1. Chapter 14 Acids and Bases
Bases are sometimes referred to as alkalis because strong bases are made with alkali metals.

2. Arrhenius: 1st to Define Acids
Acids H+(aq)
Bases  OH-(aq)
Limitations:
Only aqueous solutions
Only one kind of base: OH-
NH3 ammonia could not be an Arrhenius base.
Acids produce H ions in aqueous solutions, bases produce OH ions in aqueous solutions. Limitation are there is only 1 kind of base and other substances like ammonia could not be bases.
Acid: Substance that, when dissolved in water, increases the concentration of hydrogen ions, H+ or H3O+
Base: Substance that, when dissolved in water, increases the concentration of hydroxide ions, OH-
1883

3. Brønsted-Lowry Definitions
Acid is proton (H+) donor, must have a removable (acidic) proton
Base is proton acceptor, must have a pair of nonbonding electrons
Acids and bases always come in pairs.
Example HCl(g) + H2O(l) ⇌ H3O+ + Cl-
Water as base becomes an Hydronium ion
Remember – when acids are added to water the process is Dissociation
12.0 M HCl is 37% by mass

4. Bronsted Lowry This is equilibrium
Acid
Base
Conjugate acid
Conjugate base
This is equilibrium
HA & A- are generic term for Bronsted Lowry acid & base
Competition for H+ between H2O and A-
In any acid-base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base.
If H2O is a stronger base it takes the H+
p nd paragraph
Equilibrium moves to right.
Bronsted is Danish and Lowry is English co-discoverers in 1923
In every acid-base reaction the position of the equilibrium favors transfer of the proton to the stronger base
BL 11e p.669 last paragraph

5. Acid Dissociation Constant Ka
HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq)
The equilibrium expression is:
Ka is the acid dissociation constant
H3O+ is often written H+
Ignore the water in equation because it’s [l] remains constant (p. 625)
In Ch. 13 concentration of pure solids and pure liquids are omitted in the equilibrium expression. In a dilute solution we can assume the concentration of water remains essentially constant when an acid is dissolved. Thus the species water is not included in the expression for Ka p.625

6. Acid Dissociation Constant Ka
HA(aq) ⇆ H+(aq) + A-(aq)
We can write the expression for any acid.
Strong acids dissociate completely.
Equilibrium far to right in strong acids
Conjugate base must be weak.

7. Write the dissociation Reaction omitting water for the following:
Example Hydrochloric acid:
HCl(aq) ⇔ H+(aq) + Cl-(aq)
Acetic Acid HC2H3O2
Ammonium ion
Hydrated aluminum (III) ion [Al(H2O)6]+3
Sample exercise 14.1

8. Various Ways to Describe Acid Strength
Strength of acid is defined by the equilibrium position of its dissociated reaction.
Property
Strong Acid
Weak Acid
Ka Value
Large
Small
Position of dissociation equilibrium
Far 
Far ←
Equilibrium Comparison [H+] : [HA]o
[H+] ≈ [HA]o
[H+] << [HA]o
Strength of Conjugate Base Compared to H2O
A- much weaker
A- much stronger
Table 14.1 p.627
Table 14.2 Values for some Ka values
See figure 14.4 p. 626

9. Acid and Base Strength
Strong acids are completely dissociated in water.
Their conjugate bases are quite weak.
Weak acids only dissociate partially in water.
Their conjugate bases are weak bases.
Substances with negligible acidity do not dissociate in water.
Their conjugate bases are exceedingly strong.
From Brown LeMay

10. Strong Acids Strong Soluble Bases
Perchloric Acid >3.2 x 109
Hydroiodic Acid 3.2 x 109
Hydrobromic Acid 1.0 x 109
Hydrochloric Acid 1.3 x 106
Sulfuric Acid 1.0 x 103
Chloric Acid 1.0 x 103
Nitric Acid 2.4 x 101
Lithium Hydroxide
Sodium Hydroxide
Potassium Hydroxide
Rubidium Hydroxide
Cesium Hydroxide
Calcium Hydroxide
Strontium Hydroxide
Barium Hydroxide
Source Brown LeMay
KNOW THESE

11. Rules for Writing and Naming Acids review p.67
All inorganic acids starting with H
Binary Acids
Two elements or H plus non-oxy polyatomic anion
Name comes from root of anion
Begins w/ “hydro-” & ends “-ic”
Ex. HCl: hydrochloric acid
Ex. HCN: hydrocyanic acid

