1. Acids, Bases, Salts, Solubility, And stuff like that!
Chapter 6
Acids, Bases, Salts, Solubility,
And stuff like that!

2. Definitions Solubility:
those compounds with low solubility are said to be insoluble,
those compounds with higher solubility are said to be soluble

3. More Definitions saturated solution: unsaturated solution:
supersaturated solution:

4. Soluble or Insoluble
Explain why some substances are soluble and other substances are not soluble by giving one example of each. Used balanced equations in you discussion.
You may use the solubility rules – Use your intelligence and understanding if the internet to find them! – These are observation based – no explanation needed at this time.

5. Solvation
What happens when substances dissolve? What forces are involved? Use water as a solvent for specific examples.
Ionic?
Covalent?

6. Water as a Solvent
How water dissolves ionic compounds
water is a
ions

7. Water as a Solvent How water dissolves molecular compounds
nonpolar covalent molecules
polar covalent molecules dissolve because
Each individual molecule is

8. Electrolytes Video Link-electrolytes and non-electrolytes
Video Link – Weak and strong electrolytes

9. Electrolytes cations migrate to the negative electrode (the cathode)
anions migrate to the positive electrode (the anode)
the movement of ions constitutes an electric current
electrolyte:
nonelectrolyte
strong electrolyte:
weak electrolyte:

10. Arrhenius Acids and Bases
In 1884, Svante Arrhenius proposed these definitions
acid: a substance that produces H3O+ ions aqueous solution
base: a substance that produces OH- ions in aqueous solution

11. Arrhenius Acids and Bases
when HCl, for example, dissolves in water, its reacts with water to give hydronium ion and chloride ion
we use curved arrows to show the change in position of electron pairs during this reaction

12. Arrhenius Acids and Bases
With bases, the situation is slightly different
many bases are metal hydroxides such as KOH, NaOH, Mg(OH)2, and Ca(OH)2
these compounds are ionic solids and when they dissolve in water, their ions merely separate
other bases are not hydroxides; these bases produce OH- by reacting with water molecules

13. Arrhenius Acids and Bases
we use curved arrows to show the transfer of a proton from water to ammonia

14. Acid and Base Strength
Strong acid: one that reacts completely or almost completely with water to form H3O+ ions
Strong base: one that reacts completely or almost completely with water to form OH- ions
here are the six most common strong acids and the four most common strong bases

15. Acid and Base Strength
Weak acid: a substance that dissociates only partially in water to produce H3O+ ions
acetic acid, for example, is a weak acid; in water, only 4 out every 1000 molecules are converted to acetate ions
Weak base: a substance that dissociates only partially in water to produce OH- ions
ammonia, for example, is a weak base

16. Brønsted-Lowry Acids & Bases
Acid: a proton donor
Base: a proton acceptor
Acid-base reaction: a proton transfer reaction
Conjugate acid-base pair: any pair of molecules or ions that can be interconverted by transfer of a proton

17. Brønsted-Lowry Acids & Bases
Brønsted-Lowry definitions do not require water as a reactant

18. Brønsted-Lowry Acids & Bases
we can use curved arrows to show the transfer of a proton from acetic acid to ammonia

20. Brønsted-Lowry Acids & Bases
Note the following about the conjugate acid-base pairs in the table
1. an acid can be positively charged, neutral, or negatively charged; examples of each type are H3O+, H2CO3, and H2PO4-
2. a base can be negatively charged or neutral; examples are OH-, Cl-, and NH3
3. acids are classified a monoprotic, diprotic, or triprotic depending on the number of protons each may give up; examples are HCl, H2CO3, and H3PO4

21. Brønsted-Lowry Acids & Bases
carbonic acid, for example can give up one proton to become bicarbonate ion, and then the second proton to become carbonate ion
4. several molecules and ions appear in both the acid and conjugate base columns; that is, each can function as either an acid or a base

22. Brønsted-Lowry Acids & Bases
the HCO3- ion, for example, can give up a proton to become CO32-, or it can accept a proton to become H2CO3
a substance that can act as either an acid or a base is said to be amphiprotic
the most important amphiprotic substance in Table 8.2 is H2O; it can accept a proton to become H3O+, or lose a proton to become OH-
5. a substance cannot be a Brønsted-Lowry acid unless it contains a hydrogen atom, but not all hydrogen atoms in most compounds can be given up
acetic acid, for example, gives up only one proton

23. Brønsted-Lowry Acids & Bases
6. there is an inverse relationship between the strength of an acid and the strength of its conjugate base
the stronger the acid, the weaker its conjugate base
HI, for example, is the strongest acid in Table 8.2, and its conjugate base, I-, is the weakest base in the table
CH3COOH (acetic acid) is a stronger acid that H2CO3 (carbonic acid); conversely, CH3COO- (acetate ion) is a weaker base that HCO3- (bicarbonate ion)

