Chemistry – Dr. Hazlett
Acids/Bases - Chapter 19
Study Guide
I. Introduction
A. Acids (A) and Bases (B) are in solutions (solns)
1. Solns are homogeneous mixtures of 2+ substances
a. Solute and solvent
2. Electrolytes and non-electrolytes
a. Aqueous (aq) solns of ionic compounds conduct electricity – i.e.
flow of electrons, are electrolytes
b. Pure (distilled) water not a good conductor – no ionization and
therefore, no free electrons
c. A’s ionize in water – thus are electrolytes
(1) Ionization varies up to 100%
(2) Weaker acids ionize less than 100%
3. Ionization
a. Aqueous soln (water)  end up with H+ and OH- ions
b. The amounts of these ions in soln determine whether A or B
c. Water is universal solvent
(1) Undergoes self-ionization
(a) H2O  H+ + OH(b) H2O + H2O  H3O+ + OHII. Acids and Bases in General
A. Characteristics
1. Acids - Physical:
a. Taste sour
b. React with some metals (Al, Mg, Zn) to produce H2 (g)
c. React with carbonates to produce CO2 (g)
d. React with B to form salts and water
e. Change indicators to differing colors – most common is
litmus blue to red
f. Yields H+ ion or forms H3O+
2. Bases – Physical:
a. Taste bitter
b. Fell soapy/slippery
c. Changes red litmus paper to blue
d. Reacts with A’s to form salts and water (neutralization)
e. Reacts with some metals to produce H2 (g)
f. Yields the OH- ion
g. Alkalines
B. History and Definitions
(Source:
http://library.thinkquest.org/C006669/data/Chem/acidbases/development).
Acids and Bases : Development of Acids and Bases
1661 - Robert Boyle
Characterized acids and alkalies (bases) as the following:

Acids:
1. Sour taste
2. Corrosive
3. Change litmus (dye extracted from lichens) from blue
to red
4. Become less acidic when combined with alkalies.

Alkalies (Bases):
1. Feel slippery
2. Change litmus from red to blue
3. Become less alkaline when combined with acids.
Antoine Lavoisier

Believed that all acids contained oxygen after studying several
acids
e.g. H2SO4 - sulfuric acid, HNO3 - nitric acid
1811 - Humphry Davy


Questioned Lavoisier's theory, noting that hydrochloric acid
(HCl) did not contain oxygen yet is an acid.
Soon thereafter, several more acids without oxygen were
found.
e.g. HBr, HF, HI
1838 - Justig Liebig
Suggested that acids contain one or more hydrogen atoms which can be replaced
by metal atoms to produce salts.
e.g. HSCN is an acid because the H atom can be replace by a metal to form a salt,
such as NaSCN.
1884-1887 - Svante Arrhenius


In 1884, Arrhenius proposed that salts dissociate when they
dissolve in water to give charged particles which he called
ions.
In 1887, Arrhenius extended this his idea by defining acids and
bases as the following:
o Arrhenius acid - Any substance that ionizes when it
dissolves in water to give the H+ ion.
e.g.
o
Arrhenius base - Any substance that ionizes when it
dissolves in water to give the OH- ion.
e.g.
Arrhenius's theory helped to explain why acids have similar characteristics, since they all
give H+ ions when they dissolve in water. It also explained why acids are neutralized by
bases and why bases are neutralized by acids; the H+ ions from acids combine with the
OH- ions from bases to form water:
Though the Arrhenius theory helped to explain more about acids and bases, there were
still several drawbacks to this theory.
1. The theory can only classify substances when they are dissolved in water since
the definitions are based upon the dissociation of compounds in water.
2. It does not explain why some compounds containing hydrogen such as HCl
dissolve in water to give acidic solutions and why others such as CH4 do not.
3. The theory can only classify substances as bases if they contain the OH- ion and
cannot explain why some compounds that don't contain the OH- such as Na2CO3
have base-like characteristics.
To extend the Arrhenius theory a little further, consider the formation of water from the
combination of an H+ ion and an OH- ion:
This reaction is actually reversible, represented by the forward/backward arrow in the
following reaction:
Based on the fact that the above reaction is reversible, we can conclude the following
operational definitions for acids and bases.


Acid - Any substance that increases the concentration of the H+ ion when it
dissolves in water.
Base - Any substance that increase the concentration of the OH- ion when it
dissolves in water.
NOTE: A common shorthand notation for the concentration of a substance is placing it
in brackets:
[H+] = concentration of the H+ ion
Now, substances not containing H+ or OH- ions can be classified as acids or bases if they
alter the [H+] or [OH-] when they dissolve in water.
e.g. CO2 cannot dissociate to give H+ but it does increases [H+] when it is dissolves in
water.
e.g. CaO cannot dissociate to give OH- but it does increase [OH-] when it dissolves in
water.
1923 - Johannes Brønsted and Thomas Lowry
Johannes
Brønsted
In 1923, Johannes Brønsted and Thomas Lowry separately proposed a
new set of defintions for acids and bases which are known as either
Brønsted acids and bases or Brønsted-Lowry acids and bases.

Brønsted Acid - Any substance that can donate a proton, H+
ion to a base.
aka: hydrogen-ion donors or proton donors

Brønsted Base - Any substance that can accept a proton, H+
ion from an acid.
aka: hydrogen-ion acceptor or proton acceptor
Thomas
Lowry
In the above reaction, the H from HCl is donated to H2O which
accepts the H to form H3O+, leaving a Cl- ion.
The dissociation of water can be represented as follows:
The Brønsted-Lowry model of acids and bases brings rise to the concept of conjugate
acid-base pairs. The part of the acid remaining when an acid donates a H+ ion is called
the conjugate base. The acid formed when a base accepts a H+ ion is called the conjugate
acid. For the generic acid HA:
For the generic base A-:
More examples of conjugate acid-base pairs:
In the following reactions, it is shown how H2PO4- and H2O can act as both acids and
bases. Such compounds are said to be amphoteric.


Strong acids have weak conjugate bases.
Strong bases have weak conjugate acids.
Water has the tendency to equalize the strengths of all strong acids and strong bases,
regardless of the strength of the acid itself. This is known as the leveling effect. Acids
are limited to the strength of the H3O+ ions that they form when they lose H+ ions when
they dissolve in water. Likewise, bases are limited to the strength of the OH- ions that
they form when they gain H+ ions when they dissolve in water.
The relative strength of an acid is described as an acid-dissociation equilibrium constant.

