acids-and-bases-lecture.ppt

• There are three classes of strong electrolytes.
1 Strong Water Soluble Acids
2 Strong Water Soluble Bases
3 Most Water Soluble Salts
Acids and bases
• Weak acids and bases ionize or dissociate partially,
much less than 100%, and is often less than 10%.

Monovalent Divalent Trivalent
Hydride H- Oxide O2- Nitride N3-
Fluoride Fl- Peroxide O2
2-
Chloride Cl- Sulfide S2-
Bromide Br- Selenide Se2-
Iodide I- Oxalate C2O4
2-
Hydroxide OH- Chromate CrO4
2-
Permangante MnO4
- Dichromate Cr2O7
2-
Cyanide CN- Tungstate WO4
2-
Thiocynate SCN- Molybdate MoO4
2-
Acetate CH3COO- Tetrathionate S4O6
2-
Nitrate NO3
- Thiosulfate S2O3
2-
Bisulfite HSO3
- Sulfite SO3
2-
Bisulfate HSO4
- Sulfate SO4
2-
Bicarbonate HCO3
- Carbonate CO3
2-
Dihydrogen phosphate H2PO4
- Hydrogen phosphate HPO4
2- Phosphate PO4
3-
Nitrite NO2
-
Amide NH2
-
Hypochlorite ClO-
Chlorite ClO2
-
Chlorate ClO3
-
Perchlorate ClO4
-
Table of Common Ions
Common Negative Ions (Anions)

Monovalent Divalent Trivalent
Hydronium (aqueous) H3O+ Magnesium Mg2+ Aluminium Al3+
Hydrogen (proton) H+ Calcium Ca2+ Antimony III Sb3+
Lithium Li+ Strontium Sr2+ Bismuth III Bi3+
Sodium Na+ Beryllium Be2+
Potassium K+ Manganese II Mn2+
Rubidium Rb+ Barium Ba2+
Cesium Cs+ Zinc Zn2+
Francium Fr+ Cadmium Cd2+
Silver Ag+ Nickel II Ni2+
Ammonium NH4
+ Palladium II Pd2+
Thalium Tl+ Platinum II Pt2+
Copper I Cu+ Copper II Cu2+
Mercury II Hg2+
Mercury I Hg2
2+
Iron II Fe2+ Iron III Fe3+
Cobalt II Co2+ Cobalt III Co3+
Chromium II Cr2+ Chromium III Cr3+
Lead II Pb2+
Tin II Sn2+
Table of Common Ions
Common Positive Ions (Cations)

Water Solubility of Ionic Compounds
If one ion from the
“Soluble Compound” list
is present in a compound,
the compound is water
soluble.

• Most salts of strong or weak electrolytes
can dissolve in water to produce a neutral,
basic, or acidic solution, depending on
whether it contains the conjugate base of a
weak acid as the anion (A–) or the
conjugate acid of a weak base as the cation
(BH+), or possibly both.
• Salts that contain small, highly charged
metal ions produce acidic solutions in
H2O.
Acids and bases

The reaction of a salt with water to
produce an acidic or basic solution
is called a hydrolysis reaction,
which is just an acid-base reaction
in which the acid is a cation or the
base is an anion.

1. Arrhenius definition:
• An acid is a substance that dissociates in
water to produce H+ ions (protons), and
• A base is a substance that dissociates in
water to produce OH– ions (hydroxide);
• An acid-base reaction involves the
reaction of a proton with the hydroxide
ion to form water.
Acids and bases can be defined in
different ways:

Strong and Weak Acids
• Strong acids ionize completely in aqueous
solution:
• HCl(aq) + H2O  H3O+
(aq) + Cl-
(aq);
• H2SO4(aq) + H2O  H3O+
(aq) + HSO4
-
(aq);
• Weak acids ionize only partially in aqueous
solution:
• HF(aq) + H2O ⇌ H3O+
(aq) + F-
(aq);
• HOCl(aq) + H2O ⇌ H3O+
(aq) + ClO-
(aq).

