1. Acids and Bases
Chapter 16

2. A special solution Acids and bases are ALWAYS in a water solution.
Your body has water in it so they are always dangerous to living things.
Bases are just as dangerous as acids.
In low concentrations they are not that dangerous and found all over your house.

3. D B Acids and Bases Although they can be dangerous, acids
and bases do not react with or “eat”
everything.
Neither has an effect on glass for example. D B
Aci Re
less than 7
sour
metals and bases
ase
lue
more than 7
bitter and feel slippery
oils and acids
turn litmus paper
have a pH
taste
react with

4. Common places to find acids and bases
Vinegar- acetic acid
citrus fruits- citric acid
carbonated drinks- carbonic acid
Your stomach- hydrochloric acid
Bases
Antacid tablets (calcium hydroxide)
Windex- ammonia
Oven cleaner- sodium hydroxide
Draino – sodium hydroxide

5. Homework
Using the litmus paper provided in class check to see if two common items found in your house are acidic or basic.
Report your findings on a piece of paper, and staple the litmus paper with it.
You need to report what the items are and if they are acidic or basic.
Please exercise caution and common sense. Do NOT test anything dangerous!

6. Definitions Acid- a proton (H+) donor [force feeder]
Acids produce H3O+ (hydronium) in water
Base- a proton (H+) acceptor [thief]
Bases produce OH- (hydroxide) in water

7. Heat of solution
Normally dissolving a substance is an exothermic process.
You are normally increasing the state of entropy (measure of disorder)
Which normally means you will release heat.
There are exceptions, dissolving ammonium nitrate is an endothermic process

8. Always do what you oughta …
Always add acid to water
Dissolving the acid in water releases heat
If you have a lot of acid and a little water on top, the water typically boils quickly causing the hot acid to spray out.
A lot of water on the bottom typically doesn’t boil if the acid is added slowly enough.

10. Acid Nomenclature Review
No Oxygen
w/Oxygen
An easy way to remember which goes with which…
“In the cafeteria, you ATE something ICky”

12. Acid Nomenclature Review
HBr (aq)
H2CO3
H2SO3
 hydrobromic acid
 carbonic acid
 sulfurous acid

13. Name Them:
HI (aq)
HCl (aq)
H2SO3
HNO3
HClO2

14. Self dissociation of water.
Some water will dissociate itself
H2O +H2O  H3O+ +OH-
in “pure” water you will find
H3O+ has concentration of 1 x 10-7 M
OH- has concentration of 1 x 10-7 M
The product of the conc. of H3O+ and OH- is always 1 x [ ]-conc.
[H3O+] [OH-] = 1 x 10-14

15. pH
In any solution the H3O+ and OH- concentration is always very small.
pH- method of representing the H3O+ concentration in a solution.
pH = -log [H3O+]
So the pH of water is…
pH = - log 1 x10-7
pH = 7

16. What is a log log stands for logarithm
~we can use them to solve for an exponent.
log xy = y log x
For example log 1 x10-7 = -7
the log key on your calculator is log10 meaning it will cancel out a 10^.
To reverse a log10 raise the whole thing to the 10th power (10^), this is an antilog
The reversed pH equation is
[H3O+] = 10^(-pH)

17. pH values pH of 7 is neutral- equal [H3O+] and [OH-]
below 7 is acidic, higher [H3O+] than [OH-]
above 7 is basic or alkaline, higher [OH-] than [H3O+]

18. Acid Base Equations [H3O+] [OH-] = 1 x 10-14 pH = -log [H3O+]
[H3O+] = 10^(-pH)

19. Sig Figs and pH
The number of decimal places in the log value, pH value, is equal to the number of significant figures in the number that we took the logarithm of, concentration.
So [H3O+] = 2.45 x10-4 M 3 sig figs
pH = -log 2.45 x10-4 M = 3.611
3 decimal places

20. Reversing that
Having a pH of 4.32 (2 decimal places) gives you a hydronium concentration of…
[H3O+] = 10^(-4.32) = 4.8 x10-5 M
(2 sig figs)

21. pH problems What is the pH of a 2.4 x 10-4 M H3O+?
pH = - log 2.4 x 10-4
pH = 3.62
What is the OH- concentration?
[H3O+] [OH-] = 1 x 10-14
2.4 x10-4 [OH-] = 1 x 10-14
[OH-] = 4.2 x10-11 M