24. 6.4. DISSOCIATION OF ACIDS AND BASES IN WATER
Table 6.1. Dissociation of Acids
% dissociated
Formula Name Common uses in 1 M solution Strength
H2SO4 Sulfuric Industrial chemical 100 Strong
HNO3 Nitric Industrial chemical 100 Strong
H3PO4 Phosphoric Fertilizer, food 8 Moderately
additive weak
H3C6 H5O7 Citric Fruit drinks 3 Weak
CH3CO2H Acetic Foods, industry 0.4 Weak
HClO Hypochlorous Disinfectant Weak
HCN Hydrocyanic Very poisonous Very weak
industrial chemical
electroplating waste
H3BO3 Boric acid Antiseptic, ceramics Very weak

25. Acid-Base Equilibria
we know that HCl is a strong acid, which means that the position of this equilibrium lies very far to the right
in contrast, acetic acid is a weak acid, and the position of its equilibrium lies very far to the left
but what if the base is not water? How can we determine which are the major species present?

26. Acid-Base Equilibria
To predict the position of an acid-base equilibrium such as this, we do the following
identify the two acids in the equilibrium; one on the left and one on the right
using the information in Table 10.1, determine which is the stronger acid and which is the weaker acid
also determine which is the stronger base and which is the weaker base; remember that the stronger acid gives the weaker conjugate base, and the weaker acid gives the stronger conjugate base
the stronger acid reacts with the stronger base to give the weaker acid and weaker base; equilibrium lies on the side of the weaker acid and weaker base

27. Acid-Base Equilibria
identify the two acids and bases, and their relative strengths
the position of this equilibrium lies to the right

28. Acid-Base Equilibria
Example: predict the position of equilibrium in this acid-base reaction

29. Acid-Base Equilibria
Example: predict the position of equilibrium in this acid-base reaction
Solution: the position of this equilibrium lies to the right

30. Acid Ionization Constants
when a weak acid, HA, dissolves in water
the equilibrium constant, Keq, for this ionization is
because water is the solvent and its concentration changes very little when we add HA to it, we treat [H2O] as a constant equal to 1000 g/L or 55.5 mol/L
we combine the two constants to give a new constant, which we call an acid ionization constant, Ka

31. Acid Ionization Constants
Ka for acetic acid, for example is 1.8 x 10-5
because the acid ionization constants for weak acids are numbers with negative exponents, we commonly express acid strengths as pKa where
the value of pKa for acetic acid is 4.75
values of Ka and pKa for some weak acids are given in Table 10.2
as you study the entries in this table, note the inverse relationship between values of Ka and pKa
the weaker the acid, the smaller its Ka, but the larger its pKa

33. Properties of Acids & Bases
Neutralization
acids and bases react with each other in a process called neutralization.
Reaction of acids with metals
strong acids react with certain metals (called active metals) to produce a salt and hydrogen gas, H2

34. Properties of Acids & Bases
Reaction with metal hydroxides
reaction of an acid with a metal hydroxide gives a salt plus water
the reaction is more accurately written as
omitting spectator ions gives this net ionic equation

35. Properties of Acids & Bases
Reaction with metal oxides
strong acids react with metal oxides to give water plus a salt

36. Properties of Acids & Bases
Reaction with carbonates and bicarbonates
strong acids react with carbonates to give carbonic acid, which rapidly decomposes to CO2 and H2O
strong acids react similarly with bicarbonates

37. Properties of Acids & Bases
Reaction with ammonia and amines
any acid stronger than NH4+ is strong enough to react with NH3 to give a salt

38. Self-Ionization of Water
pure water contains a very small number of H3O+ ions and OH- ions formed by proton transfer from one water molecule to another
the equilibrium expression for this reaction is
we can treat [H2O] as a constant = 55.5 mol/L

39. Self-Ionization of Water
combining these constants gives a new constant called the ion product of water, Kw
in pure water, the value of Kw is 1.0 x 10-14
this means that in pure water

40. Self-Ionization of Water
the product of [H3O+] and [OH-] in any aqueous solution is equal to 1.0 x for solutions as well.
for example, if we add mole of HCl to 1 liter of pure water, it reacts completely with water to give mole of H3O+
in this solution, [H3O+] is or 1.0 x 10-2
this means that the concentration of hydroxide ion is

41. pH and pOH we commonly express these concentrations as pH, where
pH = -log [H3O+]
we can now state the definitions of acidic and basic solutions in terms of pH
acidic solution: one whose pH is less than 7.0
basic solution: one whose pH is greater than 7.0
neutral solution: one whose pH is equal to 7.0

