Neutralization Titration

1. 1

2. 1. Hydrolysis of salts,
2. Relative strength and its effect on titration,
3. Law of mass action
4. Ionic product of Water
5. Common ion effect,
6. Henderson- hesselbach equation,
7. Ph & buffers,
8. Neutralization curve,
9. Acid-base indicators,
10.Theory of indicators,
2

3. • An analytical technique based on measuring the amount of a
reagent of known concentration that is consumed by an
analyte in a chemical or electrochemical reaction.
• Volumetric titrations involve measuring the volume of a solution of
known concentration that is needed to react completely with the analyte.
• In Gravimetric titrations, the mass of the reagent is measured
instead of its volume.
• In coulometric titrations, the “reagent” is a constant direct
electrical current of known magnitude that consumes the analyte.
3

4. • Water is the most plentiful solvent on Earth, is easily purified, and is not
toxic. It is, therefore, widely used as a medium for chemical analyses.
• Most of the solutes we will discuss are electrolytes, which form ions
when dissolved in water (or certain other solvents) and thus produce
solutions that conduct electricity.
• Strong electrolytes ionize essentially completely in a solvent, but
weak electrolytes ionize only partially.
Classification of Electrolytes
Strong Electrolytes Weak Electrolytes
1. Inorganic acids such as HNO3,
HClO4, H2SO4*, HCl, HI, HBr,
HClO3, HBrO3
2. Alkali and alkaline-earth hydroxides
3. Most salts
1. Many inorganic acids, including
H2CO3, H3BO3, H3PO4, H2S,
H2SO3
2. Most organic acids
3. Ammonia and most organic bases
4. Halides, cyanides, and thiocyanates
of Hg, Zn, and Cd
4

5. • An acid donates protons. A base accepts protons.
• A conjugate base is formed when an
• acid loses a proton. A conjugate acid is formed when a base
accepts a proton.
• For example, acetate ion is the conjugate base of acetic acid.
Similarly, ammonium ion is the conjugate acid of the base
ammonia.
• Many solvents are proton donors or proton acceptors and can thus
induce basic or acidic behavior in solutes dissolved in them.
• For example, in an aqueous solution of ammonia, water can donate
a proton and acts as an acid with respect to the solute
• Similarly, for Nitric acid
5

6. • The conjugate acid of water is the hydrated proton written as H3O+. This species
is called the hydronium ion, and it consists of a proton covalently bonded
to a single water molecule. Higher hydrates such as H5O2,
+ H9O4
+, and the
dodecahedral cage structure shown in Figure may also appear in aqueous
solutions of protons. For convenience, however, we generally use the
notation H3O+, or more simply H1, when we write chemical equations
containing the hydrated proton.
6

7. • Amphiprotic Species
• Species that have both acidic and basic properties are amphiprotic.
• An example is dihydrogen phosphate ion, H2PO4 -, which behaves as a
base in the presence of a proton donor such as H3O+.
• The simple amino acids are an important class of amphiprotic compounds
that contain both a weak acid and a weak base functional group. When
dissolved in water, an amino acid, such as glycine, undergoes a kind of
internal acid/base reaction to produce a zwitterion—a species that has
both a positive and a negative charge.
• Water is the classic example of an amphiprotic solvent, that is, a solvent
that can act either as an acid or as a base depending on the solute. Other
common amphiprotic solvents are methanol, ethanol, and anhydrous acetic
acid.
7

8. • Amphiprotic solvents undergo self-ionization, or autoprotolysis, to form a
pair of ionic species.
• Autoprotolysis is yet another example of acid/base behavior.
• The extent to which water undergoes autoprotolysis at room temperature is
slight.
• Thus, the hydronium and hydroxide ion concentrations in pure water are
only about 1027 M. Despite the small values of these concentrations, this
dissociation reaction is of utmost importance in understanding the behavior
of aqueous solutions.
8

9. • For strong acids the reaction with the solvent is sufficiently complete that
no undissociated solute molecules are left in aqueous solution.
• The rest are weak acids, which react incompletely with water to give solutions
containing significant quantities of both the parent acid and its conjugate base.
• the weakest acid forms the strongest conjugate base, that is, ammonia has a
much stronger affinity for protons than any base above it.
• The tendency of a solvent to accept or donate protons determines the strength
of a solute acid or base dissolved in it.
9

