1. Acids and Bases

2. Acid Definition: aqueous solution of Hydrogen containing compounds
Name the following:
HCl H2SO4 H2SO3 H2S
Remember:
H with an element: hydro_______ic acid
H with polyatomic ion ending with –ate: _______ic acid
H with polyatomic ion ending with –ite: _______ous acid

3. Acid Definition: aqueous solution of Hydrogen containing compounds
Name the following:
HCl hydrochloric acid
H2SO4 sulfuric acid
H2SO3 sulfurous acid
H2S hydrosulfuric acid

4. Base Many contain the polyatomic ion OH- (hydroxide)
Name the following bases:
NaOH Ca(OH)2 Cu(OH)2 NH4OH
Follow ionic naming rules: name metal, name nonmetal, do NOT use prefixes! If the metal is a transition metal, use roman numeral for its charge

5. Base
Definition (for now): Contains the polyatomic ion OH-1 (hydroxide) Name the following bases: NaOH sodium hydroxide Ca(OH)2 calcium hydroxide Cu(OH)2 copper (II) hydroxide NH4OH ammonium hydroxide

6. Properties of Acids Ionize in water  H3O+
Corrosive to metals and skin. (React with most metals to form hydrogen gas.)
Taste sour (like lemons)
Frequently feel "sticky"
pH less than 7.
Neutralizes bases producing salt and water.
Electrolytes.

7. Properties of Bases Ionize in water  OH- Feel "slippery".
Taste bitter (like soap)
Electrolytes
pH greater than 7.
Neutralizes acids producing a salt and water

8. Indicators
Indicators are added to chemicals to determine the pH. They change colors in different pH’s.
You need to memorize the indicators and colors!

9. Acid Base Indicators Type Acid Neutral Base Red Litmus Red Blue
Blue Litmus
Phenolphthalein
Clear
Hot pink
Bromothymol Blue
Yellow
Green

10. Uses of Common Acids H2SO4 : Car batteries
HNO3 : Used to make rubber, plastics, and pharmaceuticals
H3PO4 : Flavoring agent in beverages (soda)
HCl : Stomach acid
CH3COOH: Vinegar

11. Uses of Common Bases Mg(OH)2 : antacids, milk of magnesia
NaOH : Cleaners (e.g. drain cleaners)
Ca(OH)2 : slaked lime (used in mortars, plasters, and cements)

12. Acids and Bases are electrolytes!
Why are acids and bases electrolytes?
Because they ionize in water to form ions, therefore allowing them to conduct electricity when dissolved!

13. Ionization equation for HBr:
Ionization equation for NH3:

14. Ionization equation for HBr: HBr + H2O  H3O+ + Br -
(H3O+ is called the hydronium ion and is many times used interchangeably with H+)
Ionization equation for NH3:
NH3 + H2O  NH4+ + OH -

15. Strength vs. Concentration
Concentration: amount of acid or base ÷ amount of water or solution (MOLARITY: mol/L)
Strength: how well that particular acid or base ionizes (partially or completely into H+ or OH-)
Strong acid/base: an acid/base that ionizes almost completely
Weak acid/base: an acid/base that only slightly ionizes

16. List of Strong and Weak acids
HCl
H2SO4 HF
H2SO3
HBr
HClO3
H2S
H3PO4 HI
HClO4
HNO2
HNO3
Organic acids: end in –COOH, weak acid
Ex: vinegar  CH3COOH

17. List of Strong and Weak Bases
Strong bases completely dissociates and are strong electrolytes
Weak bases only dissociate partially and are weak electrolytes
Strong
Weak
Group I and II metals with hydroxides
NH3 and any non-group I or II hydroxide

18. KOH – Strong Base KOH(s)K+(aq) + OH-(aq)
Cu(OH)2 – Weak Base Cu(OH)2(aq)↔ Cu2+(aq) + 2OH-(aq)

19. Strong vs. Weak
Strong acids or bases with completely ionize when dissolved in water. The ionization equation will have  arrow.
Weak acids or bases only partially dissociate and the dissociation equations for weak bases are always reversible .

