MKG BSMT-IB
CHAPTER 14:
NEUTRALIZATION TITRATION
 Used to determine the amounts of acids and bases
 Used to monitor the progress of reactions that
produce or consume hydrogen ions.
 Depends on a chemical reaction of the analyte with
a standard reagent
 TYPES OF ACID/BASE TITRATIONS:
 Titration of STRONG ACID ( HCl, H2SO4) with
STRONG BASE (NaOH)
 Titration of WEAK ACID (acetic acid, lactic
acid) with STRONG BASE
 Titration of WEAK BASES (sodium cyanide,
sodium salicylate) with strong acid
 Methods of determining/ locate the end point:
 Use of a chemical indicator
 Instrumental method
STANDARD SOLUTIONS
 The standard solutions used in neutralization
titrations :
 STRONG ACIDS (HCl, HClO4, H2SO4)
 STRONG BASES (NaOH, KOH)
o because these substances react more
completely with an analyte
o and Produce sharper endpoints
 Standard solution of acids are prepared by diluting
CONCENTRATED:
 Hydrochloric acid
 Perchloric acid
 Sulfuric acid
 Nitric acid is seldom used because its oxidizing
properties offer the potential for undesirable side
reactions
 Hot concentrated perchloric and sulfuric acids are
potent oxidizing agents and are very hazardous
 Cold dilute solutions of these reagents are safe to
use in analytical laboratory without any precautions
other than eye protection
 Standard solutions of bases are prepared from
SOLID:
 Sodium hydroxide
 Potassium hydroxide
 (occasionally) barium hydroxides
ACID/BASE INDICATORS
 Naturally occurring or synthetic compounds
that exhibit colors depending on the pH of the
solutions in which they are dissolved
 Indicate alkalinity or acidity
 It is a weak organic acid or weak organic base
whose undissociated form differs in color from
its conjugate
 In the dissociation of an acid-type indicator, the
hydronium ion is proportional to the ratio of the
concentration of the acid form to the
concentration of the base form of the indicator,
which in turn controls the color of the solution.
 The human eye is not very sensitive to color
diffrences in a solution containing a mixture of
HIn and In-, especially when the ration is greater
than 10 or smaller than 0.1.
 HIn, exhibits its pure acid color when the ratio is
equal to or greater than 10
 HIn, exhibits its base color when the ratio equal
to or lesser than 0.1
 The range of hydronium ion concentration for
full acid color
[𝐻3 𝑂+ ] = 10𝐾𝑎
 The range of hydronium ion concentration for
full base color
[𝐻3 𝑂 + ] = 0.1𝐾𝑎
 Indicator pH range
𝑝𝐻 (𝑎𝑐𝑖𝑑 𝑐𝑜𝑙𝑜𝑟) = 𝑝𝐾𝑎 + 1
𝑝𝐻 (𝑏𝑎𝑠𝑒 𝑐𝑜𝑙𝑜𝑟) = 𝑝𝐾𝑎 − 1
𝑖𝑛𝑑𝑖𝑐𝑎𝑡𝑜𝑟 𝑝𝐻 𝑟𝑎𝑛𝑔𝑒 = 𝑝𝐾𝑎 ± 1
TITRATION ERRORS WITH ACID/BASE INDICATORS
Two types of errors:
1. Determinate errors
 occurs when the pH at which the indicator
changes differ from the pH at equivalence point
 can be minimized by choosing the indicator
carefully or by making a blank correction
2. Indeterminate error
MKG BSMT-IB
 originates from the limited ability of the human
eye to distinguish reproducibly the intermediate
color of the indicator
 the magnitude of this error depends on:
-change in pH per milliliter of reagent at the
equivalence point
-concentration of the indicator
-Sensitivity of the eye to the two indicator colors
 The visual uncertainty with an acid/base indicator is
in the range of ±0.5 𝑡𝑜 1𝑝𝐻 𝑢𝑛𝑖𝑡
 This uncertainty can be decreased to as little as
±1𝑝𝐻 𝑢𝑛𝑖𝑡 by matching the color of the solution
being titrated with that of a reference standard
containing a similar amount of indicator at
appropriate pH
Variables that influence the behavior of indicators
 Temperature
 Ionic strength of the medium
 Presence of organic solvents and colloidal
particles
- The last two can cause the transition range to
shift by one or more pH units
Common Acid/Base Indicators
 For solutions of strong base, the concentration
of OH- is equal to the analytical concentration of
the bases.
TITRATION OF A STRONG BASE WITH a STRONG
ACID
 Three types of calculations must be done to
construct the hypothetical curve for titrating a
solution of a strong acid with a strong base.