Complete Ionization of Hydrochloric Acid

Dissociations of Strong and Weak Acids

Strong and Weak Bases
• Strong bases ionize completely in aqueous
solution:
NaOH(aq)  Na+
(aq) + OH-
(aq);
Ba(OH)2(aq)  Ba2+
(aq) + 2 OH-
(aq);
• Weak bases ionize only partially in aqueous
solution:
NH3(aq) + H2O  NH4
+
(aq) + OH-
(aq);
PO4
3-
(aq) + H2O  HPO4
2-
(aq) + OH-
(aq)

– Three limitations
1. Definition applied only to
substances in aqueous solutions.
2. Definition restricted to substances
that produce H+ and OH– ions
3. Definition does not explain why
some compounds containing
hydrogen such as CH4 dissolve in
water and do not give acidic
solutions

Neutralization happens because
hydrogen ions and hydroxide ions
react to produce water.

In the sodium hydroxide case, hydrogen ions
from the acid are reacting with hydroxide ions
from the sodium hydroxide - in line with the
Arrhenius theory.
Ammonia reacts with water like this:
This is a reversible reaction, and in a typical
dilute ammonia solution, about 99% of the
ammonia remains as ammonia molecules.
Nevertheless, there are hydroxide ions there and
those react with hydrogen ions in just the same
way as hydroxide ions from sodium hydroxide.

 An acid is a proton (hydrogen ion) donor.
 A base is a proton (hydrogen ion)
acceptor.
2. Brønsted-Lowry definition
The relationship between the Bronsted-
Lowry theory and the Arrhenius theory
Hydroxide ions are still bases because they
accept hydrogen ions from acids and form water.
An acid produces hydrogen ions in solution
because it reacts with the water molecules by
giving a proton to them.

When hydrogen chloride gas dissolves in water
to produce hydrochloric acid, the hydrogen
chloride molecule gives a proton (a hydrogen ion)
to a water molecule. A co-ordinate (dative
covalent) bond is formed between one of the lone
pairs on the oxygen and the hydrogen from the
HCl. Hydroxonium ions, H3O+, are produced.

When an acid in solution reacts with a base,
what is actually functioning as the acid is the
hydroxonium ion. For example, a proton is
transferred from a hydroxonium ion to a
hydroxide ion to make water.
Showing the electrons, but leaving out the
inner ones:

Brønsted-Lowry: Conjugate Acids & Bases
• Consider the following equilibrium:
HA + B ⇌ BH+
+ A-
;
Acid1 Base2 Conjugate Conjugate
acid2 base1
• A-
is the conjugate base to acid HA;
HA & A-
are conjugate acid-base pair;
• BH+ is the conjugate acid to base B;
BH+ & B are also conjugate acid-base pair.

Brønsted-Lowry’s Acid-Base Reaction

Brønsted’s Conjugate Acid-Base Pairs

Examples of Conjugate Acid-Base
Pairs
HNO3 – NO3
-
H2SO4 – HSO4
-
H3O+
– H2O
HF – F-
H3PO4 – H2PO4
-
CH3COOH – CH3COO-
H2PO4
-
– HPO4
2-
NH4
+
– NH3
H2O – OH-

Exercise #1: Conjugate Acids & Bases
Write the formulas of the conjugate bases
for the following acids:
(a) H2CO3
(b) HSO4
-
(c) Al(H2O)6
3+
(d) Cr(OH)3(H2O)3
Answer: (a) HCO3-; (b) SO4
2-
;
(c) Al(H2O)5(OH)2+
;
(d) Cr(OH)4(H2O)2
-

Exercise #2: Conjugate Acids and Bases
Write the formulas of the conjugate acids for
the following bases:
(a) NH3 (b) CO3
2-
(c) C5H5N (d) Al(OH)3(H2O)3
(Answer: (a) NH4
+
;
(b) HCO3
-
;
(c) C5H5NH+
;
(d) Al(OH)2(H2O)4
+
)

The Lewis Theory of acids and bases
This theory extends well beyond the things
you normally think of as acids and bases.
The theory
 An acid is an electron pair acceptor.
 A base is an electron pair donor.

Ammonia reacts with BF3 by using its lone
pair to form a co-ordinate bond with the
empty orbital on the boron.
The relationship between the Lewis theory
and the Bronsted-Lowry theory

As far as the ammonia is concerned, it
is behaving exactly the same as when
it reacts with a hydrogen ion - it is
using its lone pair to form a co-ordinate
bond.