Acid-dissociation equilibrium constant (Ka) - For the generic acid
reaction with water:
The acid-dissociation equilibrium constant is the mathematical product of the equilibrium
concentrations of the products of this reaction divided by the equilibrium concentration of
the original acid:
Strong acids dissociate almost completely in water and therefore have relatively large Ka
values. Weak acids, however, dissociate only slightly in water and therefore have
relatively small Ka values. To distinguish between strong and weak acids, the following
guidelines are used:
The relative strengths of acids and bases is discussed further in the Equilibrium notes.
Advantages to the Brønsted-Lowry model of acids and bases







Acids and bases can now be ions or neutral molecules.
Acids and bases can now be any molecule with at least one pair of nonbonding
electrons.
It explains the role of water in acid-base reactions; H2O accepts H+ ions from
acids to form H3O+ ions.
It can be applied to solutions with solvents other than water and even in reactions
that occur in the gas or solid phases.
It relates acids and bases to each other with conjugate acid-base pairs and can
explain their relative strengths.
It can explain the relative strengths of pairs of acids or pairs of bases.
It can explain the leveling effect of water.
1923 - G.N. Lewis
Proposed another method of defining acids and bases.

Lewis acid - Any substance that can accept a pair of
nonbonding electrons.
aka: electron-pair acceptor

Lewis base - Any substance that can donate a pair of
nonbonding electrons.
aka: electron-pair donor
In the following example, the Al3+ ion acts as an acid, accepting electron pairs from water
which acts as a base, an electron-pair donor. The two combine to form Al(H2O)63+, an
acid-base complex or a complex ion.
Acids and Bases : Typical Acids and Bases
Acids

Nonmetal hydrides - One or more hydrogen with a +1 oxidation state bound to
a nonmetal
e.g. HF, HCl, HBr, HI, H2S, HCN, and HSCN

Nonmetal oxides - One or more oxygen with a -2 oxidation state bound to a
nonmetal
e.g. CO2, NO2, SO2, and SO3

Nonmetal hydroxides (Oxyacids) - One or more hydroxide groups with an
overall oxidation state of -1 bound to a nonmetal
e.g. HOCl, HONO2, O2S(OH)2, and OP(OH)3
or more commonly known as: HClO, HNO3, H2SO4, and H3PO4
Bases

Metal hydrirdes - One or more hydrogen with a +1 oxidation state bound to a
metal
e.g. LiH, NaH, KH, MgH2, and CaH2

Metal oxides - One or more oxygen with a -2 oxidation state bound to a metal
e.g. Li2O, Na2O, K2O, MgO, and CaO

Metal hydroxides - One or more hydroxide groups with an overall oxidation
state of -1 bound to a metal
e.g. LiOH, NaOH, KOH, Ca(OH)2, and Ba(OH)2
Metal hydroxides are bases because the difference in the electronegativities of between
the metal and the oxygen is relatively large. The electronegativity of the oxygen is much
larger than that of the metal, so the electrons tend to remain with the oxygen atom,
resulting in the formation of a positive metal ion and the OH- ion.
Nonmetal hydroxides, however, are acids because the difference in the electronegativities
between the nonmetal and the oxygen is relatively small. The electronegativities of the
oxygen is more or less equal to that of the nonmetal, so the electrons are considered to be
shared. The electrons in the O-H bond remain with the oxygen, resulting in the formation
of the H+ ion andn oxyanion.
Amphoteric - Compounds that can act as either acids or bases. Ex.: WATER!!!!!
e.g. H2O, Al2O3, and AL(OH)2
Common Acids
Name
Formula
Where It's Found
acetic acid
HC2H3O2
vinegar
acetylsalicylic acid
HC9H7O4
aspirin
ascorbic acid
H2C6H6O6
vitamin C
carbonic acid
H2CO3
carbonated beverages
citric acid
H3C6H5O7
citrus fruits
hydrochloric acid
HCl
stomach acid
sulfuric acid
H2SO4
batteries
Common Bases
Name
aluminum hydroxide
Formula
Al(OH)3
Where It's Found
water purification
ammonia (aqueous solution)
NH3
household cleaners
calcium hydroxide
Ca(OH)2
mortar
magnesium hydroxide
Mg(OH)2
milk of magnesia (antacid and laxative)
potassium hydroxide
KOH
soap and glass making
sodium hydroxide
NaOH
drain and oven cleaners
II. Better Defined A and B
A. Bronsted-Lowry Definition
1. A is a cmpd that donates a p+; and B is a cmpd that accepts p+
a. Remember – p+ is nothing but a H+ ion
2. Remember – if it does/can do either – amphiprotic
a. ex. H2O  H+ + OH2H2O  H3O+ + OH3. General Formula:
a. HA (generic acid with H) + B  BH+ (accepted H+) + A- (lost p+)
b. A-B rxn is nothing but a p+ exchange rxn
4. A and B have conjugates:
a. A1 + B2  B1 (or Conjugate B) + A2 (or Conjugate A)
A gave p+ to B
Made when B took p+ from A
b. Acid (AH+) ---------- becomes --------- Conjugate B since lost p+
Base (OH-) ---------- becomes -------- Conjugate A since took p+
p+ “sinks” to form CA (BH+)
c. ex.: HF
A
+
H2 O
B

H3O+
CA
+
FCB
5. Rules:
a. H3O+ is strongest acid that can exist in water
b. All strong A are equally strong in water since they dissociate 100%
Chemistry Virtual Textbook → Acid-base concepts →
Proton donors and acceptors
Proton donors and acceptors
Acid-base reactions à la Brønsted
Acids are proton donors, bases are proton
acceptors.
1 Proton donors and acceptors
The older Arrhenius theory of acids and bases viewed them as substances
which produce hydrogen ions or hydroxide ions on dissociation. As useful
a concept as this has been, it was unable to explain why NH3, which
contains no OH– ions, is a base and not an acid, why a solution of FeCl3 is
acidic, or why a solution of Na2S is alkaline.
A more general theory of acids and bases was developed by Franklin in 1905,
who suggested that the solvent plays a central role. According to this view, an
acid is a solute that gives rise to a cation (positive ion) characteristic of the
solvent, and a base is a solute that yields a anion (negative ion) which is also
characteristic of the solvent. The most important of these solvents is of course
H2O, but Franklin's insight extended the realm of acid-base chemistry into nonaqueous systems as we shall see in a later lesson.
Brønsted acids and bases
In 1923 the Danish chemist J.N. Brønsted, building on Franklin's theory,
proposed that
An acid is a proton donor; a base is a proton acceptor.
In the same year the English chemist T.M. Lowry published a paper setting forth
some similar ideas without producing a definition; in a later paper Lowry himself
points out that Brønsted deserves the major credit, but the concept is still widely
known as the Brønsted-Lowry theory.
These definitions carry a very important implication: a substance cannot act as an
acid without the presence of a base to accept the proton, and vice versa. As a very
simple example, consider the equation that Arrhenius wrote to describe the behavior
of hydrochloric acid:
HCl → H+ + A–
This is fine as far as it goes, and chemists still write such an equation as a shortcut.
But in order to represent this more realistically as a proton donor-acceptor reaction,
we now depict the behavior of HCl in water by
in which the acid HCl donates its proton to the acceptor (base) H2O.
"Nothing new here", you might say, noting that we are simply replacing a shorter
equation by a longer one. But consider how we might explain the alkaline solution
that is created when ammonia gas NH3 dissolves in water. An alkaline solution
contains an excess of hydroxide ions, so ammonia is clearly a base, but because
there are no OH– ions in NH3, it is clearly not an Arrhenius base. It is, however, a
Brønsted base:
In this case, the water molecule acts as the acid, donating a proton to the base NH 3
to create the ammonium ion NH4+.
The foregoing examples illustrate several important aspects of the Brønsted-Lowry
concept of acids and bases:

A substance cannot act as an acid unless a proton acceptor (base) is present to
receive the proton;

A substance cannot act as a base unless a proton donor (acid) is present to supply the
proton;

Water plays a dual role in many acid-base reactions; H2O can act as a proton acceptor
(base) for an acid, or it can serve as a proton donor (acid) for a base (as we saw for
ammonia.