12. Rules for Writing and Naming Acids
Oxyacids
Anion contains Oxygen
Named from root name of anion
Use suffix “ic” or “-ous”
–ate Anion becomes -ic acid
–ite Anion becomes -ous acid
Ex. Nitrate ⇒ HNO3 is Nitric Acid
Ex. Nitrite ⇒ HNO2 is Nitrous Acid

13. Types of Acids
Polyprotic Acids- more than 1 acidic hydrogen (diprotic, triprotic)
Oxyacids - Proton attached to oxygen of ion
Organic acids contain the Carboxyl group -COOH with the H attached to O
See Table 22.4 p 1011
Generally very weak
Table 14.2 Common Monoprotic Acids – note strong acids not listed
See space filled models of acids p.627 the hydrogen atom is attached to oxygen part of polyatomic ion. Strong acids not listed on 14.2 because their Ka is can not be measured accurately see text

14. Amphoteric: Behave As Acid & Base
Example: Autoionization of Water (most common)
2H2O(l) ⇆ H3O+(aq) + OH-(aq)
Kw= [H3O+][OH-]=[H+][OH-]
KW is dissociation constant for water
In pure 25ºC KW = 1.0 x10-14
In EVERY aqueous 25oC, no matter what it contains, the product of [H+][OH-] is always 1.0 x10-14
room temperature 2 out of 109 molecules are ionized
NH3 + H2O  NH4+ + OH- water is acid
HCl + H2O  H3O+ + Cl water is base

15. Meaning of Kw In each case: KW = 1.0X10-14
In any aqueous solution, we know KW = 1.0X10-14
There are 3 possible situations:
Neutral solution [H+] = [OH-]= 1.0 x10-7
Acidic solution [H+] > [OH-]
Basic solution [H+] < [OH-]
In each case: KW = 1.0X10-14

16. Calculate [H+] or [OH-]
Calculate [H+] for the following solution at 25oC and state whether the solution is neutral, acidic or basic:
[OH-] = 1.0X10-5M
Answer:
KW=[H+] [OH-]=1.0X10-14 ,
since [OH-] = 1.0X10-5M, solving for [H+] gives:
Since: 1.0X10-5 [OH-] > 1.0X10-9 [H+] the solution is basic
Carnegie Time #38 p.674

17. The pH Scale pH= -log[H+]
Use pH scale because [H+] is usually very small
Sig Figs Rule: Number of decimal places in log = number of sig figs in original number
[H+] = 1.0 x pH= 8.00 (2 decimals in pH)
Similar for other quantities:
pOH= -log[OH-]
pK = -log K
2 sig figs
p. 631

18. These are the pH values of several common substances.

19. Lots of Relationships pH = -log[H+] pOH = -log[OH-] pK = -log K
KW = [H+][OH-]
log KW = log [H+]+ log[OH-]
-log KW = -log[H+] - log[OH-]
pKW = pH + pOH
Since KW = 1.0 x10-14
then pKW = -log (1.0 x10-14 )=14.00
pH + pOH = 14.00
Given any one of these values: [H+], [OH-] ,pH, and pOH we can find the other three.

20. 100
10-1
10-3
10-5
10-7
10-9
10-11
10-13
10-14
[H+] 1 3 5 7 9 11 13 14 pH
Acidic
Neutral
Basic 1 3 5 7 9 11 13 14
pOH
Basic
100
10-1
10-3
10-5
10-7
10-9
10-11
10-13
10-14
[OH-]

21. Calculating pH, pOH, [H+], & [OH-]
The pH of a sample of blood was measured to be 7.41 at 25oC. Calculate the pOH, [H+], and [OH-] for the sample. Since pH + pOH = pOH = – pH = – 7.41 = 6.59 Since pH = -log[H+] 7.41 = -log[H+] take the antilog of -7.41=3.9x10-8 Similarly for pOH = -log[OH-] 6.59 = -log[OH-] take the antilog of -6.59=2.6x10-7 Do #44 p. 674
To calculate [H+] on TI-36x given the pH  2nd yx 3rd sci = x = 3.9 x 10-8
To calculate pOH on TI-36x given the pH = 1.0 x 10-7 – pH = pOH
To calculate pH on TI-36x given the [H+] pH = - log [H+] x log  ( got to find the log of the number and then change its sign)

