1. Unit 9 Chemical Equilibrium & Acid-Base Chemistry

2. Reversible Reactions
A chemical reaction in which the products can react to re-form the reactants is called a reversible reaction.
A reversible reaction is written using double arrows to show that the reaction is proceeding in both directions.
Example:

3. Dynamic Equilibrium
A reversible reaction reaches dynamic equilibrium when the rate of its forward reaction equals the rate of its reverse reaction and the concentrations of its products and reactants remain unchanged.
At equilibrium, both reactions continue, but there is no net change in the composition of the system.
Visual Concept

4. Equilibrium  Equal
At equilibrium, the rates of the forward and reverse reactions are equal. But the concentrations aren’t necessarily equal.
Some reactions reach equilibrium only after almost all reactants are consumed (products are favored.)
Others reach equilibrium when only a small percentage of reactants are consumed (reactants are favored.)

5. Le Châtelier’s Principle
Le Châtelier’s Principle: When a system at equilibrium is disturbed, the system shifts in a direction that minimizes the disturbance.
A shift in equilibrium will result from a change to any of the following:
Concentration
Volume/Pressure
Temperature

6. Change in Concentration
If you increase the concentration by adding more of a reactant or product, the system will shift to produce less of that substance.
If you lower the concentration by removing some of a reactant or product, the system will shift to produce more of that substance.

7. Change in Volume/Pressure
When you increase the pressure (usually by decreasing volume of the container), the system shifts so the least number of gas molecules are formed (less collisions = lower pressure.)
When you decrease the pressure, the system will shift so the greatest number of gas molecules are formed.

8. Change in Temperature
For every reversible reaction, one direction is endothermic and the other is exothermic.
If the temperature is increased, the endothermic reaction will be favored (because it takes in some of the excess heat.)
If the temperature is decreased, the exothermic reaction will be favored (produces heat.)

9. Le Châtelier’s Principle Sample Problem
2 SO2(g) + O2(g) Û 2 SO3(g) DH° = -198 kJ
How will the reversible reaction above shift in response to each of the following stresses?
adding more O2 to the container
condensing and removing SO3
compressing the gases
cooling the container
doubling the volume of the container
warming the mixture
Shift right
Shift right
Shift right
Shift right
Shift left
Shift left

10. The Law of Mass Action [C]c [D]d K = [A]a [B]b
The relationship between the chemical equation and the concentrations of reactants and products is called the Law of Mass Action.
for the general equation aA + bB  cC + dD, the Law of Mass Action is:
Lowercase letters represent coefficients.
Always products over reactants.
Pure solids and pure liquids are not included.
[C]c [D]d
K =
[A]a [B]b

11. The Equilibrium Constant
The equilibrium constant (K) reflects how the concentrations of the reactants and products compare at equilibrium.
It can also be a ratio of pressures (in atmospheres) if a reaction involves gaseous reactants and/or products.
K is unitless.

12. The Value of K
K > 1 more product molecules present than reactant molecules (the position of equilibrium favors products.)
K < 1 more reactant molecules present than product molecules (the position of equilibrium favors reactants.)
K = 1 reactant and product particles are present in exact equal concentrations at equilibrium.

13. Equilibrium Constant Sample Problem
Equilibrium concentrations of [H2] = M, [I2] = 0.53 M and [HI] = M were observed at 445oC for the reaction:
H2(g) + I2(g) Û 2 HI(g)
Write an equilibrium expression for the above reaction.
2. Calculate the value of Kc for this reaction at 445oC.
[HI]2
K =
[H2][I2]
[HI]2
[0.934]2
Kc = = =
49.9
[H2][I2]
[0.033]
[0.53]

14. The Reaction Quotient [C]c [D]d Q = [A]a [B]b
When a reaction is not at equilibrium, how do you know in which direction it will proceed?
the answer is to compare the equilibrium constant to a ratio of current concentrations called the reaction quotient (Q).
for the general equation aA + bB  cC + dD:
The non-equilibrium concentrations (or pressures) are used.
[C]c [D]d
Q =
[A]a [B]b

15. Q vs. K We calculate Q in order to compare it with K.
Q < K means the reaction will proceed in the forward direction ([products] increase and [reactants] decrease.)
Q > K means the reaction will proceed in the reverse direction ([products] decrease and [reactants] increase.)
Q = K means the reaction is at equilibrium ([products] and [reactants] will not change.)