The hydronium ion H3O+ plays a central role in the acid-base chemistry of aqueous
solutions.
2 The Hydronium ion
Hydrogen ions cannot exist in water
There is another serious problem with the Arrhenius view of an acid as a substance
that dissociates in water to produce a hydrogen ion. The hydrogen ion is no more
than a proton, a bare nucleus. Although it carries only a single unit of positive
charge, this charge is concentrated into a volume of space that is only about a
hundred-millionth as large as the volume occupied by the smallest atom. (Think of a
pebble sitting in the middle of a sports stadium!) The resulting extraordinarily high
charge density of the proton strongly attracts it to any part of a nearby atom or
molecule in which there is an exess of negative charge. In the case of water, this will
be the lone pair (unshared) electrons of the oxygen atom; the tiny proton will be
buried within the lone pair and will form a shared-electron (coordinate) bond with it,
creating a hydronium ion, H3O+. In a sense, H2O is acting as a base here, and the
product H3O+ is the conjugate acid of water:
Owing to the overwhelming excess of H2O
molecules in aqueous solutions, a bare hydrogen
ion has no chance of surviving in water.
Although other kinds of dissolved ions have water molecules bound to them
more or less tightly, the interaction between H+ and H2O is so strong that writing
“H+(aq)” hardly does it justice, although it is formally correct. The formula H3O+
more adequately conveys the sense that it is both a molecule in its own right,
and is also the conjugate acid of water.
The equation "HA → H+ + A–" is so much easier to write that chemists still use it to
represent acid-base reactions in contexts in which the proton donor-acceptor
mechanism does not need to be emphasized. Thus it is permissible to talk about
“hydrogen ions” and use the formula H+ in writing chemical equations as long as you
remember that they are not to be taken literally in the context of aqueous solutions.
Interestingly, experiments indicate that the proton does not stick to a single H2O
molecule, but changes partners many times per second. This molecular
promiscuity, a consequence of the uniquely small size and mass the proton,
allows it to move through the solution by rapidly hopping from one H2O molecule
to the next, creating a new H3O+ ion as it goes. The overall effect is the same as
if the H3O+ ion itself were moving. Similarly, a hydroxide ion, which can be
considered to be a “proton hole” in the water, serves as a landing point for a
proton from another H2O molecule, so that the OH– ion hops about in the same
way.
Because hydronium- and hydroxide ions can “move without actually moving” and
thus without having to plow their way through the solution by shoving aside
water molecules as do other ions, solutions which are acidic or alkaline have
extraordinarily high electrical conductivities.
3 Acid-base reactions à la Brønsted
According to the Brønsted concept, the process that was previously written as a
simple dissociation of a generic acid HA ("HA → H+ + A–)" is now an acid-base
reaction in its own right:
HA + H2O → A–+ H3O+
The idea, again, is that the proton, once it leaves the acid, must end up somewhere;
it cannot simply float around as a free hydrogen ion.
Conjugate pairs
A reaction of an acid with a base is thus a proton exchange reaction; if the acid is
denoted by AH and the base by B, then we can write a generalized acid-base
reaction as
AH + B → A– + BH+
Notice that the reverse of this reaction,
BH+ + A– → B + AH+
is also an acid-base reaction. Because all simple reactions can take place in both
directions to some extent, it follows that transfer of a proton from an acid to a base
must necessarily create a new pair of species that can, at least in principle,
constitute an acid-base pair of their own.
In this schematic reaction, base1 is said to be conjugate to acid1, and acid2 is
conjugate to base2. The term conjugate means “connected with”, the implication
being that any species and its conjugate species are related by the gain or loss of
one proton.
The table below shows the conjugate pairs of a number of typical acid-base systems.
Some common conjugate acid-base pairs
acid
base
hydrochloric
acid
HCl
acetic acid
CH3CH2COOH CH3CH2COO–
acetate ion
nitric acid
HNO3
NO3–
nitrate ion
dihydrogen
phosphate ion
H2PO4–
HPO4–
monohydrogen
phosphate ion
hydrogen
sulfate ion
HSO4–
SO42–
sulfate ion
Cl–
chloride ion
hydrogen
carbonate
HCO3–
("bicarbonate")
ion
CO32–
carbonate ion
ammonium ion NH4+
NH3
ammonia
iron(III)
("ferric") ion
Fe(H2O)63+
Fe(H2O)5OH2+
water
H2O
OH–
hydroxide ion
hydronium ion
H3O+
H2O
water
Strong acids and weak acids
We can look upon the generalized acid-base reaction
as a competition of two bases for a proton:
Definition of a "strong" acid
If the base H2O overwhelmingly wins this tug-of-war, then the acid HA is said to be a
strong acid. This is what happens with hydrochloric acid and the other common
strong "mineral acids" H2SO4, HNO3, and HClO4:
hydrochloric acid
HCl + H2O → Cl– + H3O+
sulfuric acid
H2SO4 + H2O → HSO4– + H3O+
nitric acid
HNO3 + H2O → NO3– + H3O+
perchloric acid
HClO4 + H2O → ClO4– + H3O+
Solutions of these acids in water are really solutions of the ionic species shown in
heavy type on the right. This being the case, it follows that what we call a 1 M
solution of "hydrochloric acid" in water, for example, does not really contain a
significant concentration of HCl at all; the only real a acid present in such a solution
is H3O+!
These considerations give rise to two important rules:
H3O+ is the strongest acid that can exist in water;
All strong acids appear to be equally strong in water.
The leveling effect
The second of these statements is called the leveling effect. It means that although
the inherent proton-donor strengths of the strong acids differ, they are all completely
dissociated in water. Chemists say that their strengths are "leveled" by the solvent
water.
A comparable effect would be seen if one attempted to judge the strengths of
several adults by conducting a series of tug-of-war contests with a young child.
One would expect the adults to win overwhelmingly on each trial; their strengths
would have been "leveled" by that of the child.
Weak acids
Most acids, however, are able to hold on to their protons more tightly, so only a
small fraction of the acid is dissociated. Thus hydrocyanic acid, HCN, is a weak acid
in water because the proton is able to share the lone pair electrons of the cyanide ion
CN– more effectively than it can with those of H2O, so the reaction
HCN + H2O → H3O+ + CN–
proceeds to only a very small extent.