Acid Strength and Ionization
Constants
• For the ionization or dissociation equilibrium
of an acid in aqueous solution, such as:
HA(aq) + H2O  H3O+
(aq) + A-
(aq);
The ionization or dissociation constant is
expressed as follows:
The values of Ka indicate the relative
strength of the acids. Strong acids have
very large Ka, while weak acids have small
Ka’s (Ka << 1)
[HA]
]
][A
O
[H -
3


a
K

Relative Strength of Acids and Their
Conjugate Bases
Acids Conjugate Bases
Very Strong Very Weak
Strong Weak
Weak Strong
Very Weak Very Strong
______________________________
• Strong acids lose protons very readily  weak
conjugate bases;
• Weak acids do not lose protons very readily 
strong conjugate bases.

pH, a Concentration Scale
pH: a way to express acidity -- the concentration of H+ in solution.
Low pH: high [H+] High pH: low [H+]
Acidic solution pH < 7
Neutral pH = 7
Basic solution pH > 7
pH = log (1/ [H+]) = - log [H+]
Acid Formula pH at half
equivalence
point
Acetic CH3COOH 4.7
Nitrous HNO2 3.3
Hydrofluoric HF 3.1
Hypochlorous HClO 7.4
Hydrocyanic HCN 9

• A convenient way to express the acidity and
basicity of a solution is the pH and pOH scales.
• The pH of an aqueous solution is defined as:
Acid Formula pH
Acetic CH3COOH 4.7
Nitrous HNO2 3.3
Hydrofluoric HF 3.1
Hypochlorous HClO 7.4
Hydrocyanic HCN 9
pH, a Concentration Scale

Acid-Base Properties of Water
• Auto-ionization of water:
2H2O  H3O+
(aq) + OH-
(aq)
Kw = [H3O+
][OH-
] = 1.0 x 10-14
at 25o
C
• Water ionizes to produce both H3O+
and OH-
,
thus it has both acid and base properties. Kw is
called water ionization constant.
• Pure water at 25o
C: [H3O+
] = [OH-
] = 1.0 x 10-7
M

Expressing Acidity in pH Scale
• pH = -log[H+
]
• (note: [H+
] = [H3O+
])
• pOH = -log[OH-
]
• pKw = -log(Kw);
• pKa = -log(Ka);
• pKb = -log(Kb)

For water, Kw = [H3O+
][OH-
] = 1.0 x 10-14
-log(Kw) = -log [H3O+
] + (-log[OH-
])
pKw = pH + pOH = 14.00
At 25o
C, pOH = 14 – pH

Acidity and pH Range
• Acidic solutions:
[H3O+
] > 1.0 x 10-7
M;
pH < 7;
• Basic solutions:
[OH-
] > 1.0 x 10-7
M or [H3O+
] < 1.0 x 10-7
M
pH > 7;
• Neutral solutions:
[H3O+
] = [OH-
] = 1.0 x 10-7
M; pH = 7.00

Relationship between acidity of solution and pH

[H3O+
] and pH of Strong Acids
• Strong acids like HCl and HClO4 ionize
completely in aqueous solution:
HCl(aq) + H2O  H3O+
(aq) + Cl-
(aq);
HClO4(aq) + H2O  H3O+
(aq) + ClO4
-
(aq);
• In solutions of strong monoprotic acids HA,
such as HCl and HClO4,
[H3O+
] = [HA]0
• For example, in 0.10 M HCl, [H3O+
] = 0.10 M,
and pH = -log(0.10) = 1.00

[OH-
] and pH of Strong Bases
• Like strong acids, strong bases also ionize
completely in aqueous solution.
• Examples: NaOH(aq)  Na+
(aq) + OH-
(aq);
;
• In a base solution such as 0.050 M NaOH,
[OH-
] = [NaOH]0 = 0.050 M;
pOH = -log(0.050 M) = 1.30; pH = 14.00 -
1.30 = 12.70

In a base solution such as 0.050 M
Ba(OH)2,
[OH-
] = 2 x [Ba(OH)2]0 = 0.10 M;
pOH = -log(0.10) = 1.00;
pH = 14.00 - 1.00 = 13.00
Ba(OH)2(aq)  Ba2+
(aq) + 2 OH-
(aq)