16. Q, K, and the Direction of Reaction

17. Reaction Quotient Sample Problem
For the reaction below, which direction will it proceed if PI2 = atm, PCl2 = atm & PICl = atm?
I2(g) + Cl2(g)  2 ICl(g) Kp = 81.9
First calculate Q:
Then, compare it with K:
(ICl)2
(0.355)2
Q = = =
10.8
(I2)(Cl2)
(0.114)
(0.102) Q
< K
Reaction will proceed to the right
10.8
81.9

18. Properties of Acids Taste sour. React with metals to release H2 gas.
React with bases to produce salts and water.
Change the color of acid-base indicators.
Conduct electric current.

19. Properties of Bases Taste bitter. Feel slippery.
React with acids to produce salts and water.
Change the color of acid-base indicators.
Conduct electric current.

20. Arrhenius Acids and Bases
An Arrhenius acid produces hydrogen ions, H+, in aqueous solution.
An Arrhenius base produces hydroxide ions, OH−, in aqueous solution.
A strong acid (or base) ionizes completely.
A weak acid (or base) releases only a few ions.

21. Arrhenius Theory HCl ionizes in water, producing H+ and Cl– ions
NaOH dissociates in water,
producing Na+ and OH– ions

22. Hydronium Ion
The H+ ions (protons) produced by the acid are so reactive they cannot exist in water.
instead, they react with a water molecule to form a hydronium ion, H3O+.
H+ + H2O  H3O+
Chemists use H+ and H3O+ interchangeably.

23. Arrhenius Acid-Base Reactions
The H+ from the acid combines with the OH- from the base to make a molecule of H2O.
It is often helpful to think of H2O as H-OH.
The cation from the base combines with the anion from the acid to make a salt.
acid + base → salt + water
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

24. Brønsted-Lowry Acids and Bases
In a Brønsted-Lowry Acid-Base reaction, an H+ ion (proton) is transferred.
Does not have to take place in aqueous solution.
Broader definition than Arrhenius.
A Brønsted-Lowry acid is a molecule or ion that is a proton donor.
A Brønsted-Lowry base is a molecule or ion that is a proton acceptor.

25. Conjugate Pairs Brønsted-Lowry theory allows for reversible reactions.
The original base has an extra H+ after the reaction. It will act as an acid in the reverse process.
The original acid has a lone pair of electrons after the reaction. It will act as a base in the reverse process.
each reactant and the product it becomes is called a conjugate pair.

26. Brønsted-Lowry Acid-Base Reactions Sample Problem
Identify the Brønsted-Lowry Acids and Bases and Their Conjugates in the Reactions below: a. b.
H2SO H2O  HSO4– H3O+
acid
base
conjugate base
conjugate acid
HCO3– H2O  H2CO3 + HO–
base
acid
conjugate acid
conjugate base

27. Amphoteric Compounds
An amphoteric substance is one that can react as either an acid or a base.
Example: water
Water can act as an acid.
Water can act as a base.
base acid
acid base
Visual Concept

28. Polyprotic Acids
Molecules with more than one ionizable H are called polyprotic acids.
1 H = monoprotic, 2 H = diprotic, 3 H = triprotic (Ex: HCl = monoprotic, H2SO4 = diprotic, H3PO4 = triprotic)
Polyprotic acids ionize in steps (each ionizable H removed sequentially.)
Removing the first H automatically makes removing the second H harder. (Ex: H2SO4 is a stronger acid than HSO4)

29. Lewis Acids and Bases
A Lewis acid is an atom, ion, or molecule that accepts an electron pair to form a covalent bond.
A Lewis base is an atom, ion, or molecule that donates an electron pair to form a covalent bond.
The Lewis definition is the broadest of the three definitions.

30. Comparing the Three Definitions
Visual Concept

31. Strong and Weak Acids and Bases
Strong acid – fully dissociates in water (almost every molecule breaks up to form H+ ions.
Weak acid – partially dissociates in water.
Strong base – fully dissociates in water (almost every molecule breaks up to form OH- ions.
Weak base – partially dissociates in water.
A Strong Acid
A Weak Acid

32. Ionization Constants for Weak Acids and Bases
The acid ionization constant (Ka) is the equilibrium constant for the ionization reaction of a weak acid:
A base ionization constant (Kb) can also be created for the ionization reaction of a weak base:
HA(aq) + H2O(l)  H3O+(aq) + A-(aq)
[H3O+] [A-]
Ka =
[HA]
B(aq) + H2O(l)  BH+(aq) + OH-(aq)
[BH+] [OH-]
Kb =
[B]