Since a strong acid binds its proton only weakly, while a weak acid binds it tightly,
we can say that
Strong acids are "weak"; Weak acids are "strong"
If you are able to explain this apparent paradox, you understand one of the most
important ideas in acid-base chemistry!
Examples of proton donor-acceptor reactions
reaction
acid
base
conjugate conjugate
acid
base
1) autoionization of water H2O
H2O
H2O
H3O+
OH–
2) ionization of hydrocyanic acid
HCN
HCN
H2O
H3O+
CN–
3) ionization of ammonia NH3 in
water
NH3
H2O
NH4+
OH–
4) hydrolysis of ammonium
chloride NH4Cl
NH4+
H2O
H3O+
NH3
5) hydrolysis of sodium acetate
CH3COO- Na+
H2O
CH3COO–
CH3COOH
OH–
6) neutralization of HCl by NaOH
HCl
OH–
H2O
Cl–
7) neutralization of NH3 by acetic
CH3COOH NH3
acid
NH4+
CH3COO–
8) dissolution of BiOCl (bismuth
oxychloride) by HCl
2 H3O+
BiOCl
Bi(H2O)3+
H2O, Cl–
9) decomposition of Ag(NH3)2+
by HNO3
2 H3O+
Ag(NH3)2+ NH4+
10) displacement of HCN by
CH3COOH
CH3COOH CN–
HCN
H2O
CH3COO–
Strong acids have weak conjugate bases
This is just a re-statement of what is implicit in what has been said above about the
distinction between strong acids and weak acids. The fact that HCl is a strong acid
implies that its conjugate base Cl– is too weak a base to hold onto the proton in
competition with either H2O or H3O+. Similarly, the CN– ion binds strongly to a
proton, making HCN a weak acid.
Salts of weak acids give alkaline solutions
The fact that HCN is a weak acid implies that the cyanide ion CN– reacts readily with
protons, and is thus is a relatively good base. As evidence of this, a salt such as
KCN, when dissolved in water, yields a slightly alkaline solution:
CN– + H2O → HCN + OH–
This reaction is still sometimes referred to by its old name hydrolysis ("water
splitting"), which is literally correct but tends to obscure its identity as just another
acid-base reaction. Reactions of this type take place only to a small extent; a 0.1M
solution of KCN is still, for all practical purposes, 0.1M in cyanide ion.
In general, the weaker the acid, the more alkaline will be a solution of its
salt. However, it would be going to far to say that "ordinary weak acids have
strong conjugate bases." The only really strong base is hydroxide ion, OH–,
so the above statement would be true only for the very weak acid H2O.
Strong bases and weak bases
The only really strong bases you are likely to encounter in day-to-day chemistry are
alkali-metal hydroxides such as NaOH and KOH, which are essentially solutions of
the hydroxide ion.
Most other compounds containing hydroxide ions such as Fe(OH)3 and Ca(OH)2
are not sufficiently soluble in water to give highly alkaline solutions, so they are
not usually thought of as strong bases.
There are actually a number of bases that are stronger than the hydroxide ion —
best known are the oxide ion O2– and the amide ion
NH2–, but these are so strong that they can rob water of a proton:
O2– + H2O → 2 OH–
NH2– + H2O → NH3 + OH–
This gives rise to the same kind of leveling effect we described for acids, with the
consequence that
Hydroxide ion is the strongest base that can exist in aqueous solution.
Salts of weak bases give acidic solutions
The most common example of this is ammonium chloride, NH4Cl, whose
aqueous solutions are distinctly acidic:
NH4+ + H2O → NH3 + H3O+
Because this (and similar) reactions take place only to a small extent, a solution
of ammonium chloride will only be slightly acidic.
Autoprotolysis
From some of the examples given above, we see that water can act as an acid
CN– + H2O → HCN + OH–
and as a base
NH4+ + H2O → NH3 + H3O+
If this is so, then there is no reason why "water-the-acid" cannot donate a proton to
"water-the-base":
This reaction is known as the autoprotolysis of water.
Chemists still often refer to this reaction as the "dissociation" of water and use
the Arrhenius-style equation H2O → H+ + OH– as a kind of shorthand.
This process occurs to only a tiny extent. It does mean, however, that hydronium
and hydroxide ions are present in any aqueous solution.
Can other liquids exhibit autoprotolysis? The answer is yes. The most well-known
example is liquid ammonia:
2 NH3 → NH4+ + NH2–
Even pure liquid sulfuric acid can play the game:
2 H2SO4→ H3SO4+ + HSO4–
Each of these solvents can be the basis of its own acid-base "system", parallel to the
familiar "water system".
Ampholytes
Water, which can act as either an acid or a base, is said to be amphiprotic: it can
"swing both ways". A substance such as water that is amphiprotic is called an
ampholyte.
As indicated here, the hydroxide ion can also be an ampholyte, but not in aqueous
solution in which the oxide ion cannot exist.
It is of course the amphiprotic nature of water that allows it to play its special role in
ordinary aquatic acid-base chemistry. But many other amphiprotic substances can
also exist in aqueous solutons. Any such substance will always have a conjugate acid
and a conjugate base, so if you can recognize these two conjugates of a substance,
you will know it is amphiprotic.
The carbonate system
For example, the triplet set {carbonic acid, bicarbonate ion, carbonate ion}
constitutes an amphiprotric series in which the bicarbonate ion is the ampholyte,
differing from either of its neighbors by the addition or removal of one proton:
If the bicarbonate ion is both an acid and a base, it should be able to exchange a
proton with itself in an autoprotolysis reaction:
HCO3– + HCO3– → H2CO3 + CO32–
Your very life depends on the above reaction! CO2, a metabolic by-product of every cell in your body,
reacts with water to form carbonic acid H2CO3 which, if it were allowed to accumulate, would make your
blood fatally acidic. However, the blood also contains carbonate ion, which reacts according to the reverse
of the above equation to produce bicarbonate which can be safely carried by the blood to the lungs. At this
location the autoprotolysis reaction runs in the forward direction, producing H2CO3 which loses water to
form CO2 which gets expelled in the breath. The carbonate ion is recycled back into the blood to eventually
pick up another CO2 molecule.
If you can write an autoprotolysis reaction for a substance, then that substance is amphiprotic.
Concept Map:
(Source: www.chem1.com/acad/webtext/abcon/abcon-3.html)
B. Lewis AB Definition
1. Definition is based on bonding and structure
a. Does not require H ion or a solvent
2. A is a substance that accepts a pair of e- and forms a covalent bond with
the supplier; and a B donates a pair of unshared e- to the recipient with which
the e-‘s can be shared
1 Lewis A-B: Protons and electron-pairs
According to Lewis,