Acid-Base Equilibrium Constants:
Ka, Kb, pKa, and pKb
• The magnitude of the equilibrium constant for an ionization
reaction can determine the relative strengths of acids and bases
• The general equation for the ionization of a weak acid in water,
where HA is the parent acid and A– is its conjugate base, is
HA(aq) + H2O(l) ⇋ H3O+(aq) + A–(aq)
• The equilibrium constant for this dissociation is
K = [H3O+] [A–]
[H2O] [HA]
• The concentration of water is constant for all reactions in
aqueous solution, so [H2O] can be incorporated into a new
quantity, the acid ionization constant (Ka):
Ka = K[H2O] = [H3O+] [A–]
[HA]

• Strong acids and bases ionize essentially completely in water;
the percent ionization is always approximately 100%,
regardless of the concentration
• The percent ionization in solutions of weak acids and bases is
small and depends on the analytical concentration of the weak
acid or base; percent ionization of a
weak acid or a weak base actually
increases as its analytical
concentration decreases and
percent ionization increases as the
magnitude of the ionization
constants Ka and Kb increases
Ionization Constants for Weak
Monoprotic Acids and Bases

• Let’s look at the dissolution of acetic acid, a weak acid, in water as an
example.
• The equation for the ionization of acetic acid is:
• The equilibrium constant for this ionization is expressed as:

• The water concentration in dilute aqueous solutions is very high.
• 1 L of water contains 55.5 moles of water.
• Thus in dilute aqueous solutions:
• The water concentration is many orders of magnitude greater than the ion
concentrations.
• Thus the water concentration is essentially that of pure water.
– Recall that the activity of pure water is 1.

• We can define a new equilibrium constant for weak acid equilibria that uses
the previous definition.
– This equilibrium constant is called the acid ionization constant.
– The symbol for the ionization constant is Ka.

• The ionization constant values for several acids are given below.
– Which acid is the strongest?
– Are all of these acids weak acids?
– What is the relationship between Ka and strength?
– What is the relationship between pKa and strength?
– What is the relationship between pH and strength?
Acid Formula Ka value pKa value
-log Ka
pH of 1M
analytical [HA]
Acetic CH3COOH 1.8 x 10-5
4.7 2.4
Nitrous HNO2 4.5 x 10-4
3.3 1.7
Hydrofluoric HF 7.2 x 10-4
3.1 1.6
Hypochlorous HClO 3.5 x 10-8
7.5 3.7
Hydrocyanic HCN 4.0 x 10-10
9.4 4.7

• Complete the algebra and solve for the concentrations of the species.
• Note that the properly applied simplifying assumption gives the same result as
solving the quadratic equation does.

Acid-Base Properties of
Solutions of Salts
• A salt can dissolve in water to produce a neutral, basic, or acidic solution,
depending on whether it contains the conjugate base of a weak acid as the
anion (A–) or the conjugate acid of a weak base as the cation (BH+), or
both.
• Salts that contain small, highly charged metal ions produce acidic solutions
in water.
• The most important parameter for predicting the effect of a metal ion on the
acidity of coordinated water molecules is the charge-to-radius ratio of the
metal ion.
• The reaction of a salt with water to produce an acidic or basic solution is
called a hydrolysis reaction, which is just an acid-base reaction in which
the acid is a cation or the base is an anion.

• Acid and Base strengths can be compared using Ka and Kb values. The larger the Ka or Kb
value the more product favored the dissociation.
• An acid-base equilibrium always favors the side with the weaker acid and base.
• In an acid-base reaction the proton always reacts with the strongest base until totally
consumed before reacting with any weaker bases.
• Any substance whose anion is the conjugate base of a weak acid weaker than OH- reacts
quantitatively with water to form more hydroxide ions.
Step 1. NaCH3COO → Na+ + CH3COO-
Acetate ion is the conjugate base of acetic acid, a weak acid.
Step 2. CH3COO- + H2O CH3COOH + OH-
• Hydrolysis: Aqueous solutions of salts that dissociate into both:
1. A strong conjugate acid and a strong conjugate base are neutral (KNO3).
2. A strong conjugate acid and a weak conjugate base are acidic (HCl).
3. A strong conjugate base and a weak conjugate acid are basic (NaOH).
4. A weak conjugate base and a weak conjugate acid can be neutral, basic or acidic:
• The comparison of the values of Ka and Kb determine the pH of these solutions.
a. Kbase = Kacid make neutral solutions (NH4CH3OO)
b. Kbase > Kacid make basic solutions (NH4ClO)
c. Kbase < Kacid make acidic solutions (CH3)3NHF
stronger acid + stronger base weaker acid + weaker base
H2O
Acids, Bases, and ionization constants