33. Acid Ionization Constant Sample Problem
Calculate the Ka of a M solution of acetic acid with a measured [H3O+] of 1.34 x 10-3 M.
HC2H3O2(aq) + H2O(l)  H3O+(aq) + C2H3O2-(aq)
For every H3O+ produced, there is also a C2H3O2- produced, so the concentrations must be the same.
The equilibrium concentration of original acid is the original concentration decreased by the amount ionized (0.100M – 1.34 x 10-3 M = M)
[H3O+][C2H3O2-]
[1.34 x 10-3M]
[1.34 x 10-3M]
Ka= = =
1.82 x 10-5
[HC2H3O2]
[ M]

34. Ka & Kb and Strength of Acids/Bases
The strength of an acid or base is measured by the size of its equilibrium constant when it reacts with H2O.

35. Six Strong Acids and Bases
Because these acids and bases are known to dissociate (ionize) to essentially 100% completion, it is meaningless to connect them to equilibrium:
6 Strong Acids
6 Strong Bases
HCl – Hydrochloric Acid
Ca(OH)2 – Calcium Hydroxide
HBr – Hydrobromic Acid
Sr(OH)2 – Strontium Hydroxide
HI – Hydroiodic Acid
Ba(OH)2 – Barium Hydroxide
H2SO4 – Sulfuric Acid
LiOH – Lithium Hydroxide
HNO3 – Nitric Acid
NaOH – Sodium Hydroxide
HClO4 – Perchloric Acid
KOH – Potassium Hydroxide

36. Strengths of Conjugate Pairs
the stronger an acid is at donating H, the weaker the conjugate base is at accepting H. (i.e. strong acids have weak conjugate bases, and weak acids have strong conjugate bases.)
Increasing Basicity
Increasing Acidity

37. Predicting Acid Strength
The strength of an acid depends on its tendency to ionize (let go of its hydrogen.)
For binary acids, the strength of the acid depends on two factors:
The stronger the bond, the weaker the acid.
The more polar the bond, the stronger the acid.

38. Periodic Trends are a BEAR
All of these increase in the direction toward their letter:
B=Basicity E=Electronegativity, ionization Energy, & Electron Affinity A=Acidity R=Radius B E A R

39. Predicting Acid Strength (continued)
The strength of Oxyacids of the form H-O-Y, where Y is any atom (besides H) bonded to O, depends on two factors:
The more electronegative the element Y, the stronger the acid.
The greater the number of oxygen atoms bonded to Y, the stronger the acid.

40. Predicting Acid Strength Sample Problem
Predict the relative strengths of the following acids:
HCl, HBr, and HI
HClO, HBrO, and HIO
HNO3 and HNO2
Predict the relative strengths of the following bases:
Cl-, Br- and I-
H2PO3- and H2PO4-
LiOH and Mg(OH)2
HCl < HBr < HI
HIO < HBrO < HClO
HNO2 < HNO3
I- < Br- < Cl-
H2PO4- < H2PO3-
LiOH < Mg(OH)2

41. Autoionization of Water
Water is actually an extremely weak electrolyte.
About 1 out of every 10 million water molecules form ions through a process called autoionization.
H2O + H2O Û H3O+ + OH–
In pure water at 25°C, [H3O+] = [OH–] = 10-7M.

42. Ion Product of Water Kw = [H3O+] [OH-] = 1 x 10-14
The product of the H3O+ and OH– concentrations is always the same number, called the ion product of water (Kw).
at 25°C
If you measure one of the concentrations, you can calculate the other.
As [H3O+] increases the [OH–] must decrease so the product stays constant.
Kw =
[H3O+] [OH-]
= 1 x 10-14

43. Acidic and Basic Solutions
Neutral solutions have equal [H3O+] and [OH–]
[H3O+] = [OH–] = 1 x 10-7
Acidic solutions have a larger [H3O+] than [OH–]
[H3O+] > 1 x 10-7; [OH–] < 1 x 10-7
Basic solutions have a larger [OH–] than [H3O+]
[H3O+] < 1 x 10-7; [OH–] > 1 x 10-7

44. Acidic, Basic or Neutral? Sample Problem
Calculate [OH] at 25°C when [H3O+] = 1.5 x 10-9 M, and determine if the solution is acidic, basic, or neutral.
First calculate [OH-]:
Then, compare [H3O+] with [OH-]:
Kw =
[H3O+][OH-] Kw
1.0 x 10-14
[OH-] = = =
6.7 x 10-6 M
1.5 x 10-9
[H3O+]
<
[H3O+]
[OH-]
The solution is basic
1.5 x 10-9 M
6.7 x 10-6 M

45. The pH Scale
The acidity/basicity of a solution is often expressed as pH.
pH is defined as the negative of the common logarithm of the hydronium ion concentration.
pH = −log [H3O+]
pH < 7 is acidic; pH > 7 is basic, pH = 7 is neutral.

