1. Chapter 14 Acids and Bases

2. Types of Electrolytes salts = water soluble ionic compounds
all strong electrolytes
acids = form H+1 ions in water solution
bases = combine with H+1 ions in water solution
increases the OH-1 concentration
may either directly release OH-1 or pull H+1 off H2O

3. Properties of Acids Sour taste react with “active” metals
i.e. Al, Zn, Fe, but not Cu, Ag or Au
2 Al + 6 HCl ® 2 AlCl3 + 3 H2
corrosive
react with carbonates, producing CO2
marble, baking soda, chalk, limestone
CaCO3 + 2 HCl ® CaCl2 + CO2 + H2O
change color of vegetable dyes
blue litmus turns red
react with bases to form ionic salts

4. Common Acids

5. Structures of Acids
binary acids have acid hydrogens attached to a nonmetal atom
HCl, HF
Hydrofluoric acid

6. Structure of Acids
oxy acids have acid hydrogens attached to an oxygen atom
H2SO4, HNO3

7. Structure of Acids carboxylic acids have COOH group
HC2H3O2, H3C6H5O3
only the first H in the formula is acidic
the H is on the COOH

8. Properties of Bases also known as alkalis taste bitter
alkaloids = plant product that is alkaline
often poisonous
solutions feel slippery
change color of vegetable dyes
different color than acid
red litmus turns blue
react with acids to form ionic salts
neutralization

9. Common Bases

10. Structure of Bases most ionic bases contain OH ions
NaOH, Ca(OH)2
some contain CO32- ions
CaCO3 NaHCO3
molecular bases contain structures that react with H+
mostly amine groups

11. Arrhenius Theory
bases dissociate in water to produce OH- ions and cations
ionic substances dissociate in water
NaOH(aq) → Na+(aq) + OH–(aq)
acids ionize in water to produce H+ ions and anions
because molecular acids are not made of ions, they cannot dissociate
they must be pulled apart, or ionized, by the water
HCl(aq) → H+(aq) + Cl–(aq)
in formula, ionizable H written in front
HC2H3O2(aq) → H+(aq) + C2H3O2–(aq)

12. 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)

13. Problems with Arrhenius Theory
does not explain why molecular substances, like NH3, dissolve in water to form basic solutions – even though they do not contain OH– ions
does not explain acid-base reactions that do not take place in aqueous solution
H+ ions do not exist in water. Acid solutions contain H3O+ ions
H+ = a proton!
H3O+ = hydronium ions

14. Brønsted-Lowery Theory
in a Brønsted-Lowery Acid-Base reaction, an H+ is transferred
does not have to take place in aqueous solution
broader definition than Arrhenius
acid is H donor, base is H acceptor
base structure must contain an atom with an unshared pair of electrons
in the reaction, the acid molecule gives an H+ to the base molecule
H–A + :B  :A– + H–B+

15. Amphoteric Substances
amphoteric substances can act as either an acid or a base
have both transferable H and atom with lone pair
HCl(aq) is acidic because HCl transfers an H+ to H2O, forming H3O+ ions
water acts as base, accepting H+
HCl(aq) + H2O(l) → Cl–(aq) + H3O+(aq)
NH3(aq) is basic because NH3 accepts an H+ from H2O, forming OH–(aq)
water acts as acid, donating H+
NH3(aq) + H2O(l)  NH4+(aq) + OH–(aq)

16. Brønsted-Lowery Acid-Base Reactions
one of the advantages of Brønsted-Lowery theory is that it allows reactions to be reversible
H–A + :B → :A– + H–B+
the original base has an extra H+ after the reaction – so it could act as an acid in the reverse process
and the original acid has a lone pair of electrons after the reaction – so it could act as a base in the reverse process
:A– + H–B+ → H–A + :B
a double arrow, , is usually used to indicate a process that is reversible

17. Conjugate Pairs
In a Brønsted-Lowery Acid-Base reaction, the original base becomes an acid in the reverse reaction, and the original acid becomes a base in the reverse process
each reactant and the product it becomes is called a conjugate pair
the original base becomes the conjugate acid; and the original acid becomes the conjugate base

18. Brønsted-Lowery Acid-Base Reactions
H–A :B  :A– + H–B+
acid base conjugate conjugate
base acid
HCHO H2O  CHO2– + H3O+
acid base conjugate conjugate
base acid
H2O NH3  HO– + NH4+
acid base conjugate conjugate
base acid

19. Conjugate Pairs In the reaction H2O + NH3  HO– + NH4+
H2O and HO– constitute an Acid/Conjugate Base pair
NH3 and NH4+ constitute a Base/Conjugate Acid pair