 Stages of titration:
1. Preequivalence
2. Equivalence
3. Postequivalence
 In preequivalence stage: compute the
concentration of the acid from its starting stage
concentration and the amount of base added
 At equivalence point: the hydronium and
hydroxide ions are present in equal
concentrations.
The
hydronium
ion
concentration can be calculated directly from
the ion-product constant for water Kw
 In postequivalence stage: the analytical
concentration of the excess base is computed,
and the hydroxide ion concentration is assumed
to be equal or to a multiple of the analytical
concentration
𝑝𝐾𝑤 = 𝑝𝐻 + 𝑝𝑂𝐻
−𝑙𝑜𝑔10−14 = 𝑝𝐻 + 𝑝𝑂𝐻
Choosing an indicator
The selection of an indicator is not critical when the
reagent concentration is approximately 0.1M
Bromocresol green is unsuited for a titration involving
the 0.001 M reagent because the color change occurs
over a 5ml range well before the equivalence point.
Phenolphtalein is subject to similar objections
TITRATION OF STRONG ACIDS and BASES
 Sources of hydronium ions in aqueous solution
of a strong acid:
 Reaction of the acid with water
 Dissociation of water itself
 For solutions of strong acid that are more
concentrated than about 1x10-6 M, at
equilibrium the concentration of H3O+ is equal to
the analytical concentration of the acid
Only bromothymol blue provides a satisfactory end point
with a minimal systematic error in the titration of
0.001M NaOH
TITRATION OF A STRONG BASE WITH a STRONG ACID
 Titration curves are calculated in a similar way to
those for strong acids
 At preequivalence, the solution is basic
 At equivalence, solution is neutral
MKG BSMT-IB
 At postequivalence the solution becomes acidic
and the hydronium ion concentration is equal to
the analytical concentration of the excess strong
acid
𝑝𝐻 = 𝑝𝐾𝑎
At the half-titration point in a weak-base titration,
[𝑂𝐻 − ] = 𝐾𝑏
𝑝𝑂𝐻 = 𝑝𝐾𝑏
TITRATION CURVES FOR WEAK ACIDS
 Four distinctly different types of calculations are
needed to compute values for a weak acid (or
weak base) titration curve:
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 (up to, but not including, the equivalence
point), 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 concentrations of the weak acid
or base that remains.
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 suppresses 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
Titration curves for strong and weak acids are identical
just slightly beyond the equivalence point. The same is
true for strong and weak bases
The pH at equivalence point in the titration of a weak
acid with a strong base is greater than 7. The solution is
basic. A solution of the salt of a weak acid is always basic
Half-titration points- used to determine the dissociation
constants
At the half-titration point in a weak-acid titration,
[𝐻3 𝑂 + ] = 𝐾𝑎
*recall 𝐾𝑏 = 𝐾𝑤/𝐾𝑎
The effect of concentration
-
-
-
-
In calculating the values for more dilute acids, a
solution of a quadratic equation is necessary for
each point of the curve until after the
equivalence point
In postequivalence point region, the excess OHpredominates, simple calculatins can be used.
The initial pH values are higher and the
equivalence-point pH is lower for the more
dilute solution.
At intermediate titrant volumes the pH values
differ only slightly because of the buffering
action of weak acids that is present in the
regions.
The change in [OH] in the vicinity of the
equivalence point becomes smaller with lower
analyte and reagent concentrations.
Effect of Reaction Completeness
The pH change in the equivalence-point region becomes
smaller as the acid becomes weaker ( reaction between
the acid and the base becomes less complete)
Choosing and indicator
 The choice of indicator is more limited for the
titration of a weak acid
 Bromecresol green is unsuited for titration of
0.1000M of weak acid. Bromothymol blue does
not work either because of its full color change
occurs over a range of titration volume from
about 47 ml to 50 ml of 0.1000M base
 Phenolphtalein, provides a sharp end point with
a minimal titration error.
 Using an indicator with a transition range
between phenolphthalein and of bromothymol
blue in conjuction with a suitable color
MKG BSMT-IB
comparison standard makes it possible to
establish an endpoint.
 When titrating a weak acid use an indicator with
a mostly basic transition range
TITRATION CURVES FOR WEAK BASES
 Calculations needed to draw the titration curve for
weak base are analogous to those of weak acid
 When titrating a weak base use an indicator with a
mostly acidic transition range.
THE COMPOSITION OF SOLUTIONS DURING ACID/BASE
TITRATIONS
The changes in composition during titration can be
visualized by plotting the relative equilibrium
concentration of the weak acid as well as the relative
equilibrium concentration of the conjugate base as
functions of the pH of solution.