An acid is a
substance that
accepts a pair of
electrons, and in
doing so, forms a covalent bond with the entity that supplies the
electrons;

A base is a substance that donates an unshared pair of electrons to a
recipient species with which the electrons can be shared.
In modern chemistry, electron donors are often referred to as nucleophiles, while
acceptors are electrophiles.
Proton-transfer reactions involve electron-pair transfer
Just as any Arrhenius acid is also a Brønsted acid, any Brønsted acid is also a Lewis
acid, so the various acid-base concepts are all "upward compatible". Although we
don't really need to think about electron-pair transfers when we deal with ordinary
aqueous-solution acid-base reactions, it is important to understand that it is the
opportunity for electron-pair sharing that enables proton transfer to take place.
This equation for a simple acid-base neutralization shows how the Brønsted and
Lewis definitions are really just different views of the same process. Take special
note of the following points:

The arrow shows the movement of a proton from the hydronium ion to the
hydroxide ion.

Note carefully that the electron-pairs themselves do not move; they remain
attached to their central atoms. The electron pair on the base is "donated" to
the acceptor (the proton) only in the sense that it ends up being shared with
the acceptor, rather than being the exclusive property of the oxygen atom in
the hydroxide ion.

Although the hydronium ion is the nominal Lewis acid here, it does not itself
accept an electron pair, but acts merely as the source of the proton that
coordinates with the Lewis base.
The point about the electron-pair remaining on the donor species is especially
important to bear in mind. For one thing, it distinguishes a Lewis acid-base
reaction from an oxidation-reduction reaction, in which a physical transfer of
one or more electrons from donor to acceptor does occur.
The product of a Lewis acid-base reaction is known formally as an "adduct" or
"complex", although we don't ordinarily use these terms for simple proton-transfer
reactions such as the one in the above example. Here, the proton combines with the
hydroxide ion to form the "adduct" H2O.
The following examples illustrate these points for some other proton-transfer
reactions that you should already be familiar with.
2 Acid-base reactions without protons
The major utility of the Lewis definition is that it extends the concept of acids and
bases beyond the realm of proton transfer reactions.
The classic example, given in every textbook, is the reaction of boron trifluoride with
ammonia to form an adduct:
BF3 + NH3 → F3B-NH3
One of the most commonly-encountered kinds of Lewis acid-base reactions occurs
when electron-donating ligands form coordination complexes with transition-metal
ions.
In this example, the tin atom in SnCl 4 can expand its valence shell by utilizing a pair
of d-orbitals, changing its hybridization from sp3 to sp3d2.
Here are several more examples of Lewis acid-base reactions that cannot be
accommodated within the Brønsted or Arrhenius models; be sure you are able to
identify the Lewis acid and Lewis base in each reaction.
Al(OH)3 + OH– → Al(OH)4–
SnS2 + S2–→ SnS32–
Cd(CN)2 + 2 CN–→ Cd(CN)42+
AgCl + 2 NH3→ Ag(NH3)2+ + Cl–
Fe2+ + NO → Fe(NO)2+
Ni2+ + 6 NH3→ Ni(NH3)52+
Applications to organic reaction mechanisms
Although organic chemistry is beyond the scope of these lessons, it is instructive to
see how electron donors and acceptors play a role in chemical reactions. The
following two diagrams (from this site by William Reusch of Michigan State U.) show
the mechanisms of two common types of reactions initiated by simple inorganic
Lewis acids:
In each case, the species labeled "Complex" is an intermediate that decomposes into
the products, which are conjugates of the original acid and base pairs. The electric
charges indicated in the complexes are formal charges, but those in the products are
"real".
In (1), the incomplete octet of the aluminum atom in AlCl 3 serves as a better
electron acceptor to the chlorine atom than does the isobutyl part of the base. In
(2), the pair of non-bonding electrons on the dimethyl ether coordinates with the
electron-deficient boron atom, leading to a complex that breaks down by releasing a
bromide ion.
3. Non-aqueous protonic acid-base systems
We ordinarily think of Brønsted-Lowry acid-base reactions as taking place in aqueous
solutions, but this need not always be the case. A more general view encompasses a
variety of acid-base solvent systems, of which the water system is only one. Each of
these has as its basis an amphiprotic solvent (one capable of undergoing
autoprotolysis), in parallel with the familiar case of water.
The ammonia system
This is the one you are most likely to encounter if you do advanced work in
Chemistry. Liquid ammonia boils at –33° C, and can conveniently be maintained as a
liquid by cooling with dry ice (–77° C.) It is a good solvent for substances that also
dissolve in water, such as ionic salts and organic compounds capable of forming
hydrogen bonds.
Many other familiar substances can serve as the basis of protonic solvent systems, a
few of which are listed here:
solvent
autoprotolysis reaction
pKap
water
2 H2O → H3O + OH
ammonia
2 NH3 → NH4 + NH2
acetic acid
2 CH3COOH → CH3COOH2 +
CH3COO–
13
ethanol
2 C2H5OH → C2H5OH2+ + C2H5O–
19
hydrogen
peroxide
2 HO-OH → HO-OH2+ + HO-O–
13
hydrofluoric acid
2 HF → H2F+ + F–
10
sulfuric acid
2 H2SO4 → H3SO4 +
+
–
+
14
–
33
+
+
HSO4–
3.5
Un-leveling the strong acids
One use of nonaqueous acid-base systems is to examine the relative strengths of the
strong acids and bases, whose strengths are "leveled" by the fact that they are all
totally converted into H3O+ or OH– ions in water. By studying them in appropriate
non-aqueous solvents which are poorer acceptors or donors of protons, their relative
strengths can be determined.
Concept Map
(Source: www.chem1.com/acad/webtext/abcon/abcon-5.html)
III. Titrations and pH/pOH Scales
A. Titration
1. A method of finding the amount of an A by finding out how many
moles of B needed to neutralize it
a. Substance whos concentration is being found is “analyte”
b. Additive is called “Titrate”
c. Use pH indicators to find the “equivalence point”
d. Graph results on titration curve to discover neutralization point
B. Measuring AB with pH and pOH
1. Based on the dissociation (Do) of water
2. Rxn of A and B depends on H+ ions combining with OH- ions to form water
a. This is the Do of water
b. Therefore, the scales measure the molal concentration (mol/L) of the
H+ and OH- ions
c. [H+] acid and [OH-] base = 1.008 x 10-7 mol/L = pH of 7 = neutral
(1) [H+] x [OH-] = 1.00 x 10-14 = pH of 14
(2) If:
[H+] > [OH-] = A
[H+] < [OH-] = B
[H+] = [OH-] = pH = 7 and soln is neutral
Acids and Bases : pH
1909 - S. P. L. Sørenson - Danish biochemist who suggested the use of a logarithmic
scale to express the concentration of the H3O+ ion
pH - The negative of the log of the H3O+ (hydronium) ion concentration:
pOH - The negative of the log of the OH- (hydroxide) concentration:
Water-dissociation equilibrium constant (Kw) - The product of the
equilibrium concentration of the H3O+ and OH- ions in an aqueous solution is
equal to 1.00 x 10-14 at 25 C.
NOTE: When the pH is doubled, the [H3O+ ] decreases by a factor of 100.
When the pH is quadrupled, the [H3O+ ] decreases by a factor of 10,000