42. pH and pOH
just as pH is a convenient way to designate the concentration of H3O+, pOH is a convenient way to designate the concentration of OH-
pOH = -log[OH-]
the ion product of water, Kw, is 1.0 x 10-14
taking the logarithm of this equation gives
pH + pOH = 14
thus, if we know the pH of an aqueous solution, we can easily calculate its pOH

43. pH and pOH
pH of some common materials

44. pH of Salt Solutions
When some salts dissolve in pure water, there is no change in pH from that of pure water
Many salts, however, are acidic or basic and cause a change the pH when they dissolve
We are concerned in this section with basic salts and acidic salts

45. pH of Salt Solutions Basic salt: raises the pH
as an example of a basic salt is sodium acetate
when this salt dissolves in water, it ionizes; Na+ ions do not react with water, but CH3COO- ions do
the position of equilibrium lies to the left
nevertheless, there are enough OH- ions present in 0.10 M sodium acetate to raise the pH to 8.88

46. pH of Salt Solutions Acidic salt: lowers the pH
an example of an acidic salt is ammonium chloride
chloride ion does not react with water, but the ammonium ion does
although the position of this equilibrium lies to the left, there are enough H3O+ ions present to make the solution acidic

47. Acid-Base Titrations
Titration: an analytical procedure in which a solute in a solution of known concentration reacts with a known stoichiometry with a substance whose concentration is to be determined

48. Acid-Base Titrations
An acid-base titration must meet these requirement
1. we must know the equation for the reaction so that we can determine the stoichiometric ratio of reactants to use in our calculations
2. the reaction must be rapid and complete
3. there must be a clear-cut change in a measurable property at the end point (when the reagents have combined exactly)
4. we must have precise measurements of the amount of each reactant

49. Acid-Base Titrations
As an example, let us use M H2SO4 to determine the concentration of a NaOH solution
requirement 1: we know the balanced equation
requirement 2: the reaction between H3O+ and OH- is rapid and complete
requirement 3: we can use either an acid-base indicator or a pH meter to observe the sudden change in pH that occurs at the end point of the titration
requirement 4: we use volumetric glassware

50. Acid-Base Titrations
experimental measurements
doing the calculations

51. pH Buffers
pH buffer: a solution that resists change in pH when limited amounts of acid or base are added to it
a pH buffer as an acid or base “shock absorber”
a pH buffer is common called simply a buffer
the most common buffers consist of approximately equal molar amounts of a weak acid and a salt of the conjugate base of the weak acid
for example, if we dissolve 1.0 mole of acetic acid and 1.0 mole of its conjugate base (in the form of sodium acetate) in water, we have an acetate buffer

52. pH Buffers How an acetate buffer resists changes in pH
if we add a strong acid, such as HCl, added H3O+ ions react with acetate ions and are removed from solution
if we add a strong base, such as NaOH, added OH- ions react with acetic acid and are removed from solution

53. pH Buffers The effect of a buffer can be quite dramatic
consider a phosphate buffer prepared by dissolving 0.10 mole of NaH2PO4 (a weak acid) and 0.10 mole of Na2HPO4 (the salt of its conjugate base) in enough water to make 1 liter of solution

54. pH Buffers
Buffer pH
if we mix equal molar amounts of a weak acid and a salt of its conjugate base, the pH of the solution will be equal to the pKa of the weak acid
if we want a buffer of pH 9.14, for example, we can mix equal molar amounts of boric acid (H3BO3), pKa 9.14, and sodium dihydrogen borate (NaH2BO3), the salt of its conjugate base

55. pH Buffers
Buffer capacity depends both its pH and its concentration

56. Blood Buffers The average pH of human blood is 7.4
any change larger than 0.10 pH unit in either direction can cause illness
To maintain this pH, the body uses three buffer systems
carbonate buffer: H2CO3 and its conjugate base, HCO3-
phosphate buffer: H2PO4- and its conjugate base, HPO42-
proteins: discussed in Chapter 21

57. Henderson-Hasselbalch Eg.
Henderson-Hasselbalch equation: a mathematical relationship between
pH,
pKa of the weak acid, HA
concentrations HA, and its conjugate base, A-
It is derived in the following way
taking the logarithm of this equation gives

58. Henderson-Hasselbalch Eg.
multiplying through by -1 gives
-log Ka is by definition pKa, and -log [H3O+] is by definition pH; making these substitutions gives
rearranging terms gives

59. Henderson-Hasselbalch Eg.
Example: what is the pH of a phosphate buffer solution containing 1.0 mole of NaH2PO4 and 0.50 mole of Na2HPO4 dissolved in enough water to make 1.0 liter of solution