20. Practice – Identify the Brønsted-Lowery Acids and Bases and their Conjugates in each Reaction
H2SO H2O  HSO4– + H3O+
HCO3– H2O  H2CO3 + HO–

21. Neutralization Reactions
H+ + OH- H2O
acid + base salt + water
double displacement reactions
salt = cation from base + anion from acid
cation and anion charges stay constant
H2SO4 + Ca(OH)2 → CaSO4 + 2 H2O
some neutralization reactions are gas evolving where H2CO3 decomposes into CO2 and H2O
H2SO4 + 2 NaHCO3 → Na2SO4 + 2 H2O + 2 CO2

22. Nonmetal Oxides are Acidic
nonmetal oxides react with water to form acids
causes acid rain
CO2 (g) + H2O(l) → H2CO3(aq)
2 SO2(g) + O2(g) + 2 H2O(l) → 2 H2SO4(aq)
4 NO2(g) + O2(g) + 2 H2O(l) → 4 HNO3(aq)

23. Acid Reactions Acids React with Metals
acids react with many metals
but not all!!
when acids react with metals, they produce a salt and hydrogen gas
3 H2SO4(aq) + 2 Al(s) → Al2(SO4)3(aq) + 3 H2(g)

24. Acid Reactions Acids React with Metal Oxides
when acids react with metal oxides, they produce a salt and water
3 H2SO4 + Al2O3 → Al2(SO4)3 + 3 H2O

25. 2 NaOH + 2 Al + 6 H2O → 2 NaAl(OH)4 + 3 H2
Base Reactions
the reaction all bases have is common is neutralization of acids
strong bases will react with Al metal to form sodium aluminate and hydrogen gas
2 NaOH + 2 Al + 6 H2O → 2 NaAl(OH)4 + 3 H2

26. Titration
using reaction stoichiometry to determine the concentration of an unknown solution
Titrant (unknown solution) added from a buret
indicators are chemicals added to help determine when a reaction is complete
the endpoint of the titration occurs when the reaction is complete

27. Titration

28. Titration The base solution is the titrant in the buret.
As the base is added to
the acid, the H+ reacts with
the OH– to form water.
But there is still excess
acid present so the color
does not change.
At the titration’s endpoint,
just enough base has been
added to neutralize all the
acid. At this point the
indicator changes color.

29. Example 14.4 Acid-Base Titration
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?

30. Strong or Weak a strong acid is a strong electrolyte
practically all the acid molecules ionize, →
a strong base is a strong electrolyte
practically all the base molecules form OH– ions, either through dissociation or reaction with water, →
a weak acid is a weak electrolyte
only a small percentage of the molecules ionize, 
a weak base is a weak electrolyte
only a small percentage of the base molecules form OH– ions, either through dissociation or reaction with water, 

31. Strong Acids The stronger the acid, the more willing it is to donate H
use water as the standard base
strong acids donate practically all their H’s
100% ionized in water
strong electrolyte
[H3O+] = [strong acid]
[ ] = molarity
HCl ® H+ + Cl-
HCl + H2O® H3O+ + Cl-

32. Strong Acids
Pure Water
HCl solution

33. Weak Acids weak acids donate a small fraction of their H’s
most of the weak acid molecules do not donate H to water
much less than 1% ionized in water
[H3O+] << [weak acid]
HF Û H+ + F-
HF + H2O Û H3O+ + F-

34. Weak Acids
Pure Water
HF solution

35. Strong Bases The stronger the base, the more willing it is to accept H
use water as the standard acid
strong bases, practically all molecules are dissociated into OH– or accept H’s
strong electrolyte
multi-OH bases completely dissociated
[HO–] = [strong base] x (# OH)
NaOH ® Na+ + OH-

36. Weak Bases
in weak bases, only a small fraction of molecules accept H’s
weak electrolyte
most of the weak base molecules do not take H from water
much less than 1% ionization in water
[HO–] << [strong base]
NH3 + H2O Û NH4+ + OH-

37. Relationship between Strengths of Acids and their Conjugate Bases
the stronger an acid is, the weaker the attraction of the ionizable H for the rest of the molecule is
the better the acid is at donating H, the worse its conjugate base will be at accepting a H
strong acid HCl + H2O → Cl– + H3O+ weak conj. base
weak acid HF + H2O  F– + H3O+ strong conj. base

38. Autoionization of Water
Water is actually an extremely weak electrolyte
therefore there must be a few ions present
about 1 out of every 10 million water molecules form ions through a process called autoionization
H2O Û H+ + OH–
H2O + H2O Û H3O+ + OH–
all aqueous solutions contain both H+ and OH–
the concentration of H+ and OH– are equal in water
[H+] = [OH–] = 25°C

