1. Acids and Bases
Chapter 19

2. Describing Acids and Bases

3. Mini-Project
Work with a partner. Organize the following formulas into two groups with four formulas in each group:
HNO3, NaOH, H2SO4, H2CO3, Ca(OH) 2, KOH, H8PO4, Mg(OH) 2
One way to organize them into groups is:
Group One Group Two
HNO3 NaOH
H2SO4 Ca(OH) 2
H2CO3 KOH
H8PO4 Mg(OH) 2
Group One formulas represent acids.
Group Two formulas represent bases.

4. ACIDS Electrolytes in solution Taste sour (lemon, vinegar)
React with metal (corrosion)
React with carbonates (makes bubbles of CO2
Turns blue litmus RED
In Water forms Hydrogen ION
Water
HCl H+ + Cl-

5. BASES Electrolyte in solution Taste Bitter (soap, tonic water)
Feel Slippery (soap)
Turns Red Litmus Blue
React with Acids to make water
Water
NaOH Na+ + OH-

6. Why are Some Solutions Acid & Others Base?
Acid solutions contain more H+ ions than OH- ions.
Base solutions contain more OH- ions than H+ ions.
Water is the standard for Acid/Base and is defined as NEUTRAL
Water has equal amounts of H+ and OH- ions

7. Arrhenius Model of Acids/Bases
Substance is an acid if it contains hydrogen and dissociation causes hydrogen ions to form in solution
Substance is a base if it contains a hydroxide and dissociates to produce hydroxide ions in solution

8. Bronsted-Lowry Model Acid is a proton (hydrogen ion) donor
Base is a proton (hydrogen ion) receptor
This is a broader definition than Arrhenius model because there are substances that cause donation or reception without having hydrogen in them.

9. Example
When an acid dissolves in water, it donates an H+ ion to a water molecule forming H3O+.
The water molecule acts as a base and accepts the H+ ion
HX + H2O ⇄ H3O+ + X-
Conjugate acid = species produced when a base accepts a hydrogen ion from an acid
Conjugate base = species produces when an acid donates a hydrogen ion to a base
Conjugate base pair = 2 substances related to each other by donating and accepting a single hydrogen ion

10. Historical views on acids
O (e.g. H2SO4) was originally thought to cause acidic properties. Later, H was implicated, but it was still not clear why CH4 was neutral.
Arrhenius made the revolutionary suggestion that some solutions contain ions & that acids produce H3O+ (hydronium) ions in solution.
Ionization + H O H O Cl Cl H + +
The more recent Bronsted-Lowry concept is that acids are H+ (proton) donors and bases are proton acceptors

11. The Bronsted-Lowry concept
In this idea, the ionization of an acid by water is just one example of an acid-base reaction. + Cl H O
acid
base
conjugate acid
conjugate base
conjugate acid-base pairs
Acids and bases are identified based on whether they donate or accept H+.
“Conjugate” acids and bases are found on the products side of the equation. A conjugate base is the same as the starting acid minus H+.

12. Practice problems
Identify the acid, base, conjugate acid, conjugate base, and conjugate acid-base pairs:
HC2H3O2(aq) + H2O(l)  C2H3O2–(aq) + H3O+(aq)
acid
base
conjugate base
conjugate acid
conjugate acid-base pairs
OH –(aq) + HCO3–(aq)  CO32–(aq) + H2O(l)
base
acid
conjugate base
conjugate acid
conjugate acid-base pairs

13. Answers: question 18 (a) HF(aq) + SO32–(aq)  F–(aq) + HSO3–(aq) acid
base
conjugate base
conjugate acid
conjugate acid-base pairs
(b)
CO32–(aq) + HC2H3O2(aq)  C2H3O2–(aq) + HCO3–(aq)
base
acid
conjugate base
conjugate acid
(c)
conjugate acid-base pairs
H3PO4(aq) + OCl –(aq)  H2PO4–(aq) + HOCl(aq)
acid
base
conjugate base
conjugate acid
conjugate acid-base pairs

14. More
Amphoteric = substances like water that can act like either an acid or a base

15. Monoprotic and Polyprotic Acids
Monoprotic – acids based on formula that can donate only one hydrogen ions
CH3COOH + H2O ⇄ H3O+ + CH3COO-
Polyprotic – acids that can donate multiple hydrogen ions
H3PO4 + H2O ⇄ H3O+ + H2PO4+
H2PO4+ + H2O ⇄ H3O+ + HPO4+2
HPO4+2 + H2O ⇄ H3O+ + PO4+3
Anhydride = oxides that can become acids or bases by adding elements contained in water