22. Calculating pH of Solutions
Always write down the major species in solution, this is first step in all problems
Remember these are at equilibrium.
Remember the chemistry.
Don’t just memorize

23. Calculate pH of Strong Acid Solutions
Common Strong Acids: HCl, HNO3, H2SO4, HClO4
When working with a Strong Acid (SA) solution, focus on solution species
Determine which species are significant and which can be ignored.
MAJOR SPECIES = COMPONENTS PRESENT IN RELATIVELY LARGE AMOUNTS.
Ex. 1.0M HCl Calculate pH
p nd full paragraph

24. Calculate pH of Strong Acid Solutions
No HCl in solution because all SA completely dissociate
Focus on Major Species: H+ Cl H2O
Because want pH, focus on H+
Consider H+ from H2O, the dissociation of HCl drives H+ contribution from water to left, so H+ from H2O is negligible
Therefore [H+] is 1.0 and pH = -log[1.0] = 0

25. Calc. pH of 0.10 M HNO3 Solution:
HNO3 strong acid, Major species: H+, NO3-, H2O
[HNO3] virtually 0 because SA
[OH-] very small because H2O(l)  [H+](aq) + [OH-](aq) is driven to left by H+ from acid (Le Chatelier’s Principle)
Source of H+ are:
HNO3
H2O, negligible from H2O (see #3 above)
Leaves HNO3 as source of H+ so:
[H+]=0.10 M and pH = -log(0.10) = 1.00
CHT #48 p.674

26. Solving Weak Acid Equilibrium Problems
List major species
Choose the species that can produce H+ & write balanced equation for reactions producing H+
Ka vs. Kw decide which equilibrium will dominate in producing H+
Write Equilibrium Expression for dominate equilibrium.
Build ICE Table: List the Initial conc. of species participating in dominate equilibrium
Define change needed to achieve equilibrium in terms of x
Initial conc. + Change = Equilibrium Concentration
Sub. Equil. Conc. into equilibrium expression #4
Solve for x using Approximation Rule
Calculate [H+] & pH
p. 637

27. Example Study Sample Exercise 14.8 p. 638
Pay close attention to comments between equations and calculations
Step 5-7 is nICE
Step 9-10 is the Approximation Rule (Ch. 13 p.604)
You will use these steps in ALL problems in next couple of chapters, so learn them well
Do # 53 p.675

28. Find the pH of a 0.100 M HClO2 solution See Solution Handout
Finding The pH Of A Weak Acid Solution in Cases Where “x is small Approximation” Does Not Work
Find the pH of a M HClO2 solution See Solution Handout
Chemistry A Molecular Approach by Nivaldo J. Tro p.680
Ka = 1.1 x 10-2

29. The initial concentration and Ka’s of several weak acid solutions are listed below. For which of these is the “x is small approximation” least likely to work in finding the pH of the solution?
Initial [HA] = 0.100M Ka = 1.0 x 10-5
Initial [HA] = 1.00M Ka = 1.0 x 10-6
Initial [HA] = M Ka = 1.0 x 10-3
Initial [HA] = 1.0M Ka = 1.5 x 10-3
The validity of the x is small approximation depends on both the value of the equilibrium constant and the initial concentration - the closer that these are to one another, the less likely the approximation will be valid c)
Chemistry A Molecular Approach by Nivaldo J. Tro p.680

30. pH of a Mixture of Weak Acids
Calc. pH of a solution that contains M HCN (Ka=6.2x10-10) and 5.00 M HNO2 (Ka=4.0x10-4). Also calc. the concentration of the (CN)- at equilibrium.
Major Species:
HCN, HNO2, andH2O
All 3 produce H+
HCN Ka=6.2x10-10
HNO2 Ka=4.0x10-4
H2O Kw=1.0x10-14
Pick the strongest of the weak acids
Sample Exercise 14.9 p.639

31. pH of a Mixture of Weak Acids
HNO2 is much stronger than the other 2 species. Thus HNO2 is assumed to be the dominant producer of H+
Set up ICE Table and write the Equilibrium Expression
Plug Equilibrium Concentrations into Expression and using the Approximation Rule, solve for x
x is the [H+] and you can use it to find [CN-]
Sample Exercise 14.9 p remainder of the problem is explained on the bottom of 640 and the top of 641
Answer: [CN-] = 1.4 x 10-8 M