22. Backwards problem
What is the [H3O+] and [OH-] of a solution with a pH of 8.75?
[H3O+] = 10^(-pH)
[H3O+] =
[H3O+] = 1.8 x 10-9 M
1.78…x10-9 [OH-] = 1 x 10-14
[OH-] = 5.6 x 10-6 M

23. Last one
What is the pH and [H3O+] of a solution with a [OH-] conc. of x10-4 M?

24. Homework
What is the pH and [H3O+] of a solution with a [OH-] of 5.92 x10-5 M?

25. Neutralization of an acid or base.

26. Mixing acids and bases ~creates water H3O+ + OH-  2 H2O
this is called neutralizing the solution
a neutralized solution is no longer dangerous.
The point where neutralization is complete is called the equivalence point

27. Salts ~the byproduct of an acid and a base. NaOH + HCl  H2O + NaCl
(base) (acid) (water) (salt)
there are several more than just table salt.
HNO3 + NH4OH  H2O + NH4NO3
Acid Base water salt

28. Gases can be created this depends on the reactants (not all will)
sodium bicarbonate (baking soda) will pretty much always release a gas
NaHCO3 + H2SO4 H2O + NaHSO4 + CO2
Salt
Gas

29. Titration
~mixing an acid and base perfectly to make a neutral solution.
You normally need some kind of indicator for this.
Phenolphthalein- when in solution turns red if basic and is clear if acidic.
You can also use a pH probe

30. Graph of titration
equivalence point pH
Volume strong base added

31. Using math
To neutralize a solution you will need to add an equal amount of H3O+ / OH- to what was already present.
so that
mol H3O+ = mol OH-
This is used if and only if you are at the equivalence point (completely neutral solution)!

32. Problem
If 94 mL of 4.0 M NaOH neutralizes 6.0 L of an unknown strong acid, what was the H3O+concentration of the unknown?
4 M NaOH x .094 L = .376 mol NaOH
.376 mol H3O+/ 6.0 L = .063 M H3O+

33. Another problem
If 127 mL of 2.0 M NaOH neutralizes 4.1 L of an unknown acid, what is the initial concentration of the acid?
2.0M(.127 L) = .254 mol NaOH
=.254 mol OH- = .254 mol H3O+
4.1 L
=.062 M

34. Equilibrium
Chapter 17

35. Rates of reaction Different reactions happen at different speeds.
There are ways to speed up or slow down a reaction.
Changing the surface area of the reactants
Powders react more quickly than “chunks”
Changing the amount/concentration of reactants
More reactant speeds up the reaction
Changing the temperature
Warmer reactions tend to go faster
Amount of rate change depends on the reaction.

36. Catalysts and Inhibitors
Catalyst- something that increases the rate of a reaction without changing the products of the reaction.
Catalytic converter speeding the reaction of emissions of a car to less dangerous products
Inhibitor- something that slows or stops a reaction
-food preservatives

37. Catalyst example
2 O3  3 O2
Ozone will decompose into elemental oxygen, however this process is very slow.
Chlorine acts as a catalyst as shown in this two step reaction
2 O3 + 3 Cl2 6 ClO
6 ClO  3 O2 + 3 Cl2
ClO is an intermediate, something formed in the middle of the reaction that is later consumed.
Chlorine is a catalyst because it is a reactant in the first step, but a product in the last step. So it isn’t used up during the reaction.

38. Forwards and backwards
Most reactions can go forwards or backwards
Neutralization equation
H3O+ + OH-  2 H2O
Self ionization of water
2 H2O  H3O+ + OH-

39. Equilibrium
In water, both of those reactions are occurring simultaneously.
Equilibrium is when the forward and backward reactions are occurring at the same rate.
This will cause a stable amount of product and reactant to be present. No net change is occurring when it is at equilibrium. (dynamic equilibrium)
The amount of product and reactant do NOT have to be equal!

40. Representing equilibrium
It is normally represented with a double arrow
2 H2O H3O+ + OH-
This reaction comes to equilibrium when [H3O+ ] = 1 x10-7 M and [OH- ] = 1 x10-7 M (assuming the solution is neutral)
you won’t have to calculate this.

41. Le Châtelier’s Principle
~whenever stress is applied/changes are made to a system at equilibrium, a new equilibrium will be obtained to balance this stress/change.
stress is a change in temperature, pressure, or concentration of some component.
This will change the rate of reaction of either the forward or backward reaction
So you will see an increase in the concentration of the substances on one side of the equation, and a decrease on the other.
This will “shift” the equation to the right or left.