39. Ion Product of Water
the product of the H+ and OH– concentrations is always the same number
the number is called the ion product of water and has the symbol Kw
[H+] x [OH–] = 1 x = Kw
as [H+] increases the [OH–] must decrease so the product stays constant
inversely proportional

40. Acidic and Basic Solutions
neutral solutions have equal [H+] and [OH–]
[H+] = [OH–] = 1 x 10-7
acidic solutions have a larger [H+] than [OH–]
[H+] > 1 x 10-7; [OH–] < 1 x 10-7
basic solutions have a larger [OH–] than [H+]
[H+] < 1 x 10-7; [OH–] > 1 x 10-7

41. Ba(OH)2 = Ba2+ + 2 OH– therefore
Example - Determine the [H+1] for a M Ba(OH)2 and determine whether the solution is acidic, basic or neutral
Ba(OH)2 = Ba OH– therefore
[OH–] = 2 x = = 4.0 x 10-4 M
[H+] = 2.5 x M

42. Practice - Determine the [H+1] concentration and whether the solution is acidic, basic or neutral for the following
[OH–] = M
[OH–] = 3.50 x 10-8 M
Ca(OH)2 = 0.20 M

43. pH the acidity/basicity of a solution is often expressed as pH
pH = -log[H+], [H+] = 10-pH
exponent on 10 with a positive sign
pHwater = -log[10-7] = 7
need to know the [H+] concentration to find pH
pH < 7 is acidic; pH > 7 is basic, pH = 7 is neutral

44. pH
the lower the pH, the more acidic the solution; the higher the pH, the more basic the solution
1 pH unit corresponds to a factor of 10 difference in acidity
normal range 0 to 14
pH 0 is [H+] = 1 M, pH 14 is [OH–] = 1 M
pH can be negative (very acidic) or larger than 14 (very alkaline)

45. pH of Common Substances
1.0 M HCl
0.0
0.1 M HCl
1.0
stomach acid
1.0 to 3.0
lemons
2.2 to 2.4
soft drinks
2.0 to 4.0
plums
2.8 to 3.0
apples
2.9 to 3.3
cherries
3.2 to 4.0
unpolluted rainwater
5.6
human blood
7.3 to 7.4
egg whites
7.6 to 8.0
milk of magnesia (sat’d Mg(OH)2)
10.5
household ammonia
10.5 to 11.5
1.0 M NaOH 14

46. Example - Calculate the pH of a 0
Example - Calculate the pH of a M Ba(OH)2 solution & determine if is acidic, basic or neutral
Ba(OH)2 = Ba OH- therefore
[OH-] = 2 x = = 2.0 x 10-3 M
[H+] =
1 x 10-14
2.0 x 10-3
= 5.0 x 10-12M
pH = -log [H+] = -log (5.0 x 10-12)
pH = 11.3
pH > 7 therefore basic

47. Practice - Calculate the pH of the following strong acid or base solutions
M HCl
M Ca(OH)2
0.25 M HNO3

48. Sample - Calculate the concentration of [H+] for a solution with pH 3
means < [H+1] < 0.001
[H+] = 2 x 10-4 M = M

49. Practice - Determine the [H+] for each of the following
pH = 2.7
pH = 12
pH = 0.60

50. Buffers
buffers are solutions that resist changing pH when small amounts of acid or base are added
they resist changing pH by neutralizing added acid or base
buffers are made by mixing together a weak acid and its conjugate base
or weak base and it conjugate acid

51. How Buffers Work
the weak acid present in the buffer mixture can neutralize added base
the conjugate base present in the buffer mixture can neutralize added acid
the net result is little to no change in the solution pH

52. What is Acid Rain? natural rain water has a pH of 5.6
naturally slightly acidic due mainly to CO2
rain water with a pH lower than 5.6 is called acid rain
acid rain is linked to damage in ecosystems and structures

53. What Causes Acid Rain?
many natural and pollutant gases dissolved in the air are nonmetal oxides
CO2, SO2, NO2
nonmetal oxides are acidic
CO2 + H2O  H2CO3
2 SO2 + O2 + 2 H2O  2 H2SO4
processes that produce nonmetal oxide gases as waste increase the acidity of the rain
natural – volcanoes and some bacterial action
man-made – combustion of fuel
weather patterns may cause rain to be acidic in regions other than where the nonmetal oxide is produced

54. Damage from Acid Rain
acids react with metals, and materials that contain carbonates
acid rain damages bridges, cars and other metallic structures
acid rain damages buildings and other structures made of limestone or cement

55. Damage from Acid Rain
circa 1935
circa 1995