1. Acids and Bases

2. What do lemon juice, lime juice, sauerkraut, vinegar and sour cream have in common?

3. Acids colorless Sour Electrolytes Change the color of some indicators
litmus paper changes from blue to red
phenolphthalein is
React with metals to produce hydrogen
React with bases to produce water and a salt
colorless

4. Naming acids Binary Ternary two elements hydro-element-ic
ex: hydrochloric  HCl
hydrosulfuric  H2S
Ternary
three or more elements
oxyacids  polyatomic ions
ex: sulfuric  H2SO4
chloric  HClO3

5. Bases pink Bases have a bitter taste. Bases are slippery.
Bases are electrolytes.
Bases change the color of some indicators:
litmus paper changes from red to blue
phenolphthalein is
Bases react with acids to form water and a salt.
An example of a base is soap
Bases are called alkaline
Demo: place 10 drops of 1M HCl in 50 mL water in 600mL beaker on overhead. Add 20 drops universal. Drop in one Phillips antacid tablet. Stir slowly for 2 mins. As the pH increases the color changes from orange to blue.
pink

6. Naming Bases Name positive ion – hydroxide ex: NaOH  sodium hydroxide
NH4OH  ammonium hydroxide

7. Bases
Families I and II on the periodic table readily form bases when mixed with water.
Alkali/alkaline means basic
Family I/II + H2O  metal(OH) + H2
Na H2O  NaOH H2
Note the base quickly dissociates.

8. Naming acids and bases
Stop and practice by doing worksheet 7

9. Acid/Base/Salt Acid – hydrogen ions Base –hydroxide ions
Salt – metal-nonmetal
ionic compound

10. Bronsted-Lowry acid/base
Includes all Arrhenius and more.
Acid = hydrogen donor
It gives away hydrogen ions
Base = hydrogen acceptor
It takes hydrogen

11. Bronsted-Lowry acid/base
NH3 + H2O  NH OH-
Base Acid
Although we are mixing an acid and a base the solution
is not neutral. It is basic, since there is more –OH- on
the product side.

12. Bronsted-Lowry acid/base
HF + H2O  H3O F-
Acid Base
Although we are mixing an acid and a base the solution
is not neutral. It is acidic, since there is “extra” H3O+
on the product side.

13. Bronsted-Lowry acid/base
Notice the reverse of this equation.
HF + H2O  H3O F-
Acid Base
In the reverse equation the acid and base changes
places, this is called conjugate acid and conjugate
base pairs.

14. Bronsted-Lowry acid/base
HF + H2O  H3O F-
Acid Base conj. Acid conj. Base
Conjugate acid – formed when a base acquires a proton
H+ to form the acid.
Conjugate base – the particle that remains after the
proton has been released by the acid.

15. Bronsted-Lowry acid/base
Identify the acid, base, conj. acid, conj. base for the following reaction.
HNO NaOH  NaNO3 + H2O

16. Bronsted-Lowry acid/base
Identify the acid, base, conj. acid, conj. base for the following reaction.
NaHCO3 + HCl  NaCl + H2CO3

17. Amphoteric A substance that can be both an acid and a base.
Water has a split personality. It can be both an acid and base.

18. Summary of Acid/Base Theories
Bases
Arrhenius
releases H+ ions in water.
releases OH- ions in water.
Bronsted-Lowry
donates a proton (H+)
accepts protons. (H+)
Lewis
accepts an electron pair
donates an electron pair

19. Summary of Acid/Base Theories

20. Hydrogen Ions from Water
At times water collides with so much force that one water molecule loses a H+ to another water molecule – this is called self-ionization
H – O : + H – O :  H – O – H H – O :
l l l
H H H

21. Hydrogen Ions from Water Simplified
Water decomposes into hydrogen ion and hydroxide.
H – O :  H H – O : l H
In class you may use either interchangeably. You need to recognize both.
In reality there is so much water that H+ never
occurs it is always H3O+ (hydronium)

22. Pure water In pure water the [OH-] = [H3O+]
Neutral - any aqueous solution where
[OH-] = [H3O+].
[H3O+] = 1.0 E -7 mol/L
In class you may use either interchangeably. You need to recognize both.

23. Pure water [OH-] and [H3O+] are inversely related.
Therefore, [OH-] [H3O+] = Kw
Kw = 1.0E-14 (mol/L)2
Ion-product constant for water
In class you may use either interchangeably. You need to recognize both.

24. Acidic solutions HCl + H2O  H+ + Cl- + H+ + OH- 2H+ and 1 OH-
Remember there is so much water in an aqueous solution that the H+s are really H3O+.
Therefore, [H3O+] > [OH-]
In class you may use either interchangeably. You need to recognize both.

25. [OH-] vs. [H3O+] [H3O+] > [OH-]  acid [H3O+] < [OH-]  basic
Ex: NaOH + H2O  Na+ + OH- + H OH-
[H3O+] = [OH-]  neutral
In class you may use either interchangeably. You need to recognize both.

26. [H3O+] vs. [OH-] practice problems
Calculate [OH-] if:
[H3O+] = 1E-1 M
[H3O+] = 1E-12 M
Calculate [H3O+] if:
[OH-] = 1E-3 M
[OH-] = 1E-10 M
In class you may use either interchangeably. You need to recognize both.