32. Weak acid: percent dissociation increases as acid becomes more dilute.
Calculate the percent dissociation of acetic acid (Ka=1.8x10-5) for a 1.00M solution.
Solution:
List major species:
HC2H3O H2O
Choose species that produces most H+ : HC2H3O2 & write balanced equation: HC2H3O2(aq)  H+(aq) + C2H3O2-(aq)

33. Percent Dissociation Write Equilibrium Expression
Create ICE chart & fill in
Plug equilibrium concentrations into equilibrium expression & solve for X (use “x is small approximation rule”)
X = 4.2 X 10-3 M
HC2H3O2(aq)
H+(aq)
C2H3O2-(aq)
Initial Conc.
1.00
Change -X X
Equilibrium
1.00-X

34. Study this sample & CHT #65 p.675
Percent Dissociation
Calculate percent dissociation:
General Trend: For weak acid solutions Fig 14.10
[acid] and [H+] decrease as % dissociation increases
% dissociation increase as acid is diluted
Calculating Ka from percent dissociation is illustrated in Sample Exercise on p.643
Study this sample & CHT #65 p.675

35. Study Sample Exercise 14.12 p.645
Bases
Remember the OH- list of strong bases.
Hydroxides of the alkali metals are strong bases because they dissociate completely when dissolved.
The hydroxides of alkaline earths (Ex. Ca(OH)2) are strong dibasic (2 OH-) bases, but they don’t dissolve well in water. Not very soluble.
Many metal hydroxides are used as antacids because [OH- ] can’t build up.
Study Sample Exercise p.645
Carnegie #77 p.675

36. B(aq) + H2O(l) ⇆ BH+(aq) + OH- (aq)
Bases without OH-
Bases are proton acceptors (Brønsted-Lowry).
Example: NH3 + H2O ⇆ NH4+ + OH-
NH3 is a base. It is the lone pair on nitrogen that accepts the proton, H2O is acid
Many weak bases contain N
General Reaction:
B(aq) + H2O(l) ⇆ BH+(aq) + OH- (aq)
Base Acid Conj. Acid Conj. Base

37. Strength of Bases Equilibrium constant is Kb
Hydroxides bases are strong bases.
Others are weak.
Smaller Kb = weaker base.
Calculate the pH for 15.0 M solution of NH3 (Kb=1.8x10-5)
Refer to steps for solving weak acid handout
Create ICE table
Substitute conc. Into equilibrium expression
Sample exercise p. 647

38. Polyprotic Acids Polyprotic acids furnish more than 1 proton
Always dissociate stepwise i.e. 1 proton at a time
The first H+ comes off much easier than the second.
Denoted Ka1, Ka2, Ka3
Ka1 is much bigger than Ka2
For typical weak acid Ka1 > Ka2 > Ka3
Look at Table 14.4 on p. 651 for Ka values

39. Polyprotic Acids Carnegie #95 p. 676
Calculate the pH of a 5.0 M H3PO4 solution and the equilibrium concentration of the species H3PO4, H2PO4-, HPO4-2, and PO4-3
List major species, write balanced equations
Use Ka & Kw to decide dominant equilibrium
Write equilibrium expression
Create ICE table
Solve for x
Repeat for each successive species
Carnegie #95 p. 676

40. Acid-Base Properties of Salts
Salts are ionic compounds.
Cations of strong bases (Na+ & K+) have no effect on [H+]
Anion of strong acids (Cl- & NO3-) have no effect on [H+]
Thus these salts NaCl, KNO3 have no effect on [H+]
Aqueous solutions of these salts are neutral (pH = 7)
Anion of weak acid yield strong conjugate base, solution will be basic
Alkali metal and Alkaline Earth Metal cations have no effect
Metal cations with +2 or +3 charges will produce weakly acidic solutions
p. 660 Table 14.6

41. Acid-Base Properties of Salts
Problem
Decide whether a salt is acidic, basic, or neutral
Data / Information
ID the cation and anion
ID the nature of Cation
ID nature of Anion
Decide on acid-base properties of salt