16. Acid Rain
Acid rain comes from rain collecting gasses from the air to create acids:
Carbon Dioxide = carbonic acid
Sulfur oxides = sulfuric acid
Nitrogen oxides = nitric acid
Damages statues, buildings, kills forests, kills fish

17. Acids and Bases in Solution
Chapter 19.2

18. Acid/Base Strength In strong acids, almost all molecules ionize.
In weak acids, fewer molecules ionize.

19. Conjugate Pairs Strength
If an acid is a strong acid, its conjugate pair base is a weak base
Why?
If HX is strong acid, it ionizes completely.
The conjugate base must be a weak base because it has a greater attraction to the H+ than HX
The reaction equilibrium lies far to the right of the equation.

20. Conjugate Pair Strength
For a weak acid, the equation equilibrium lies to the right (reactant side)
Conjugate base (Y-) has a stronger attraction for the H+ ion than the base H2O
HY + H2O H3O+ + Y-

21. Acid Ionization Constants
An “ionization constant” is the tendency of an item to make ions in solution.
Higher the constant, the higher the amount of ions.
Acid ionization constant is value of the equilibrium constant expression for a weak acid
Value Ka indicates whether reactants or products are favored at equilibrium
Weak acids have low Ka values

22. Acid Ionization Constants
Substance
Formula Ka
Acetic Acid
HC2H3O2
1.7 x 10-5
Boric Acid
H3BO3
5.9 x 10-10
Carbonic Acid
H2CO3
4.3 x 10-7
HCO3-
4.8 x 10-11
Hydrogen Sulfide
H2S
8.9 x 10-8
HS-
1.2 x 10-13
Hypochlorous Acid
HClO
3.5 x 10-8
Nitrous Acid
HNO2
4.5 x 10-4
Oxalic Acid
H2C2O4
5.6 x 10-2
HC2O4-
5.1 x 10-5
Phosphoric Acid
H3PO4
6.9 x 10-3
HSO3
6.3 x 10-8

23. Base Ionization Constant
Same Basic Principle as Acid
Measures OH- concentrations

24. pH Scale
Chapter 19.3

25. Ionization Constant for Water
The ionization constant for water is:
1.0 x 10-14
[Ka][Kb]
= [1.0x10-7] [1.0 x 10-7]
Experiments show that the product of [H+] and [OH-] always equals 1.0 x at 298°C
pH scale is a way of showing this relationship of ionization constants

26. The pH Scale pH stands for ‘per Hydrion’ Low pH is Acid
High pH is base
Water is neutral (7.0)

27. pH
There are many ways to consider acids and bases. One of these is pH.
[H+] is critical in many chemical reactions.
A quick method of denoting [H+] is via pH.
By definition pH = – log [H+], [H+] = 10-pH
The pH scale, similar to the Richter scale, describes a wide range of values
An earthquake of “6” is 10 as violent as a “5”
Thus, the pH scale condenses possible values of [H+] to a 14 point scale (fig. 2, p370)
Also, it is easier to say pH = 7 vs. [H+] = 1 x 10–7

28. pH
pH = -log [H+]
[H+] = 10-pH
pOH = -log [OH-]
pH + pOH = 14

29. Calculations with pH Q: What is the pH if [H+]= 6.3 x 10–5?
pH = – log [H+]
‘(-)’, ‘log’, ‘6.3’, ’10x’, ‘(-)’, ‘5’, ‘)”, ‘)”, ‘ENTER’)
Ans: 4.2
Q: What is the [H+] if pH = 7.4?
[H+] = 10–pH mol/L
(’10x’, ‘(-)’, ‘7.4’, “)” ‘ENTER‘)
3.98 x 10–8 M

30. Calculating pH from Strong Acid Solutions
Strong acids are 100% ionized
For monoprotic acids, concentration of the acid IS the concentration of the H+ ion
Use Acid concentration as substitute for H+ ion concentration.
Use Base concentration as substitute for OH- concentration

