Study Guide
Chapter 14 Principles of Neutralization Titrations
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University:
Kentucky State University -
Course:
CHE 102 | General Chemistry II Academic year:
2020
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Views:
360
Pages:
31
Author:
Quintin Ware
Chapter 14
Principles of Neutralization
Titrations
Like any titration, neutralization titrations depend on a chemical
reaction between the analyte and a standard reagent.
The point of chemical equivalence is indicated by a chemical
indicator or an instrumental measurement.
The discussion here focuses on the types of standard solutions and the
chemical indicators that are used for neutralization titrations. Standard Solutions
The standard solutions employed in neutralization titrations are strong
acids or strong bases because these substances react more completely
with an analyte than their weaker counterparts do and thus yield
sharper end points.
Standard solutions of acids are prepared by diluting concentrated
hydrochloric, perchloric, or sulfuric acid.
Standard solutions of bases are ordinarily prepared from solid sodium
or potassium and occasionally barium hydroxides. The concentrations
of these bases must be established by standardization. Acid/Base Indicators
Many substances display colors that depend on the pH of the solutions
in which they are dissolved.
An acid/base indicator is a weak organic acid or a weak organic base
whose undissociated form differs in color from its conjugate form.
e.g., the behavior of an acid-type indicator, HIn, is described by the
equilibrium
HIn + H2O
In- + H3O+
acid color
base color
The equilibrium for a base-type indicator, In, is
In + H2O
InH+ + OHbase color
acid color …continued…
The equilibrium-constant expression for the dissociation of an acidtype indicator takes the form
+
H 3O In
Ka
HIn
Rearranging leads to
HIn
H 3O Ka In
+
The hydronium ion concentration determines the ratio of the acid to
the conjugate base form of the indicator and thus determines the color
developed by the solution. …continued…
The human eye is not very sensitive to color differences in a solution containing a
mixture of HIn and In-, particularly when the ratio [Hln/ln] is greater than about 10 or
smaller than about 0.1.
The color imparted to a solution by a typical indicator appears to the
average observer to change rapidly only within the limited
concentration ratio of approximately 10 to 0.1
The average indicator, Hln, exhibits its pure acid color when
HIn 10
In 1
and its base color when
HIn 1
In 10
The color appears to be intermediate for ratios between these two
values. These ratios vary considerably from indicator to indicator. …continued…
For the full acid color,
[H3O+] = 10 Ka
and similarly for the full base color,
[H3O+] = 0.1 Ka
To obtain the indicator pH range, we take the
negative logarithms of the two expression:
pH (acid color) = -log (10Ka) = pKa + 1
pH (basic color) = -log (0.1Ka) = pKa – 1
Indicator pH range = pKa 1 CALCULATING pH IN TITRATIONS OF STRONG ACIDS
AND STRONG BASES
The hydronium ions in an aqueous solution of a strong acid have two
sources:
(1) the reaction of the acid with water and
(2) the dissociation of water itself.
In all but the most dilute solutions, however, the contribution from the
strong acid far exceeds that from the solvent. Thus, for a solution of
HCl with a concentration greater than about 10-6 M, we can write
[H3O+]= cHCl + [H3O+] cHCl
where [OH-] represents the contribution of hydronium ions from the
dissociation of water. An analogous relationship applies for a solution
of a strong base, such as sodium hydroxide. That is,
[OH-]= cNaOH + [OH-] cNaOH Titrating a Strong Acid with a Strong Base
We will be interested here, in calculating hypothetical titration curves
of pH versus volume of titrant. Three types of calculation must be done
to construct the hypothetical titration curve for titrating a solution of a
strong acid with a strong base. Each calculation corresponds to a
distinct stage in the titration:
1. Preequivalence
2. Equivalence
3. Postequivalence
In the preequivalence stage, we compute the concentration of the acid from its
starting concentration and the amount of base that has been added.
At the equivalence point, the hydronium and hydroxide ions are present in equal
concentrations, and the hydronium ion concentration is derived directly from the ionproduct constant for water.
In the postequivalence stage, the analytical concentration of the excess base is
computed, and the hydroxide ion concentration is assumed to be equal to or a
multiple of the analytical concentration. …continued…
Kw = [H3O+][OH-]
-log Kw = -log [H3O+][OH-]
= -log [H3O+] – log [OH-]
pKw = pH + pOH
-log 10-14 = 14.00
= pH + pOH The Effect of
Concentration
Titration curves for HCI with
NaOH.
Curve A: 50.00 mL of 0.0500
M HCI with 0.1000 M NaOH.
Curve B: 50.00 mL of 0.000500
M HCI with 0.001000 M
NaOH.
Choosing an
Indicator Titrating a Strong Base with a Strong Acid TITRATION CURVES FOR WEAK ACIDS
Four distinctly different types of calculations are needed to derive a titration curve for
a weak acid (or a weak base):
1. At the beginning, the solution contains only a weak acid or a weak base, and the
pH is calculated from the concentration of that solute and its dissociation
constant.
2. After various increments of titrant have been added (in quantities up to, but not
including, an equivalent amount), the solution consists of a series of buffers. The
pH of each buffer can be calculated from the analytical concentrations of the
conjugate base or acid and the residual concentrations of the weak acid or base.
3. At the equivalence point, the solution contains only the conjugate of the weak
acid or base being titrated (that is, a salt), and the pH is calculated from the
concentration of this product.
4. Beyond the equivalence point, the excess of strong acid or base titrant represses
the acidic or basic character of the reaction product to such an extent that the pH
is governed largely by the concentration of the excess titrant. Note that the total volume of solution
is present in both numerator and
denominator and thus cancels in the
expression for [H3O+]. Calculations
similar to this provide points on the
curve throughout the buffer region.
Data from these calculations are
presented in columnable 14-3. • The analytical concentrations of acid and conjugate base are
identical when an acid has been half neutralized.
• the buffer capacities of each of the solutions are at a maximum at
this point. These points are often called the half-titration points.
At the half-titration point in the titration of a weak acid,
[H3O+] = Ka or
pH = pKa
At the half-titration point in the titration of a weak base,
[OH-] = Kb or
pOH = pKb The Effect of Concentration
• Note that the initial pH values are higher and the
equivalence-point pH is lower for the more dilute solution
(curve B).
• At intermediate titrant volumes, however, the pH values
differ only slightly because of the buffering action
• The pH of buffers is largely independent of dilution.
Choosing an Indicator: The
Feasibility of Titration The Effect of Reaction Completeness
The pH change in the equivalence-point region becomes smaller as the acid becomes
weaker-that is, as the reaction between the acid and the base becomes less complete. TITRATION CURVES FOR WEAK BASES
Chapter 14 Principles of Neutralization Titrations
of 31
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