42. Acid-Base Properties of Salts
Practice: Decide whether each of the following will give rise to an acidic, basic or neutral solution with water.
NaNO3
K3PO4
FeCl2
NaHCO3
NH4F
NaNO3 neutral pH=7 Na Alkali metal, NO3 weak conjugate base
K3PO4 basic pH>7 K Alkali metal, PO4 strong conjugate base
FeCl2 weakly acidic pH<7 Fe+2 will form a Bronsted acid, Cl- weak conjugate base
NaHCO3 depends on relative magnitude of Ka and Kb IF Ka>Kb Acidic IF Ka<Kb Basic
NH4F depends on relative magnitude of Ka and Kb IF Ka>Kb Acidic IF Ka<Kb Basic
Kotz p. 599

43. Add this to your formula list
Basic Salts
If salt solution:
Has neutral cation
And anion is conjugate base of a weak acid
is basic
Reverse is true for acidic salts
If you know one (Ka or Kb ) you can calculate the other because you always know KW
Ka x Kb = KW
Add this to your formula list

44. Salts as Weak Base Study Sample Exercise 14.18 p. 656
The anion of a weak acid is a weak base.
Calculate the pH of a 0.30 M NaF solution.
KaHF = 7.2 x 10-4
List major species
Write equilibrium reaction
Calc. Kb expression from Ka x Kb = Kw
Create ICE table
Solve for x
Calculate pH
Do #103 p. 676
Sample exercise p.656

45. Acidic salts
Salts in which anion is not a base and cation is conjugate acid of weak base produce acidic solution
The same development as bases leads to
Ka x Kb = KW
Sample Salt as Weak Acid 0.10 M NH4Cl , Kb of NH3 1.8 x 10-5
Other acidic salts are those of highly charged metal ions, Example Al+3
Sample exercise p.658

46. Hydrated metals
Highly charged metal ions pull the electrons of surrounding water molecules toward them.
Make it easier for H+ to come off. H
Al+3 O H

47. Qualitative Predictions Salts w/ Acidic & Basic Properties Compare Ka for acidic ion with Kb for basic ion
IF Ka > Kb acidic
IF Ka < Kb basic
IF Ka = Kb Neutral

48. Effect of Structure on Acid-base Properties 14.9
Any molecule with an H atom is a potential acid
Two Factors will determine if molecule w/ H-X is acid:
Strength of Bond - The stronger the H-X bond the less acidic (Table 14.7)
Polarity: The more polar the H-X bond the stronger the acid (electronegativity) The more polar H-O-X bond -stronger acid.

49. HClO4 > HClO3 > HClO2 > HClO
Strength of Oxyacids
The more oxygen hooked to the central atom, the more acidic the hydrogen.
HClO4 > HClO3 > HClO2 > HClO
Remember that the H is attached to an oxygen atom. (Table 14.8)
Oxygen is electronegative & pull electrons away from hydrogen allowing it to dissociate more completely

50. Strength of oxyacids
Electron Density Cl O H

51. Strength of oxyacids
Electron Density O Cl O H

52. Strength of oxyacids
Electron Density O Cl O H O

53. Strength of oxyacids
Electron Density O O Cl O H O

54. Acid-Base Properties of Oxides
Non-metal oxides dissolved in water
SO3 (g) + H2O(l)  H2SO4(aq)
If X-O is strong & covalent it will produce an acid. Did you ever cry while slicing an onion?
CaO(s) + H2O(l)  Ca(OH)2(aq)
If X-O is ionic produces bases.

55. Lewis Acids and Bases :N F H B F H F H Most general definition.
Acids are electron pair acceptors.
Bases are electron pair donors. F H B F :N H F H
Lewis Acid
Lewis Base

56. Lewis Acids and Bases :N F H B F H F H
Boron trifluoride wants more electrons. F H B F :N H F H

57. Lewis Acids and Bases F H F B N H F H
Boron trifluoride wants more electrons.
BF3 is Lewis Acid NH3 is a Lewis Base F H F B N H F H

58. Lewis Acids and Bases ( ) H
Al+3
+ 6 O H +3 ( ) 6 H Al O H

59. Summary of 3 Models of Acids & Bases
Definition of Acid
Definition of Base
Arrhenius
H+ producer
OH- producer
Bronsted-Lowry
H+ donor
H+ acceptor
Lewis
Electron-pair acceptor
Electron-pair donor