60. Henderson-Hasselbalch Eg.
Example: what is the pH of a phosphate buffer solution containing 1.0 mole of NaH2PO4 and 0.50 mole of Na2HPO4 in enough water to make one liter of solution
Solution
the equilibrium we are dealing with and its pKa are
substituting these values in the H-H equation gives

61. PREPARATION OF ACIDS Combination of H with nonmetal: H2 + Cl2  2HCl
Nonmetal with water: Cl2 + H2O  HCl + HClO
Nonmetal oxide plus water: SO3 + H2O  H2SO4
Evolution of volatile acid: 2NaCl(s) + H2SO4(l)  2HCl(g) + Na2SO4(s)
• HCl gas collected in water gives hydrochloric acid
Organic acids, such as acetic acid, have the carboxylic acid group:

62. PREPARATION OF BASES Active metal plus water
• 2K + H2O  2K+ + 2OH- + H2(g)
Metal oxide plus water
• CaO + H2O  Ca(OH)2
Substances that generate OH- in water
NH3 + H2O  NH4+ + OH-
Salt anions that react with water to produce OH-
• From Na2CO3: 2Na+ + CO32- + H2O  2Na+ + HCO3-+ OH-
• This reaction is a hydrolysis reaction
Organic bases, particularly amines
• (CH3)3N + H2O  (CH3)3NH+ + OH-

63. PREPARATION OF SALTS Reaction of acid with base
• 2NaOH + H2SO4  2H2O + Na2SO4 (sodium sulfate)
Reaction of metal and nonmetal
• Ca + F2  CaF2 (calcium fluoride)
Metal reacting with acid
• Mg + H2SO4  H2 + MgSO4 (magnesium sulfate)
Active metal reacting with base
• 2Al + 6NaOH  3H2(g) + Na3AlO3 (sodium aluminate)
Addition of a base to a salt to form another salt and an insoluble base
• 2KOH + MgSO4  Mg(OH)2(s) + K2SO4(aq)
Evolution of a volatile acid leaving a salt
• 2NaCl(s) + H2SO4(l)  2HCl(g) + Na2SO4(s)
Displacement of a metal from a salt, such as in cementation
• Fe(s) + CdSO4(aq)  Cd(s) + FeSO4(aq)
Specialized processes, such as the Solvay synthesis of NaHCO3
NaCl + NH3 + CO2 + H2O  NaHCO3 + NH4Cl

64. 6.10. ACID SALTS AND BASIC SALTS
Acid salts are salts that contain H and can act as acids
• NaHSO4 + NaOH  Na2SO4 + H2O
• Sodium bicarbonate: NaHCO3
• Sodium dihydrogen phosphate, NaH2PO4, used to prepare buffers
• Disodium hydrogen phosphate, Na2HPO4, buffers
• Potassium hydrogen tartrate, KH4C4H4O6, acid in baking powder
Basic salts contain OH and can react with H+ ion
Example: Calcium hydroxyapatite source of phosphorus
• Ca5OH(PO4)3
Many rock-forming minerals are basic salts

65. 6.11. WATER OF HYDRATION
Water molecules bound to other compounds, typically salts
• Example: Sodium carbonate decahydrate, Na2CO3•10H2O

66. 6.12. NAMES OF ACIDS, BASES, AND SALTS
H and a nonmetal: Hydro-ic acid
• Hydrochloric acid, HCl
Oxygen-containing acids
• H2SO3, sulfurous acid
Table 6.7. Names of Oxyacids of Chlorine
Formula Name Anion name
HClO4 perchloric acid perchlorate
HClO3 chloric acid chlorate
HClO2 chlorous acid chlorite
HClO hypochlorous acid hypochlorite
Bases
For ionic bases containing OH, name of cation followed by hydroxide
• NaOH, sodium hydroxide Ca(OH)2, calcium hydroxide

67. Names of Salts Formulas of Salts
Name of cation followed by name of anion
See Table 6.8 for some important ions and their names
Examples:
• Na2SO4, sodium sulfate
• KH2PO4, potassium dihydrogen sulfate
• Ca(ClO)2, calcium hypochlorite
Formulas of Salts
Sum of charge on cations times their subscripts plus sum of charge on anions times their subscripts must equal zero
• Example: Iron(III) sulfate
• Formula before adding subscripts: Fe(SO4)
• Cation charge: • Anion charge: -2
• 2 Fe3+ cations gives a total cation charge of 2 x 3 = 6
• 3 SO42- gives an anion charge of 3 x (-2) = -6
• Therefore, the formula is Fe2(SO4)3