A substance is acidic if the pH is less than 7.
A substance is basic if the pH is greater than 7.
A substance is neutral if the pH is equal to 7.
In pure water, the concentration of the hydronium and hydroxide ions are equal.
Relationship Between[H3O+] and [OH-]
with pH and pOH of an Aqueous Solution
[H3O+] (M)
1.0
1.0 x 10-1
1.0 x 10-2
1.0 x 10-3
1.0 x 10-4
1.0 x 10-5
[OH-] (M)
1.0 x 10-14
1.0 x 10-13
1.0 x 10-12
1.0 x 10-11
1.0 x 10-10
1.0 x 10-9
pH
0
1
2
3
4
5
pOH
14
13
12
11
10
9
1.0 x 10-6
1.0 x 10-8
6
8
-7
-7
1.0 x 10
1.0 x 10-8
1.0 x 10-9
1.0 x 10-10
1.0 x 10-11
1.0 x 10-12
1.0 x 10-13
1.0 x 10
1.0 x 10-6
1.0 x 10-5
1.0 x 10-4
1.0 x 10-3
1.0 x 10-2
1.0 x 10-1
7
8
9
10
11
12
13
7
6
5
4
3
2
1
1.0 x 10-14
1.0
14
0
Acidic
Neutral
Basic
pH of 0.1 M Solutions of Common Acids and Bases
Compound Name
Formula
pH
Compound
Name
Formula
pH
hydrochloric acid
HCl
1.1
sulfuric acid
H2SO4
1.2
sodium sulfate
Na2SO4
6.1
sodium bisulfate
NaHSO4
1.4
sodium chloride
NaCl
6.4
sulfurous acid
H2SO3
1.5
sodium acetate
NaC2H3O2
8.4
phosphoric acid
H3PO4
1.5
sodium
bicarbonate
NaHCO3
8.4
hydrofluoric acid
HF
2.1
Na2HPO4
9.3
2.9
sodium
biphosphate
acetic acid
HC2H3O2
carbonic acid
H2CO3
2.8
sodium sulfite
NaSO3
9.8
hydrogen sulfide
H2S
4.1
sodium dihydrogen
phosphate
NaH2PO4
4.4
ammonium chloride
NH4Cl
4.6
hydrocyanic acid
HCN
5.1
sodium cyanide
NaCN
11.0
ammonia
NH3
11.1
sodium
carbonate
NaCO3
11.6
sodium
phosphate
Na3PO4
12.0
sodium
hydroxide
NaOH
13.0
1 Dissociation of water
All chemical reactions that take place in a single phase (such as in a solution) are theoretically
"incomplete" and are said to be reversible.
The ability of acids to react with bases depends on the tendency of hydrogen ions to
combine with hydroxide ions to form water:
H+(aq) + OH–(aq) → H2O
(1)
This tendency happens to be very great, so the reaction is practically complete— but not
"completely" complete; a few stray H+ and OH– ions will always be present. What's more,
this is true even if you start with the purest water attainable. This means that in pure
water, the reverse reaction, the "dissociation" of water
H2O → H+(aq) + OH–(aq)
(2)
will proceed to a very slight extent. Both reactions take place simultaneously, but (1) is so
much faster than (2) that only a minute fraction of H2O molecules are dissociated.
Liquids that contain ions are able to conduct an electric current. Pure water is practically an
insulator, but careful experiments show that even the most highly purified water exhibits a
very slight conductivity that corresponds to a concentration of both the H + ion and OH– ions
of almost exactly 1.00 × 10–7mol L–1 at 25°C.
Problem Example 1
What fraction of water molecules in a litre of water are dissociated
Solution: 1 L of water has a mass of 1000 g. The number of moles in 1000 g of H 2O is
(1000 g)/(18 g mol–1) = 55.5 mol. This corresponds to (55.5 mol)(6.02E23 mol-1) = 3.34E25
H2O molecules.
An average of 10-7 mole, or (10-7)(6.02E23) = 6.0E16 H2O molecules will be dissociated at
any time. The fraction of dissociated water molecules is therefore (6.0E16)/(3.3E25) =
1.8E–9.
Thus we can say that only about two out of every billion (109) water molecules will be
dissociated.
Ion product of water
The degree of dissociation of water is so small that you might wonder why it is even
mentioned here. The reason stems from an important relationship that governs the
concentrations of H+ and OH– ions in aqueous solutions:
[H+][OH–] = 1.00 × 10–14
(3) must know this!
The quantity 1.00 x 10–14 is commonly denoted by Kw. Its value varies slightly with temperature,
pressure, and the presence of other ions in the solution.
in which the square brackets [ ] refer to the concentrations (in moles per litre) of the
substances they enclose.
This expression is known as the ion product of water, and it applies to all aqueous solutions,
not just to pure water. The consequences of this are far-reaching, because it implies that if
the concentration of H+ is large, that of OH– will be small, and vice versa. This means that
H+ ions are present in all aqueous solutions, not just acidic ones.
This leads to the following important definitions, which you must know:
acidic solution
[H+] > [OH–]
alkaline ("basic")
solution
[H+] < [OH–]
neutral solution
[H+] = [OH–] = 1.00×10–7 mol
L–1
A neutral solution is one in which the concentrations of H+ and OH– ions are identical.
The values of these concentrations are constrained by Eq. 3. Thus, in a neutral solution,
both the hydrogen- and hydroxide ion concentrations are 1.00 × 10–7 mol L–1:
[H+][OH–] = [1.00 × 10–7][1.00 × 10–7] =1.00 × 10–14
Hydrochloric acid is a typical strong acid that is totally dissociated in solution:
HCl → H+(aq) + Cl–(aq)
A 1.0M solution of HCl in water therefore does not really contain any significant
concentration of HCl molecules at all; it is a solution in of H + and Cl– in which the
concentrations of both ions are 1.0 mol L–1. The concentration of hydroxide ion in such a
solution, according to Eq 2, is
[OH–] = (Kw)/[H+] = (1.00 x 10–14) / (1 mol L–1) = 1.00 x 10–14 mol L–1.
Similarly, the concentration of hydrogen ion in a solution made by dissolving 1.0 mol of
sodium hydroxide in water will be 1.00 x 10–14 mol L–1.
2 pH
When dealing with a range of values (such as the variety of hydrogen ion concentrations
encountered in chemistry) that spans many powers of ten, it is convenient to represent
them on a more compressed logarithmic scale. By convention, we use the pH scale to denote
hydrogen ion concentrations:
pH = – log10 [H+]
(4) must know this!
or conversely, [H+] = 10–pH .
This notation was devised by the Danish chemist Søren Sørensen (18681939) in 1909. There are several accounts of why he chose "pH"; a likely
one is that the letters stand for the French term pouvoir hydrogène,
meaning "power of hydrogen"— "power" in the sense of an exponent. It has
since become common to represent other small quantities in "p-notation".
Two that you need to know in this course are the following:
pOH = – log10 [OH–]
pKw = – log Kw (= 14 when Kw = 1.00 × 10–14)
Note that pH and pOH are expressed as numbers without any units, since logarithms must
be dimensionless.
Recall from Eq 3 that
it becomes
[H+][OH–] = 1.00 × 10–14; if we write this in "p-notation"
pH + pOH = 14
(5) must know this!
In a neutral solution at 25°C, pH = pOH = 7.0. As pH increases, pOH diminishes; a higher pH
corresponds to an alkaline solution, a lower pH to an acidic solution. In a solution with [H +] =
1 M , the pH would be 0; in a 0.00010 M solution of H+, it would be 4.0. Similarly, a 0.00010
M solution of NaOH would have a pOH of 4.0, and thus a pH of 10.0. It is very important
that you thoroughly understand the pH scale, and be able to convert between [H +] or [OH–]
and pH in both directions.
Problem Example 2
The pH of blood must be held very close to 7.40. Find the hydroxide ion concentration that
corresponds to this pH.
Solution: The pOH will be (14.0 – 7.40) = 6.60.
[OH–] = 10–pOH = 10–6.6 = 2.51 x 10–7 M
The pH scale
The range of possible pH values runs from about 0 to 14.
The word "about" in the above statement reflects the fact that at very high
concentrations (10 M hydrochloric acid or sodium hydroxide, for example,) a significant
fraction of the ions will be associated into neutral pairs such as H+·Cl–, thus reducing
the concentration of “available” ions to a smaller value which we will call the effective
concentration. It is the effective concentration of H+ and OH– that determines the pH
and pOH. For solutions in which ion concentrations don't exceed 0.1 M, the formulas
pH = –log [H+] and pOH = –log[OH–] are generally reliable, but don't expect a 10.0 M
solution of a strong acid to have a pH of exactly –1.00!
Tables will help give you a general feeling for where common substances fall on the pH scale.
Notice especially that