27. How strong is your acid?

28. How strong is your acid?
A strong acid completely ionizes (dissociates) in aqueous solution.
HCl water H+ + Cl-

29. How strong is your acid? A weak acid slightly ionizes (dissociates)
in aqueous solution.
HC2H3O2 water H+ + C2H3O2-

30. pH and pOH

31. pH and pOH

32. pH Indicates the hydronium ion concentration.
The negative logarithm of the hydronium ion concnetration.
pH = - log [H3O+]

33. pH scale
Always use pH to determine acid/base/neutral
strong
acid
strong
base
neutral

34. pH What is the pH of pure water? U: pH K: [H3O+] = 1.0E-7M
P: pH = - log [H3O+]
S: pH = - log [1.0E-7M] =

35. pH
Calculate the pH of a solution with a hydronium ion concentration of 5.0E-6M?
U: pH
K: [H3O+] = 5.0E-6M
P: pH = - log [H3O+]
S: pH = - log [5.0E-6M]
5.30

36. pOH Indicates the hydroxide ion concentration.
The negative logarithm of the hydroxide ion concnetration.
pOH = - log [OH-]
never use pOH to determine acid/base/neutral!! (ZERO CREDIT!!)

37. Equations Kw = 1.0E-14 = [H3O+][OH-] pH = -log[H3O+] pOH = -log[OH-]
pH + pOH = 14

38. pH and pOH problems
Determine the pH and pOH of a solution which contains mole of H3O+ per liter. Is this an acid/base or neutral solution?
P: pH = -log [H3O+] ; pH + pOH = 14
S: pH = - log [0.0035M]
= 2.46
pOH = 14 – 2.46 = 11.54
acid

39. pH and pOH Problems
What is the pOH of a solution containing 0.042M KOH? Is this a/b/n?
since this is a strong base  100% ionizes
Therefore: KOH  K OH-
before 4.2E
after E-2M 4.2E-2M

40. pH and pOH Problems (cont’d)
What is the pOH of a solution containing 0.042M KOH? Is this a/b/n?
u: pOH
k: [OH-] = 4.2E-2M
p: pOH= -log [OH-]
s: pOH = -log [4.2E-2M]
pOH = 1.38
pH = 14 – 1.38 =  base

41. pH and pOH Problems
What is the pH of a solution containing 0.035M H2SO4? Is this a/b/n?
since this is a strong acid  100% ionizes
Therefore: H2SO4  2H SO4-2
before 3.5E
after E-2M 3.5E-2M

42. pH and pOH Problems (cont’d)
What is the pH of a solution containing 0.035M H2SO4? Is this a/b/n?
u: pH
k: [H+] = 7.0E-2M
p: pH= -log [H+]
s: pH = -log [7.0E-2M]
pH = 1.15  acid

43. pH and pOH Problems
Calculate the [H3O+] of a solution that has a pH of 3.70.
u: [H3O+]
K: pH = 3.70
p: pH = -log [H3O+]  pH
s: [H3O+] =
= 2.0E-4M

44. Titration

45. Titration
an experimental procedure in which the unknown concentration of a known volume of solution is determined by measuring the volume of a solution of known concentration required to react completely with it.

46. Titration
an analytical method in which a standard solution is used to determine the concentration of another solution.
Standard solution – one of known concentration.
End point – the point at which neutralization is achieved.

47. Steps for Titration
A measured amount of an acid/base of unknown concentration is added to a flask.
An appropriate indicator (phenolphthalein, BB, or methyl orange) is added to the solution.
A measured amount of base/acid (known concentration) is mixed into the first solution.

48. Titration Experiment Preview
A ring stand, a clamp, a burette, a 250mL flask, and a white sheet of paper are required materials
The materials should be assembled as the picture shows.

49. Titration Experiment Preview
Condition the burette.
Add a small amount of NaOH to the burette (about 5 mL).
Place the burette almost horizontal and turn.
Empty contents through stopcock into sink.

50. Titration Experiment Preview
Be sure the stopcock of the burette is closed!
Using the funnel/stirring rod at the top, fill the burette with the NaOH solution.
In the flask place the measured HCl solution, and 5 drops of the indicator.

51. Titration Experiment Preview
Begin the titration by opening the stopcock to make the NaOH solution flow dropwise into the flask.
Be sure to swirl the flask constantly so the solutions mix thoroughly.

52. Titration Experiment Preview
Continue to add NaOH solution to the flask until a color change begins (colorless to pink)
At this point, slow the flow even more. (It is over titrated when the solution turns fuchsia.)

53. Titration curve
Graphical depiction of pH of solution as it is neutralized by a given volume of titrate added.

54. Titration curve: strong vs. strong

55. Titration curve: strong acid vs. weak base

56. Titration curve: weak acid vs. strong base

57. Titration curve: weak vs. weak

58. Titration curve summary

59. Titration problem
MaVa = MbVb

60. Titration problem sample 1
In a laboratory experiment, 20.0 mL of NH3(aq) solution is titrated to the methyl orange endpoint using 15.65mL of a 0.200M HCl solution. What is the concentration of the aqueous solution?

61. Titration problem sample 1
U:Mb K: vHCl=15.65mL = L Ma = 0.200M HCl Vb = 20.0mL = 0.020L P: MaVa = MbVb S: (0.200M)( L) = (Mb)(0.020L) 0.157M = MNH3