10. • The tendency of a solvent to accept or donate protons determines the strength of
a solute acid or base dissolved in it.
• For example, perchloric and hydrochloric acids are strong acids in water. If
anhydrous acetic acid, a weaker proton acceptor than water, is substituted as the
solvent, neither of these acids undergoes complete dissociation.
• Instead, equilibria such as the following are established:
• Perchloric acid is, however, about 5000 times stronger than hydrochloric acid in
this solvent. Acetic acid thus acts as a differentiating solvent toward the two
acids by revealing the inherent differences in their acidities.
• Water, on the other hand, is a leveling solvent for perchloric, hydrochloric,
and nitric acids because all three are completely ionized in this solvent and
show no differences in strength. There are differentiating and leveling solvents for
bases as well.
• The common strong bases include NaOH, KOH, Ba(OH)2, and the quaternary
ammonium hydroxide R4NOH, where R is an alkyl group such as CH3 or C2H5.
• The common strong acids include HCl, HBr, HI, HClO4, HNO3, the first proton in
H2SO4, and the organic sulfonic acid RSO3H.
10

11. 11
• The concept of a dynamic equilibrium is central to quantitative discussion of most
chemical phenomena.
• A dynamic equilibrium is a state in which there appears to be nothing happening,
a state in which there is no net change, but is also a state in which chemical
reactions are taking place, often at rapid rates.
• Consider the chemical reaction
• We can follow the rate of this reaction and the extent to which it proceeds to the
right by monitoring the appearance of the orange-red color of the triiodide ion I3-.
(The other participants in the reaction are colorless.) For example, if 1 mmol of
arsenic acid, H3AsO4, is added to 100 mL of a solution containing 3 mmol of
potassium iodide, the red color of the triiodide ion appears almost immediately.
Within a few seconds, the intensity of the color becomes constant, showing that
the triiodide concentration has become constant.

12. • These reactions, however, are opposites of each other so that the net change is
no change at all. Any chemical system at equilibrium is a dynamic system in
which real reactions are occurring with real speed. And, as we shall see in a
different section, if that system should be altered in some way by external means
the reactions will operate so as to bring it back to equilibrium again.
• Le Châtelier’s principle. This principle states that the position of chemical
equilibrium always shifts in a direction that tends to relieve the effect of an
applied stress.
• For example, an increase in temperature of a system alters the concentration
relationship in the direction that tends to absorb heat, and an increase in
pressure favors those participants that occupy a smaller total volume.
• What will be the effect of adding carbon monoxide to the water gas system at
equilibrium?
• If ammonia is added to a system of ammonia and its elements at
• equilibrium, which way will the equilibrium shift? (b) What will happen
• to the temperature of the system because of this shift?
12

13. https://www.youtube.com/watch?v=xfGlEXWDRZE
• The influence of concentration or pressure (if the participants are gases) on
the position of a chemical equilibrium can be described in quantitative terms
by means of an equilibrium-constant expression.
• However Equilibrium-constant expressions provide no information about
whether a chemical reaction is fast enough to be useful in an analytical
procedure.
• Cato Guldberg (1836–1902) and Peter Waage (1833–1900) proposed this
Law, According to law ,
• The rate of chemical reaction is proportional to the active masses of the
reacting sabstances.
• Let us consider a reaction,
• For this reaction Equilibrium-constant
13

14. • The value of K depends only on the specific chemical equation and on the
temperature. It does not depend on any of the other factors that can affect the rate
of a reaction.
• For example, if different quantities of the same reactants and products are
introduced into different reaction vessels, they will react with one another until, at
equilibrium, the same ratio of concentrations, each raised to the appropriate
power, is established.
• The value of K for a reaction gives quantitative information about the extent of the
reaction.
• That is, a large value of K (usually about 104 or larger) means that, at equilibrium,
the reaction proceeds almost completely to the right;
• a small value of K (usually less than 10-4 or so) means that, at equilibrium, the
reactants have not reacted much at all (or the “products” have reacted almost
completely).
14