20. Acid and Base Theories Theory Acid Definition Base Definition
Arrhenius
Releases H+ into solution
Releases OH-into solution
Bronsted-Lowery
Proton (H+) donor
Proton (H+) acceptor

21. What happens when you mix an Arrhenius acid with an Arrhenius base?
Neutralization, producing a salt and water General neutralization reaction: Salt: Ionic compound from the cation of a base and the anion of an acid

22. Complete and net ionic equations!
LiOH(aq) + HBr(aq)  LiBr(aq) + H2O(l)

23. Complete and net ionic equations! ANSWER!
LiOH(aq) + HBr(aq)  LiBr(aq) + H2O(l)
Complete
Li+(aq) + OH –(aq) + H+(aq) + Br –(aq)  Li+(aq) + Br –(aq) + H2O(l)
Net
OH –(aq) + H+(aq)  H2O(l)

24. Name the salt produced and write the balanced reaction:
Sodium hydroxide and hydrochloric acid
Calcium hydroxide and sulfuric acid
Potassium hydroxide and nitric acid

25. Name the salt produced and write the balanced reaction: Answers!
Sodium hydroxide and hydrochloric acid NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l) salt: sodium chloride, NaCl
Calcium hydroxide and sulfuric acid Ca(OH)2(aq) + H2SO4(aq)  CaSO4(aq) + 2H2O(l) salt: calcium sulfate, CaSO4
Potassium hydroxide and nitric acid KOH(aq) + HNO3(aq)  KNO3(aq) + H2O(l) salt: potassium nitrate, KNO3

26. Bronsted-Lowery Acid and Base
HCl + NH3  NH4+ + Cl-
H2O(l) + NH3(aq) ↔ NH4+(aq) + OH-(aq)

27. Bronsted-Lowery Acid and Base
HCl + NH3  NH4+ + Cl- acid base
H2O(l) + NH3(aq) ↔ NH4+(aq) + OH-(aq) acid base

28. Definitions
Monoprotic acid: only has one ionizable hydrogen ex: HCl, HBr, HC2H3O2, HNO3
Polyprotic acid: more than one ionizable hydrogen ex: H2SO4, H3PO4
Amphoteric: substance that can act as an acid or as a base ex: H2O

29. Water as an acid or base
Base H2SO4(aq) + H2O(l)  H3O+(aq) + HSO4-(aq) Acid NH3(aq) + H2O(l) ↔ NH4+(aq) + OH-(aq)

30. B-L Conjugate Acids and Bases
Conjugate acid: The species that is formed when a B-L (Bronsted-Lowery) base gains a proton
Conjugate base: The species that remains after a B-L acid has given up a proton

31. Identify the Conjugates
NH3 + H3O+  NH4+ + H2O
CH3OH + NH2-  CH3O- + NH3

32. Try these!
OH- + H3O+  H2O + H2O
H2O + NH2-  OH- + NH3

33. Try these! OH- + H3O+  H2O + H2O base acid conj. conj. acid base
H2O + NH2-  OH- + NH3 acid base conj conj base acid

34. Write the Conjugate Base
H3O+
H2SO4
HCO3-
HOCl
NH4+

35. Write the Conjugate acid
SO32-
PO43-
C2H3O2-
H2BO3-

36. Strength of conjugates
The stronger an acid is, the weaker its conjugate base; the stronger a base is, the weaker its conjugate acid.

37. Ionization of water Kw = [H3O+][OH -] = 1.oo x10-14
Pure water ionize slightly according to:
The product of molar concentrations of the ions is equal to a constant Kw.
Kw = [H3O+][OH -] = 1.oo x10-14
[ ] means “molar concentration” or molarity
This formula is in your reference table!