42. Examples
Endothermic reactions absorb heat, i.e. they need heat to react.
If the solution is heated prior to the reaction (stress)…
It will react more quickly
So the equation will be forced to the right (product side)
If the reaction is cooled, it will be forced to the left (reactant side)

43. Equilibrium

44. Equilibrium
Systems at equilibrium are still dynamic (changing). However, no NET CHANGE will be observed.
A system is at equilibrium when the rate of the forward reaction is equal to the rate of the reverse reaction.

45. Changing concentration
2 H2O  H3O+ + OH-
If I add more water
It will force the reaction to the right
Which means more hydronium and hydroxide will be produced
This is dilution (making the ratio of hydronium/hydroxide closer)

46. Equilibrium Add water 2 H2O  H3O+ + OH- Stress + X 0 0
Shift y y y
Final X - 2y y y
*where X is the amount of H2O added
Since the stress was added to the left, we must take from the left and give to the right to relieve the stress
*y is the amount of water that “shifts” over to make more hydronium and hydroxide. For every 2 H2O molecules, one H3O+ and OH- is produced

47. What this means… Add water 2 H2O  H3O+ + OH- Final + X - 2y + y +y
The overall amount of water increased because X is always larger than y (with any coefficient).
We increased H3O+ because +y is an increase
We increased OH- because +y is an increase
The amount H3O+ increased is equal to the amount OH- increased

48. The only equilibrium calculation
That you will have to do with numbers is:
[OH-] [H3O+ ] = Kw
Kw is the equilibrium constant for water, it equals 1 x M
We have already used the equation

49. More Le Châtelier’s If I add an acid to the equilibrium…
2 H2O  H3O OH-
Stress X
Shift y y y
Final y X- y y
  
*Where X is larger than 2y
so adding acid will decrease the [OH-], only slightly increase the[H3O+ ], and increase water.

50. 2 H2O  H3O+ + OH- If I remove hydroxide from the solution…
Stress X
Shift y y y
Final y y X +y
  
*Where X is larger than 2y
So removing hydroxide increases [H3O+], only slightly decreases [OH-], and decrease the water

51. Different equation Adding ammonia, NH3, to the equilibrium
2 NH3  3 H2 + N2
Stress X
Change y y +y
Final X-2y y +y
  
*where X is larger than 2y
Everything increases
Note that the amount H2 increases 3x as much as N2

52. With heat If I cool the following equilibrium
Heat+ Co Cl-  CoCl42-
stress -x
Shift y y y
Final y y y
  
So cooling the solution will cause more Co2+ & Cl- and less CoCl42- to form

53. Conjugate acids and bases

54. Different definitions of acids and bases
Acids are proton donors (Brønsted Lowry definition)
they generate H3O+ in water (Arrhenius definition)
Bases are proton acceptors
they generate OH- in water
which is an acid/base?
HF + H2O  H3O+ + F-
NaHCO3 + H2O  Na+ +H2O + CO2 + OH-

55. Follow the proton HF + H2O  H3O+ + F-
NaHCO3 + H2O  Na++ H2O+CO2 +OH-
Joining equilibrium to acid base…What about the reverse reaction? H+ H+ H+ H+

56. Conjugate acids and bases
When you run the reverse reaction you find the products are also acids and bases. The acids and bases that are formed are called conjugate acids or bases
H2O + HF  H3O F-
base acid conjugate acid conjugate base
NaHCO3 + H2O  Na+ +H2O + CO2 +OH-
base acid CA CB

57. Label Acid, Base, Conjugate Acid, Conjugate Base
HClO3 + H2O  ClO3- +H3O+
ClO- + H2O  HClO + OH-
HSO4- + H2O  SO42- +H3O+
LiOH + H2O  Li+ + H2O + OH-

58. Strong and Weak Acids and Bases
A strong acid or base is one that completely dissociates into water, making the most possible H3O+ or OH- that is possibly can.
HCl  H+ + Cl-
A weak acid or base is one that does not completely dissociate, and is an equilibrium reaction.
HF ⇌ H+ + F-

59. Conjugate acids and bases …
Conjugate acids and bases determine if an acid or base is strong or weak.
If the conjugate acid/base readily reacts to run the reverse reaction it is a weak acid/base.
If it does not react in the reverse reaction the acid or base is strong.