46. pH of Some Common Substances

47. The pH Scale Sample Problem
Calculate the pH at 25°C when the [OH] = 1.3 x 10-2 M, and determine if the solution is acidic, basic, or neutral
First calculate [H3O+]:
Then, calculate pH:
Kw =
[H3O+][OH-] Kw
1.0 x 10-14
[H3O+] = = =
7.7 x M
1.3 x 10-2
[OH-]
pH = −log [H3O+]
= -log(7.7 x M)
pH =
12.1
The solution is basic

48. Finding the Ionization Constant from pH Sample Problem
A M weak acid (HA) solution has a pH of Find Ka for the acid.
First calculate [H3O+]:
Then, calculate Ka:
pH =
-log [H3O+]
10-pH =
5.6 x 10-5 M = =
[H3O+]
HA(aq) + H2O(l)  H3O+(aq) + A-(aq)
[H3O+] [A-]
(5.6 x 10-5)
(5.6 x 10-5)
Ka = = =
3.1 x 10-8
[HA]
( x 10-5)

49. The pOH Scale
Another way of expressing the acidity/basicity of a solution is pOH.
pOH is defined as the negative of the common logarithm of the hydroxide ion concentration.
pOH = −log [OH-]
pOH < 7 is basic; pOH > 7 is acidic, pOH = 7 is neutral

50. Relationship Between pH and pOH
The sum of the pH and pOH of a solution is 14.
pH + pOH = 14.0

51. Neutralization Reactions
A neutralization reaction is a double displacement reaction in which an acid and a base in an aqueous solution react to produce a salt and water.
A salt is an ionic compound made up of a cation from a base and an anion from an acid.

52. Acid Base Properties of Salts
Salts of strong acids and strong bases are neutral.
Ex: HCl(aq) + NaOH (aq)  NaCl (aq) + H2O (l)
Salts of strong acids and weak bases are acidic.
Ex: NH3(aq) + HCl(aq)  NH4Cl (aq)
Salts of strong bases and weak acids are basic
Ex: 2NaOH(aq)+ H2CO3 (aq)  Na2CO3 (aq) + 2H2O(l)
Salts of weak acids and weak bases can be acidic, basic or neutral depending on the relative strength of acids and bases.

53. Buffers
Buffers are solutions that resist changes in pH when an acid or base is added.
Buffers contain both a weak acid and its conjugate base (or a weak base and its conjugate acid.)
The weak acid neutralizes added base.
The conjugate base neutralizes added acid.

54. Buffering Effectiveness
A good buffer should be able to neutralize moderate amounts of added acid or base.
However, there is a limit to how much can be added before the pH changes significantly.
The buffering capacity is the amount of acid or base a buffer can neutralize.
The buffering range is the pH range over which the buffer can be effective.
The effectiveness of a buffer depends on:
The relative amounts of acid and base.
The absolute concentrations of acid and base.

55. Buffers in Human Blood
Many of the chemical reactions that occur in the body are pH-dependent.
Ideally, the pH of the blood should be maintained at a slightly basic 7.4.
pH below 6.8 or above 7.8 can be fatal.
Fortunately, we have buffers in the blood to protect against large changes in pH.

56. Acid-Base Indicators
An acid-base indicator is a chemical dye that changes colors at definite pH values.
There are a variety of indicators that change color at different pH levels.
A properly selected indicator can be used to visually "indicate" the approximate pH of a sample.

57. Litmus Paper
A common indicator is found on litmus paper. It is red below pH 4.5 and blue above pH 8.2.

58. Phenolphthalein
Phenolphthalein is an organic compound often used as an acid-base indicator.
Phenolphthalein is colorless in acidic solutions, but turns pink when the pH is greater than 8.3.

59. Bromothymol Blue
Bromothymol Blue (BTB) is a useful indicator for substances that have a relatively neutral pH (near 7).
BTB is yellow in acids, green in neutral solutions, and blue in bases.