38. Relationships
In a neutral solution:
In an acidic solution:
In a basic solution:
Basically this is saying there is more H3O+ in acidic solutions, more OH- in basic solutions. The concentrations of both are the same in neutral solutions

39. pH and pOH Neutral 0 7 14 Acid Base Neutral 0 7 14 Base Acid
pH: negative log of [H3O+], measure of “acidity” pOH: negative log of [OH-], measure of “basicity”
Neutral
Acid Base
Neutral
Base Acid

40. pH and pOH
A change in [H+] by a factor of 10 causes the pH to change by 1.
Solution with a pH of 6 has 10x the [H+] as a solution with a pH of 7.
Solution with a pH of 3 has 1000x the [H+] as a solution with a pH of 6.
What is the [H+] difference between pH of 1 and pH of 4?
4-1= 3  that’s how many zero’s
Answer= pH 1 has 1000x more

41. Important formulas 1 4 2 5 3 6

42. Find the pH of the following:
[H3O+] = 1.00 x 10-3M
[H3O+] = 6.59 x 10-10M
[H3O+] = 7.01 x 10-6M

43. Find the [H3O+1]
pH = 3
pH = 6.61
pH = 2.52

44. Find the pH for the following:
pOH = 2
pOH = 1.26
pOH = 4.98

45. Find the pH for the following:
[OH-] = 1.00 x 10-11M
[OH-] = 2.64 x 10-13M
[OH-] = 3.45 x 10-8M

46. [OH-] = 1.00 x 10-6M [OH-] = 4.97 x 10-10M [OH-] = 2.93 x 10-2M
Find the [H3O+1]
[OH-] = 1.00 x 10-6M
[OH-] = 4.97 x 10-10M
[OH-] = 2.93 x 10-2M

47. Find the ph
0.054M HCl
0.178M NaOH
0.033M H2SO4

48. pH continued Find the pH of the following:
3.0 M HF (8.1% dissociation)

49. pH continued Find the pH of the following:
3.0 M HF (8.1% dissociation) 3.0M(0.081) = 0.24M  [H+] pH = -log(0.24) = 0.62

50. Buffered Solution A solution that can resist changes in pH
Made up of a weak acid and its conjugate base

51. Titrations
An acid-base titration is a neutralization reaction that is performed in the lab in order to determine an unknown concentration of acid or base.
The reaction is complete at the equivalence point.

52. Titrations
Endpoint: When the indicator changes color; when you stop the titration
Equivalence point: When the reaction is complete. All the acid and base have reacted to form salt and water. Not always a pH of 7!

53. Titration terms
Indicator: weak acid or base; whose color changes in different pH’s (because the conjugate base/acid is a different color)
Standard solution: solution of known concentration used as the titrant

54. Titration Curves
On the following slide you will see 3 different curves.
The relationships are:
Strong Acid titrated with Strong Base
Weak Acid titrated with Strong Base
Strong Acid titrated with Weak Base
Using those graphs compare the pH values of each substance as well as the equivalence points

55. More Titration Curves Strong acid and weak base curve
Strong base and weak acid curve
Strong acid and strong base curve

56. Diprotic Acid Curve

57. Try it!
1) If it takes 54mL of 0.1M NaOH to neutralize 125mL of an HCl solution, what is the concentration of the HCl? 2) What is the molar concentration of a 50.0mL solution of NaOH that is titrated to an endpoint with 15.0mL of a M solution of H2SO4?

58. Test Topics Acid/Base Properties Naming Acids and Bases Electrolytes
Strength vs. Concentration
Memorize Strong/Weak Acids and Bases!
Arrhenius and Bronsted-Lowry acids/bases
Conjugate acids and bases
Indicators (memorize colors)
Neutralization Reactions
pH, pOH, [H+], and [OH-] calculations
Titration concepts, calculations, and curves
Common uses of acids and bases
ALL vocab
Solutions Stuff