31. Calculating pH from Strong Acid Solutions
Example: What is the pH of a 0.1M solution of HCl?
0.1 M HCl = 1 x 10-1 M
Calculate pH = 1
Example: What is pH of solution that is 7.5 x 10-4 M Ca(OH)2?
(7..5 x 10-4) x 2 = 1.5 x 10-3M
There are 2 OH- ions per molecule
Calculate pOH = -log[OH-]
= 2.8
pH = = 11.2

32. Calculating Molarity from pH
Example: what is the molarity of an acid solution with a pH of 2.37?
[H+] = 10-pH
[H+] = = 4.27 x 10-3 M

33. Neutralization
Chapter 15.4

34. Acid-Base Reactions HCl + NaOH H2O + Na+ + Cl-
Neutralization reaction is a reaction between an acid and a base
Makes Water + Salt
Solution becomes Neutral (not acid or base)
HCl + NaOH H2O + Na+ + Cl-

35. Acid-Base Reactions Mg(OH)2 + 2 HCl → MgCl2 + 2H2O Note:
Cation from base (Mg) is combined with anion from acid (Cl)
The salt is MgCl2
The H+ and OH- always combine to form water

36. Acid-Base Titration
Acid/Base Titration is the stoichiometry of acid/base reactions.
Titration is a method for determining the concentration of a solution by using another solution of known concentration
Uses an INDICATOR to show when the acid/base reaction is complete (neutral)
Indicator is a chemical that changes color as determined by acid or base conditions
There are many indicators with different pH points.

37. Acid/Base Titration Curve

38. pH Indicators Name Acid Color pH Range of Color Change Base Color
Methyl violet
Yellow
Blue
Thymol blue
Red
Methyl orange
Bromocresol green
Methyl red
Litmus
Bromothymol blue
Phenolphthalein
Colorless
Pink
Thymolphthalein
Alizarin yellow R

39. Calculating Molarity from Titration
Write the balanced equation
Calculate the number of moles used in the ‘known’ solution
Use the mole ratio from the balanced equation to calculate moles of reactant in the ‘unknown’ solution
Calculate the molarity of the ‘unknown’ solution based on moles used and liters used.

40. Salt Hydrolysis
When you put salts in water, the resulting solution can be either acid, base, or neutral
Salts will dissolve to form ions
The anions will accept hydrogens from water
The cations will accept hydroxides from water
Which way it goes depends upon the strength of the conjugate acids/bases
If conjugate acid is strong, it will be acid
If conjugate base is strong, it will be base
If both are strong, it will be neutral

41. What are Buffers?
Buffers are solutions that resist changes in pH when limited amounts of acid or base are added.
Buffer is a weak acid and its conjugate base or a weak base and it’s conjugate acid
Buffer can accept or donate Hydrogen ions and shift its equilibrium point left or right
Buffers have limits, but the work for a while
Heavily used in human body, especially blood

42. Acid Names
1. Binary or hydrohalic acids – HF, HCl, HBr, HI, etc. “hydro____ic acid” are usually strong acids
If name ends in ‘-ide’
Acid name will be “hydro ____ic acid”
HF and H2S are weak hydrohalic acid. Although the H-F bond is very polar, the bond is so strong (due to the small F atom) that the acid does not completely ionize.

43. Acid Naming 2. Oxyacids – contain a polyatomic ion
Most common form (MCF) “ic” ending – strong acids (contain 2 oxygen per hydrogen)
If chemical name ends in “-ate”
Acid name will be “___IC Acid”
HNO3 – nitric from nitrate
H3PO4 - phosphoric from phosphate
H2SO4 - sulfuric from sulfate
HClO3 - chloric from chlorate

44. Acids with 1 less oxygen than the MCF “ous” ending- weaker acids
Chemical name ends in “-ite”
Acid name is “___OUS Acid”
HNO2 – nitrous from nitrite
H3PO3 - phosphorous from phosphite
H2SO3 - sulfurous from sulfite
HClO2 - chlorous from chlorite
c. Acids with 2 less oxygen than the MCF “hypo___ous” – very weak acids
HNO - hyponitrous
H3PO2 - hypophosphorus
HClO - hypochorous

45. Organic acids – have carboxyl group -COOH - usually weak acids
d. Acids with 1 more oxygen than the MCF “per______ic” – very strong acids
HClO4 – perchloric acid
HNO4 - pernitric acid
Organic acids – have carboxyl group -COOH - usually weak acids
Acid names are based on the base organic name or common name
HC2H3O2 - acetic acid
C7H5COOH - benzoic acid