most foods are slightly acidic;

the principal "bodily fluids" are slightly alkaline, as is seawater— not surprising,
since early animal life began in the oceans.

the pH of freshly-distilled water will drift downward as it takes up carbon
dioxide from the air; CO2 reacts with water to produce carbonic acid, H2CO3.

the pH of water that occurs in nature varies over a wide range. Groundwaters
often pick up additional CO2 respired by organisms in the soil, but can also
become alkaline if they are in contact with carbonate-containing sediments.
"Acid" rain is by definition more acidic than pure water in equilibrium with
atmospheric CO2, owing mainly to sulfuric and nitric acids that originate from
fossil-fuel emissions of nitrogen oxides and SO2.
pH indicators
The colors of many dye-like compounds depend on the pH, and can serve as useful indicators
to determine whether the pH of a solution is above or below a certain value.
Universal indicators
Most indicator dyes show only one color change, and thus
are only able to determine whether the pH of a solution is
greater or less than the value that is characteristic of a
particular indicator. By combining a variety of dyes whose
color changes occur at different pHs, a "universal" indicator
can be made. Commercially-prepared pH test papers of this
kind are available for both wide and narrow pH ranges.
3 Titration
Since acids and bases readily react with each other, it is experimentally quite easy to find
the amount of acid in a solution by determining how many moles of base are required to
neutralize it. This operation is called titration, and you should already be familiar with it
from your work in the Laboratory.
We can titrate an acid with a base, or a base with an acid. The substance whose
concentration we are determining (the analyte) is the substance being titrated; the
substance we are adding in measured amounts is the titrant. The idea is to add titrant until
the titrant has reacted with all of the analyte; at this point, the number of moles of titrant
added tells us the concentration of base (or acid)
in the solution being titrated.
36.00 ml of a solution of HCl was titrated with
0.44 M KOH. The volume of KOH solution
required to neutralize the acid solution was 27.00
ml. What was the concentration of the HCl?
Solution: The number of moles of titrant added
was
(.027 L)(.44 mol L–1) = .0119 mol. Because one mole
of KOH reacts with one mole of HCl, this is also
the number of moles of HCl; its concentration is
therefore
(.0119 mol) ÷ (.036 L) = 0.33 M .
Titration curves
The course of a titration can be followed
by plotting the pH of the solution as a
function of the quantity of titrant added.
The figure shows two such curves, one for
a strong acid (HCl) and the other for a
weak acid, acetic acid, denoted by HAc.
Looking first at the HCl curve, notice how
the pH changes very slightly until the acid
is almost neutralized. At that point, which
corresponds to the vertical part of the
plot, just one additional drop of NaOH
solution will cause the pH to jump to a very
high value— almost as high as that of the pure NaOH solution.
Compare the curve for HCl with that of HAc. For a weak acid, the pH jump near the
neutralization point is less steep. Notice also that the pH of the solution at the
neutralization point is greater than 7. These two characteristics of the titration curve for a
weak acid are very important for you to know.
If the acid or base is polyprotic, there will be a jump in pH for each proton that is titrated.
In the example shown here, a solution of carbonic acid H 2CO3 is titrated with sodium
hydroxide. The first equivalence point (at which the H2CO3 has been converted entirely into
bicarbonate ion HCO3–) occurs at pH 8.3. The solution is now identical to one prepared by
dissolving an identical amount of sodium bicarbonate in water.
Addition of another mole equivalent of hydroxide ion converts the bicarbonate into
carbonate ion and is complete at pH 10.3; an identical solution could be prepared by
dissolving the appropriate amount of sodium carbonate in water.
Finding the equivalence point: indicators
When enough base has been added to react completely with the hydrogens of a monoprotic
acid, the equivalence point has been reached. If a strong acid and strong base are titrated,
the pH of the solution will be 7.0 at the equivalence point. However, if the acid is a weak
one, the pH will be greater than 7; the “neutralized” solution will not be “neutral” in terms
of pH. For a polyprotic acid, there will be an equivalence point for each titratable hydrogen
in the acid. These typically occur at pH values that are 4-5 units apart, but they are
occasionally closer, in which case they may not be readily apparent in the titration curve.
The key to a successful titration is knowing when the equivalance point has been reached.
The easiest way of finding the equivalence point is to use an indicator dye; this is a
substance whose color is sensitive to the pH. One such indicator that is commonly
encountered in the laboratory is phenolphthalein; it is colorless in acidic solution, but turns
intensely red when the solution becomes alkaline. If an acid is to be titrated, you add a few
drops of phenolphthalein to the solution before beginning the titration. As the titrant is
added, a local red color appears, but quickly dissipates as the solution is shaken or stirred.
Gradually, as the equivalence point is approached, the color dissipates more slowly; the trick
is to stop the addition of base after a single drop results in a permanently pink solution.
Different indicators change color at different pH values; see here for an illustrated list.
Since the pH of the equivalance point varies with the strength of the acid being titrated,
one tries to fit the indicator to the particular acid. One can titrate polyprotic acids by using
a suitable combination of several indicators, as is illustrated above for carbonic acid.
Concept Map (Source: www.chem1.com/acad/webtext/abcon.abcon-2.html)