15. • Aqueous solutions contain small concentrations of hydronium and hydroxide ions
as a result of the dissociation reaction
• An equilibrium constant for this reaction can be written as :
• The concentration of water in dilute aqueous solutions is enormous, however,
when compared with the concentration of hydronium and hydroxide ions.
• As a result, [H2O]2 in above Equation can be taken as constant, and we write:
• where the new constant Kw is given a special name, the ion-product constant
for water.
• At 25°C, the ion-product constant for water is 1.008 * 10-14. For convenience, we
use the approximation that at room temperature Kw = 1.00 * 10-14
15

16. Why [H2O] Does Not Appear in Equilibrium-Constant Expressions
for Aqueous Solutions ?
• In a dilute aqueous solution, the molar concentration of water is
• Suppose we have 0.1 mol of HCl in 1 L of water. The presence of this acid will shift
the equilibrium shown in above Equation to the left.
• Originally, however, there was only 10-7 mol/L OH- to consume the added protons.
Therefore, even if all the OH- ions are converted to H2O, the water concentration
will increase to only
16

17. • The common-ion effect is a mass-action effect predicted from Le Châtelier’s
principle.
• The solubility of an ionic precipitate decreases when a soluble compound
containing one of the ions of the precipitate is added to the solution. This behavior
is called the common-ion effect.
• Most, but not all, sparingly soluble salts are essentially completely dissociated in
saturated aqueous solution.
• For example, when an excess of barium iodate is equilibrated with water, the
dissociation process is adequately described by the equation
• If we add Ba(NO3)2 ,which is a strong electrolyte, in above solution as [Ba2+]
increase and according to Le Châtelier’s principle solubility of Ba(IO3)2 will
decrease so as to relieve the stress applied to system.
17

18. 18
• When a weak acid or a weak base is dissolved in water, partial dissociation
occurs.
• Because of this reactions of Bronsted acids and bases with water are equilibrium
reactions, we can write equilibrium constant expressions for these ionizations or
dissociations.
• For example, the dissociation of nitrous acid in water and its equilibrium
constant expression are as follows:
• The constant is called the acid dissociation constant.
• he subscript “a” indicates that the constant is for the ionization of an acid.
• Similarly, an equilibrium constant expression can be written for the reaction of a
weak base with water
• The dissociation constant have very small numerical values hence it is more
convinient to use logarithmic form of these values.
pk = - log K

19. • Notice that [H2O] does not appear in the denominator of either equation because
the concentration of water is so large relative to the concentration of the weak acid
or base that the dissociation does not alter [H2O] appreciably.
• Just as in the derivation of the ion-product constant for water, [H2O] is incorporated
into the equilibrium constants Ka and Kb.
Dissociation Constants for Conjugate Acid/Base Pairs
• Consider the base dissociation-constant expression for ammonia and the acid
dissociation constant expression for its conjugate acid, ammonium ion:
• By multiplying one equilibrium-constant expression by the other, we have
• There for,
• This relationship is general for all conjugate acid/base pairs. Many compilations of
equilibrium-constant data list only acid dissociation constants because it is so easy
to calculate dissociation constants for bases by using Equation.
19

20. • When the weak acid HA is dissolved in water, two equilibria produce hydronium ions:
• Normally, the hydronium ions produced from the first reaction suppress the
dissociation of water to such an extent that the contribution of hydronium ions from
the second equilibrium is negligible. Under these circumstances, one H3O+ ion is
formed for each A- ion, and we write
• Furthermore, the sum of the molar concentrations of the weak acid and its conjugate
base must equal the analytical concentration of the acid cHA because the solution
contains no other source of A- ions. Therefore,
• Substituting [H3O+ ] for [A- ] yields
20

21. • When [A- ] and [HA] are replaced by their equivalent, the equilibrium-
constant expression becomes
• This equation can frequently be simplified by making the additional
assumption that dissociation does not appreciably decrease the molar
concentration of HA.
• Thus, if [H3O+] << cHA, cHA - [H3O+] ͌ cHA, and Equation reduces to
• And
• The error introduced by the assumption that [H3O+] << cHA , increases as the
molar concentration of acid becomes smaller and its dissociation constant
becomes larger.
21

22. • We can adapt the techniques of the previous sections to calculate the hydroxide or
hydronium ion concentration in solutions of weak bases.
• Aqueous ammonia is basic as a result of the reaction
• The predominant species in this solution is certainly NH3. Nevertheless, solutions of
ammonia are still called ammonium hydroxide occasionally because at one time
chemists thought that NH4OH rather than NH3 was the undissociated form of the
base.
• We write the equilibrium constant for the reaction as
• Upon Successive Approximations just like previous derivation we can get
[OH-] = K b cNH3
22