60. Strong acids Acid formula Nitric Acid HNO3 Sulfuric Acid H2SO4
Hydrochloric acid
HCl

61. Strong Bases Name Formula Sodium Hydroxide NaOH Calcium Hydroxide
these make a lightning bolt
on the periodic table!
Name
Formula
Sodium Hydroxide
NaOH
Calcium Hydroxide
Ca(OH)2
Potassium Hydroxide
KOH
Strontium Hydroxide
Sr(OH)2
Barium Hydroxide
Ba(OH)2

62. More with conjugate acids/bases
H2SO4 + H2O  H3O+ + HSO4-
Sulfuric acid is a strong acid so its conjugate base, HSO4-, will not run the reverse reaction.
HSO4 - is actually an acid in water.
HSO4 - + H2O  H3O+ + SO42-
SO42- will run the reverse reaction, so it is a weak acid

63. Other weak acids and bases
Acetic Acid (vinegar)
Citric Acid
Ascorbic Acid (vitamin C)
Boric Acid
Carbonic Acid
Weak Bases
Sodium Bicarbonate
Ammonia
Sodium Hypochlorite (bleach)

64. Danger!!!
Strong and Weak acids and bases do NOT necessarily tell you how dangerous they are.
Concentration is the most important factor for determining danger.
Ammonia is a weak base, if it is highly concentrated it can burn you.
Dilute hydrochloric acid (less than 1 M) is not particularly dangerous

65. What is water
Water is either an acid or base depending on the situation.
Anything that is either an acid or a base is called amphoteric.
Several things are amphoteric, like parts of you.

66. Donating Protons
Hydrochloric acid (HCl) can donate 1 proton, so it is called a monoprotic acid.
Sulfuric acid (H2SO4) can donate 2 protons, so it is called a diprotic acid.
Phosphoric acid (H3PO4) can donate 3 protons, so it is called a triprotic acid.

67. Indicators

68. Weak acids and bases can act as an indicator
can be forced the other way
So ammonia…
NH3 + H2O NH4++OH-
Ammonia is a gas with a distinct odor
Ammonium and hydroxide are both odorless.
If base is added to the solution you will smell ammonia, if hydroxide is removed you won’t.

69. Pet Stain Problem Urine has ammonia in it. Most cleansers are basic
NH3 + H2O ⇌ NH4++OH-
If I stress this equilibrium by adding a base…
It shifts to the left causing more ammonia to form. Animals sense of smell is better so they fine the same spot and mark it again.
That is why there are special cleansers (acidic) for pet stains

70. Indicators
Indicators are a substance that change color in the presence of (whatever they check for)
They do this because of Le Châtelier’s principle. All you need an equilibrium reaction with different colored products and reactants.
The pen used to check for counterfeit money is a starch indicator

71. How an acid base indicator works
A generic indicator will follow this reaction, HId is the reactant indicator, and Id- is its product
HId + H2O  H3O+  + Id-
The color differences are important
in an acidic solution (high H3O+) you see reactant
HId + H2O   H3O+  + Id-
in a basic solution (low H3O+) you see product
HId + H2O  H3O+  + Id-

72. Acid Base indicators
Acid base indicators change color at certain pH levels
They don’t have to change at 7 (most don’t)
Universal indicator solution (phenolphthalein, bromthymol blue and methyl red dissolved in ethanol and water) changes color at each integral pH value

73. Other pH indicators Litmus and phenolphthalein are indicators
Red cabbage has a pigment that changes colors at different pH values

74. Buffers
Buffers are solutions that don’t change in pH when acids or bases are added.
They use weak acids/bases and Le Châtelier’s principle.
WA = weak acid
HWA + H2O  H3O+ + WA-

75. How? pH is determined by the concentration of H3O+
Concentration is measured by mol /L
Moles of H3O+ / L (primarily of) H2O

76. What it does adding H3O+ should increase [H3O+]
However, this forces the equation to the left, decreasing H3O+ and increasing H2O
so the [H3O+] remains constant
Removing H3O+ (adding a base) should decrease [H3O+]
However, this forces the equation to the right, increasing H3O+ and decreasing H2O
So again, there is no change to [H3O+]
There is a breaking point where the pH will change.

77. What does this have to do with my life?
Your blood is a buffered solution
The pH must remain between
Outside of that range can kill you
below this range is called acidosis
above is called alkalosis

78. Buffered Products medications (Bufferin) Shampoos, body soaps
All are buffered to be near your body pH so they won’t cause a major disruption.