Polyprotic Acids and Bases
• Polyprotic acids contain more than one
ionizable proton, and the protons are
lost in a stepwise manner.
• The fully protonated species is always
the strongest acid because it is easier to
remove a proton from a neutral
molecule than from a negatively
charged ion; the fully deprotonated
species is the strongest base.
• Acid strength decreases with the loss of
subsequent protons, and the pKa
increases.
• The strengths of the conjugate acids and
bases are related by pKa + pKb = pKw,
and equilibrium favors formation of the
weaker acid-base pair.

• Many weak acids contain two or more acidic hydrogens.
– Examples include H3PO4 and H3AsO4.
• The calculation of equilibria for polyprotic acids is done in a
stepwise fashion.
– There is an ionization constant for each step.
• Consider arsenic acid, H3AsO4, which has three ionization
constants.
1 Ka1 = 2.5 x 10-4
2 Ka2 = 5.6 x 10-8
3 Ka3 = 3.0 x 10-13
• Notice that the ionization constants vary in the following fashion:
• This is a general relationship.
– For weak polyprotic acids the Ka1 is always > Ka2, etc.
a3
a2
a1 K
K
K 

Polyprotic Acids

Polyprotic Acids
• A comparison of the various species in 0.100 M H3AsO4 solution follows.
Species Concentration
H3AsO4 0.095 M
H+ 4.9 x 10-3 M
H2AsO4
- 4.9 x 10-3 M
HAsO4
2- 5.6 x 10-8 M
AsO4
3- 3.4 x 10-18 M
OH- 2.0 x 10-12 M
When a strong base is added to a solution of a polyprotic acid, the
neutralization reaction occurs in stages.
1. The most acidic group is titrated first, followed by the next most acidic, and so
forth
2. If the pKa values are separated by at least three pKa units, then the overall
titration curve shows well-resolved “steps” corresponding to the titration of
each proton

• Silver chloride, AgCl,is rather insoluble in water.
• Careful experiments show that if solid AgCl is placed in pure water and
vigorously stirred, a small amount of the AgCl dissolves in the water.
• The equilibrium constant expression for this dissolution is called a
solubility product constant.
– Ksp = solubility product constant
The Solubility Product, Ksp

The Common Ion Effect in Solubility
Calculations
• Calculate the molar solubility of barium sulfate, BaSO4, in 0.010 M sodium
sulfate, Na2SO4, solution at 25oC. Compare this to the solubility of BaSO4
in pure water. (Example 20-3). (What is the common ion? How was a
common ion problem solved in Chapter 19?)
1. Write equations to represent the equilibria.
2. Substitute the algebraic representations of the concentrations into the Ksp
expression and solve for x.
• The molar solubility of BaSO4 in 0.010 M Na2SO4 solution is 1.1 x 10-8 M.
• The molar solubility of BaSO4 in pure water is 1.0 x 10-5 M.
– BaSO4 is 900 times more soluble in pure water than in 0.010 M sodium
sulfate!
– Adding sodium sulfate to a solution is a fantastic method to remove
Ba2+ ions from solution!
• If your drinking water were suspected to have lead ions in it, suggest a
method to prove or disprove this suspicion.

The Ion Product
• The ion product (Q) of a salt is the product of the
concentrations of the ions in solution raised to the same
powers as in the solubility product expression.
• The ion product describes concentrations that are not
necessarily equilibrium concentrations, whereas Ksp describes
equilibrium concentrations.
• The process of calculating the value of the ion product and
comparing it with the magnitude of the solubility product is a
way to determine if a solution is unsaturated, saturated, or
supersaturated and whether a precipitate will form when
solutions of two soluble salts are mixed.

The Ion Product
• Three possible conditions for an aqueous solution
of an ionic solid
1. Q < Ksp: the solution is unsaturated, and more of the ionic solid will
dissolve
2. Q = Ksp: the solution is saturated and at equilibrium
3. Q > Ksp: the solution is supersaturated, and ionic solid will
precipitate

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