23. • A buffer solution resists changes in pH when it is diluted or when acids or
bases are added to it.
• Generally, buffer solutions are prepared from a conjugate acid/base pair, such as
acetic acid/sodium acetate or ammonium chloride/ammonia.
• Scientists and technologists in most areas of science and in many industries use
buffers to maintain the pH of solutions at a relatively constant and predetermined
level.
• A solution containing a weak acid, HA, and its conjugate base, A-, may be acidic,
• neutral, or basic, depending on the positions of two competitive equilibria:
• If the first equilibrium lies farther to the right than the second, the solution is acidic.
• If the second equilibrium is more favorable, the solution is basic.
• These two equilibrium constant expressions show that the relative concentrations of
the hydronium and hydroxide ions depend not only on the magnitudes of Ka and Kb
but also on the ratio between the concentrations of the acid and its conjugate base.
23

24. • To find the pH of a solution containing both an acid, HA, and its conjugate base,
NaA, we need to express the equilibrium concentrations of HA and NaA in terms of
their analytical concentrations, cHA and cNaA.
• If we look closely at the two equilibria, we find that the first reaction decreases the
concentration of HA by an amount equal to [H3O+ ], while the second increases the
HA concentration by an amount equal to [OH-].
• Thus, the species concentration of HA is related to its analytical concentration by the
equation
• Similarly, the first equilibrium will increase the concentration of A- by an amount
equal to [H3O+ ], and the second will decrease this concentration by the amount
[OH-].
• Therefore, the equilibrium concentration is given by a second equation that looks a
lot like
• Because of the inverse relationship between [H3O+ ] and [OH-], it is always possible
to eliminate one or the other from these equations. Additionally, the difference in
concentration between [H3O+ ],and [OH-] is usually so small relative to the molar
concentrations of acid and conjugate base that Equations simplify to
24

25. • If we then substitute these values into the dissociation-constant expression and
rearrange the result, we have
• The assumption sometimes breaks down with acids or bases that have
dissociation constants greater than about 10-3 or when the molar concentration of
either the acid or its conjugate base (or both) is very small. In these
circumstances, either [OH-] or [H3O+] must be retained in Equations depending
on whether the solution is acidic or basic.
• Within the limits imposed by the assumptions made in deriving final equations, it
says that the hydronium ion concentration of a solution containing a weak acid
and its conjugate base depends only on the ratio of the molar concentrations of
these two solutes. Furthermore, this ratio is independent of dilution because the
concentration of each component changes proportionally when the volume
changes.
• So, we can rewrite it as
• This is Henderson-Hasselbalch equation, which is used to calculate the pH of
buffer solutions, is frequently encountered in the biological literature and
biochemical texts
25

26. • An indicator is a substance which is used to determine the end point in a
titration. In acid-base titrations, organic substances (weak acids or weak bases)
are generally used as indicators. They change their colour within a certain pH
range. The colour change and the pH range of some common indicators are
tabulated below
• Ostwald’s theory (Acid-Base concept)
• Quinoid ( Resonsnce ) theory
26

27. 27
Methyle red
Just before end point
Correct End point
End point exceeded

28. 28

29. Methyle red
29

30. Thymol blue
30

31. • A titration curve is the plot of the pH of the analyte solution versus the
volume of the titrant added as the titration progresses.
1. Titration of a strong acid with a strong base
2. Titration of a weak acid with a strong base
3. Titration of a strong acid with a weak base
4. Titration of a weak base with a weak acid
31

32. 32

33. 33

34. 34

35. 35

36. 36

37. 37

38. • The halide ions -chloride, bromide and iodide - are too weakly
basic to react quantitatively with acetous perchloric acid.
• Addition of mercuric acetate (which is undissociated in acetic
acid solution) to a halide salt replaces the halide ion by an
equivalent quantity of acetate ion, which is a strong base in
acetic acid.
 2R.NH2.HCl ⇌ 2RNH3
+ + 2Cl−
 (CH3COO)2Hg + 2Cl− → HgCl2 + 2CH3COO−
(undissociated) (undissociated)
 2CH3COOH2
+ + 2CH3COO− ⇌ 4CH3COOH
38

39. 39