C. Titration and Buffers:
A buffer is a solution that is resistant to changes in pH. The desired buffer pH determines what
compounds are used to make buffer. To make a buffer with an acidic pH, ( pH less than 7) the
solution is made using a weak acid and a soluble salt of its anion. A buffer with a pH of 4.74 can
be made using a equal volumes of 0.10 M acetic acid, HC2H3O2, and 0.10 M sodium acetate,
NaC2H3O2. Typically the range for an acid buffer is centered around the value of the dissociation
constant for the weak acid. The pH is adjusted by controlling the ratio of the mols of weak acid
and the mols of soluble salt. The "recipe" for preparing buffers is illustrated below.
A rule of thumb is that an acid buffer can be made using a weak acid and its salt.
To make a buffer solution with a basic pH, (pH is more than 7) the solution is typically made using
a weak base and a soluble salt of its cation.
Buffers have at least two limitations. The natural pH range that matches the properties of the
solute acid or base and the capacity of the buffer. One limitation depends on the equlibrium
constant for the weak acid or base. The second comes from the limited solubility of the acids,
bases or their salts. The amount of dissolved acid, base and salt determines the mols available in
the buffer that cn react with "added" acids or bases from the outside.
HOW IT WORKS
A buffer can stabilize pH and acidity because it contains a reservoir of acid and a reservoir of
base. The acid can react with added base. The anion of the weak acid "behaves" as a base so it
reacts with added acid. A buffer of acetic acid has molecules of HC 2H3O2 that can react with any
base added to the solution. The same buffer mixture contains acetate ions, C 2H3O21- that can
react with protons(hydronium ions) from any acid that is added to the buffer.
A famous ( in biochemistry circles) relationship fo predicting pH for acid buffers is the HendersonHasselbalch equation.
The equation shows that the pH range for the buffer is tied to the Ka for the acid. The pH can be
adjusted by controlling the molarity of the acid, [HA], and the molarity for the anion of the acid,
[A1-]. This really requires control of the ratio of mols of anion to mols of acid.
Henderson-Hasselbalch equation
pH = pKa + log [A1-]/[HA]
here
[A1-] is the molarity of the anion of the weak acid;
[HA] is the molarity of the weak acid;
pKa is the -log of the weak acid dissociation
constant;
log means the base ten logarithm.
Example: How the pH for a buffer relates to pKa.
What is pH of a solution made using equal numbers of mols of acetic acid,HC 2H3O2, and mols of
sodium acetate, NaC2H3O2 ? The ratio of mols of each will be the same as the ratio of the
molarities (concentrations) because they are dissolved in the same final volume of solution.
1. The Henderson-Hasselbalch equation relates the ratio of mols of anion and mols of weak acid
to pH.
2. Recall pH = pKa + log [A1-]/[HA]
3. Collect the data
desired pH = ? ;
for acetic acid, HC2H3O, from references Ka = 1.8 x 10-5;
from references or by calculating pKa = -log Ka ; pKa= 4.7
pKa = -log(1.8 x 10-5) = 4.7
[C2H3O2 1-] = [HC2H3O2] so ratio [C2H3O2 1-]/[HC2H3O2] = 1
4. Substitute into the equation;
pH = pKa + log [A1-]/[HA]
pH = 4.7 + log [C2H3O2 1-]/[HC2H3O2]
pH = 4.7 + log (1) = 4.7 + 0 = 4.7
Recall the log of 1 is zero.
pH = 4.7
5. The ratio of the concentration of the anion and the weak acid can be adjusted to change the pH
Typically the pH can be set to one pH unit above or one pH unit below the pKa for the acid. This
matches a limit for the ratio of 1 to 10
(source: http://www.800mainstreet.com/acid_base/buffers.html)
V. Acid-Base Strength
A. Strong Acid – completely ionize/dissociate in aqueous soln
1. Only 7: HCl, HBr, HI, H2SO4, HNO3, HClO3, HClO4
a. Sometimes H3O+, HBrO3, HBrO4, HIO3, and HIO4 included
2. Weak Acids – do not ionize/dissociate completely
3. Strength depends on polarity of bond between H and the element(s) to
which it is bonded
a. Connected to the level of Do
b. Strength increases when polarity increases with decrease in
bond energy
B. Superacids – pH way beyond 14 – up to 2 x 1019 stronger than H2SO4
C. Acid Ionization Constant (Ka)
Ka = [H+] [A-]
[HA]
=
[H3O+] [OH-]
[HA]
[ ] is M conc. = (mol/L)
HA is acid
OH- is base
D. Strong Base – completely ionize/dissociate in soln
1. Only 7: LiOH, NaOH, KOH, RbOH, Ca(OH)2, Sr(OH)2, Ba(OH)2
2. Weak bases do not ionize completely
E. Base Equilibrium Constant (Kb)
Kb = [Conjugate A] [OH-]
[Weak B] [H2O]
=
[HA] [OH-]
[A-]
[ ] = M = n/vol
A- is base
Kw = [H3O+] [OH-]
Ka x Kb = Kw = 1.008 x 10-14
F. Salts from AB Rxns
1. Salt is an ionic cmpd made from cation of a B and an anion of an A
a. Metal with a nonmetal  salt
2. Salts react w/ water in process called “salt hydrolysis”
a. Anions dissociated from salt accept H+ ions from water or the
dissociated cations from the salt donate H+ ions to the water
3. Salts of a WA give an alkaline soln; of a WB give an acidic soln
a. Types:
(1) Acid salts
NaHCO3
(2) Basic/Alkaline
Pb(OH)NO3
(3) Double
CaMg(CO3)2
(4) Complex
K3Fe(CN)6
(5) Hydrates
MgSO4.7H2O
b. Zwitterions (Organic Chemistry)
(1) Amino acids ( - NH2) can accept p+; and carboxyl
grps (-COOH) can lose p+
(2) Resulting molecule is double ion called zwitterion
VI. Types of Acids and Bases
A. Binary Hydrides
1. MHn where M is any element and n is a real number
B. Hydroxy Compounds
1. M-OH where M is a metal and bonded with OH (ex. NaOH)
2. If M is nonmetal – then have an oxyacid (ex. HPO3)
C. Oxygen Compounds (M-O)
1. Binary oxides w/ no H but exhibit ab behaviors when rxn w/ H2O
a. Grps I and II bonded with OD. Organics (M-COOH)
E. Alkalies (M-OH)
1. Grps I and II bonded with OH
a. NH3 is an exception
VII. Nomenclature of AB
A. Acids
1. If simple anion – prefix of “hydro” and suffix of “ic”
2. If polyatomic anion – suffix of “ic” if the anion ends with “ate”
a. Use “ous” if it ends in “ite”
B. Bases end with “ide”
1. Exception is NH3 - ammonia