60. Acid-Base Titration
Titration is a method for determining the concentration of a solution by reacting a known volume of that solution with a solution of known concentration.
The analyte is a measured volume of an acid or base of unknown concentration.
The standard solution (titrant) is an acid or base solution whose concentration is known.
Standard solution (titrant)
analyte

61. Titration Procedure
In a titration procedure, a measured volume of an analyte is placed in a beaker or flask, and initial pH recorded.
The standard solution (titrant) is filled in a burette.
A couple of drops of an acid-base indicator are added to the flask.
The standard solution is slowly added to the unknown solution in the flask.
As the two solutions are mixed the acid and the base are neutralized.

62. Titration Procedure (cont’d)
As the base is added to the acid, H+ reacts with OH– to form water. But there is still excess acid present so the color does not change.
Once enough base has been added to neutralize all the acid, the indicator changes color.
The difficulty is determining when there has been just enough titrant added to complete the reaction…without going over!

63. The Equivalence Point
End point - The point at which an indicator changes color.
Equivalence point – The point at which the moles of acid added equals the moles of base that you started with (should be the same as the end point.)
An abrupt change in pH occurs at the equivalence point.

64. Acid-Base Titration Sample Problem
The titration of mL of HCl solution of unknown concentration requires mL of M NaOH solution to reach the end point. What is the concentration of the unknown HCl solution? HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
First, write the neutralization reaction
12.54 mL NaOH 1
0.100 1
L NaOH
mol NaOH
mol HCl
1.25 x 10-3 mol HCl = 1
1000
mL NaOH 1
mol NaOH
L NaOH
moles
1.25 x 10-3 mol HCl = =
0.125 M HCl M = L
L HCl

65. Titration of Strong Acid with Strong Base
Since the salt produced is neutral, the solution at the equivalence point has a pH of 7.
the pH starts off low and increases as you add more of the base.
The pH doesn't change very much until you get close to the equivalence point. Then it surges upwards very steeply

66. Titration of Strong Base with Strong Acid
This curve is very similar to the previous one for the titration of a strong acid with strong base.
The main difference is that the curve starts basic and then turns acidic after the equivalence point (rather than vice-versa.)

67. Titration of Weak Acid with Strong Base
The salt is basic, so equivalence point is at a pH > 7.
Before the equivalence point, the solution acts as a buffer. The start of the graph shows a relatively rapid rise in pH but this slows down due to the buffering effect.

68. Titration of Weak Base with Strong Acid
Salt formed is acidic, hence, equivalence point comes at a pH < 7.
This curve is very similar to the titration of a weak acid with a strong base.
The main differences are that the curve starts basic and has an acidic equivalence point.

69. Titration of a Polyprotic Acid
A polyprotic acid titration will have more than one equivalence point.
The first equivalence point represents the titration of the first proton, while the second equivalence point represents the titration of the second proton.

70. Titration Curves Sample Problem
Two acidic solutions were titrated with a strong base. Which curve represents a weak acid and which represents a strong acid?
Strong Acid
Weak Acid

71. Acid Rain
Rain is naturally somewhat acidic (pH ~5.6) due to atmospheric CO2. Carbon dioxide combines with rainwater to form carbonic acid: CO2 + H2O → H2CO3.
Rain water with a pH < 5.6 is called acid rain.
Acid rain is linked to damage in ecosystems and structures.

72. What Causes Acid Rain? Nonmetal oxides such as SO2 and NO2 are acidic:
2 SO2 + O2 + 2 H2O  2 H2SO4
4 NO2 + O2 + 2 H2O  4 HNO3
Processes that produce nonmetal oxide gases increase the acidity of the rain
natural – volcanoes, bacterial action.
man-made – combustion of fuel

73. Sources of SO2 from Utilities
Weather patterns may cause rain to be acidic in regions other than where the nonmetal oxide is produced.

74. Damage from Acid Rain
Acids react with metals, so acid rain damages bridges, cars, and other metallic structures.
Acid reacts with carbonates, so acid rain damages buildings and other structures made of limestone or cement.
Acidified lakes affect aquatic life.
Acid dissolves and leaches minerals from soil, which weakens and kills trees.

75. Acid Rain Legislation
The 1990 Clean Air Act was passed to reduce acid rain.
It requires industries to minimize pollutant gas emissions.
As a result, the acidity of rain in the northeast has stabilized and is beginning to be reduced.