronger the base, the weaker its conjugate acid.
The weaker the base the stronger its conjugate acid.

21. Common Strong Acids Strong Acids Hydrochloric Acid, HCl
Hydrobromic Acid, HBr
Hydroioidic Acid, HI
Nitric Acid, HNO3
Sulfuric Acid, H2SO4
Chloric Acid, HClO3
Perchloric Acid, HClO4

22. Dissociation of Strong Acids
Strong acids completely dissociation (broke down) to the ions that make them up.
HCl  H+1 + Cl-1
HNO3  H+1 + NO3-1
If you know the concentration or molarity of a strong acid, you also know the amount of H+ ions and can find the pH (-log [H+])
0.1 M HCl = 0.1 M H+, pH = -log 0.1 = 1.0

23. Common Strong Bases Strong Bases
Any Group 1 Hydroxide and any Group 2 Hydroxide below Mg.
Sodium Hydroxide, NaOH
Potassium Hydroxide, KOH
*Barium Hydroxide, Ba(OH)2
*Calcium Hydroxide, Ca(OH)2
*While strong bases they are not very soluble

24. Dissociation of Strong Bases
Like strong acids, strong bases also completely dissociate to the ions that make them up.
KOH  K+1 + OH-1
Ca(OH)2  Ca OH-1
Just like with strong acids, if you know the molarity of a strong base, you can determine the pOH and therefore the pH.
0.01 M NaOH = 0.01 M OH-,
pOH = -log 0.01 = 2.0 and pH = 12.0

25. pH of Strong Acids and Strong Bases
Determine the pH of 0.25 M KOH.
Determine the pH of M H2SO4.
Determine the pH of M Ca(OH)2.

26. Dissociation of Weak Acids
Weak acids only dissociate to a very small degree, most break down to ions less than 5%. So an equilibrium is established where there is a small amount of product (H+ and conjugate base) but most of the acid does not break down.
These acids are weak electrolytes versus strong acids are strong electrolytes.

27. Dissociation of Weak Acids
To determine the pH of a weak acid, an “ICE” chart must be used.
I = Initial concentrations
C = Change
E = Equilibrium concentrations
Unless otherwise noted, the initial concentrations of the products are zero.

28. Dissociation Constants
For a generic weak acid dissociation,
the equilibrium expression would be
This equilibrium constant is called the acid-dissociation constant, Ka.
HA(aq) + H2O(l) 
A−(aq) + H3O+(aq)
[H3O+] [A−]
[HA]
Ka =

29. Dissociation of Weak Acids
Equilibrium expressions are written as products/reactants, and liquids and solids are left out of the expressions.
Ka are used to tell the relative strength of the acid.
You will be given the Ka value for each weak acid.

30. Dissociation of Weak Acids
Determine the pH of M HC2H3O2.
Ka = 1.8 X 10-5.

31. Dissociation of Weak Acids
Determine the pH of 1.0 M hypoiodous acid, HIO. Ka of HIO = 2.3 x

32. Determination of Ka for a Weak Acid
If you are given the initial concentration of the weak acid, and the equilibrium concentration of the conjugate base or H+, you can determine the Ka of the acid.
Remember the initial concentrations of both products are zero.
If you know the equilibrium concentration of the conjugate base or H+, you also know the value of “x” in the “C” step of the “ICE” chart.

33. Determination of Ka for a Weak Acid
Determine the Ka of a weak acid when the initial concentration of HA as M and the equilibrium concentration of H+ = M.
Equation: HA + H2O  A-1 + H3O+1

34. Titration
Titration – technique for determining the unknown concentration of an acid or base.
Titration involves delivery (from a buret) of a measured volume of a solution of known concentration (the titrant, usually a base) into a volume of solution of unknown concentration (the analyte, usually an acid).

35. Titration
Equivalence or stoichiometric point – point in a titration where enough titrant has been added to react exactly with the analyte.
Equivalence pt: moles of acid = moles of base

36. Titration

37. Titration
The equivalence point is often marked by an indicator (commonly phenolphthalein), which is added at the beginning of the titration. It changes color at (or just after) the equivalence point.
The point where the indicator actually changes color is called the endpoint of the titration.

38. Titration

39. Titration Calculations
2.00 mL of an unknown concentration of HCl was titrated with 0.25 M NaOH. If mL of NaOH was needed to reach the stoichiometric point, what is the concentration of the HCl?
Equation: HCl + NaOH  NaCl + H2O

40. Titration Calculations
2 drops of phenolphthalein is added to mL of an 0.75 M of HCl. If mL of NaOH is needed to completely react with all of the HCl, what is the concentration of the NaOH?