1. What is titration? Discuss the importance, principles, and procedure.

Titration is a common technique used to quantitatively determine the unknown concentration of an

identified analyte. It is also called volumetric analysis, as the measurement of volumes is critical in
titration. There are many types of titrations based on the types of reactions they exploit. The most
common types are acid-base titrations and redox titrations.

In a typical titration process, a standard solution of titrant in a burette is gradually applied to react with an
analyte with an unknown concentration in an Erlenmeyer flask. For acid-base titration, a pH indicator is
usually added in the analyte solution to indicate the endpoint of titration. Instead of adding pH indicators,
pH can also be monitored using a pH meter during a titration process and the endpoint is determined
graphically from a pH titration curve. The volume of titrant recorded at the endpoint can be used to
calculate the concentration of the analyte based on the reaction stoichiometry.

Principles:

The determination of acetic acid in vinegar is based on the principle of an acid-base titration method. The
reaction between NaOH and CH3COOH is shown in Equation 1:

CH3COOH(aq) + NaOH(aq) → H2O(l) + NaCH3CO2(aq) (1)

The standardized NaOH solution is progressively added to the vinegar with unknown acetic acid
concentration until the end point is reached. During the acid-base titration, the pH can be plotted as a
function of the volume of the titrant added. The inflection point on the curve, the point at which there is a
stoichiometric equal amount of acid and base in a solution, is called the equivalence point. Most acids
and bases are colorless, with no visible reaction occurring at the equivalence point. To observe when the
equivalence point has been reached, a pH indicator is added. The endpoint is not the equivalence point
but a point at which the pH indicator changes color. It is important to select an appropriate pH indicator so
that the end point is as close to the equivalence point of titration as possible.

At the end point of this reaction, the conjugate base NaCH3CO2 is slightly basic. Phenolphthalein indicator
has a working pH range of 8.3–10.0, which is colorless in acidic solution and magenta above pH 8.2.
Therefore, phenolphthalein is a preferred indicator as it will change from colorless to pink at this condition.
When performing the experiment, it is best to keep the concentration of pH indicator low because pH
indicators themselves are usually weak acids that react with base.

The volume of standardized NaOH solution added at end point can then be used to calculate the molar
concentrations of acetic acid based on the stoichiometry of the above equation. In this experiment, the
titrant NaOH is a strong alkaline and the analyte acetic acid is a weak acid.

Before performing the experiment, it is important to consider the hygroscopic nature of NaOH. This
property requires its solution to be standardized with a stable primary standard such as potassium
hydrogen phthalate (KHC8H4O4). The exact molar concentration of NaOH solution can then be accurately
determined after standardization. The reaction between the primary acid standard and NaOH is shown
in Equation 2:

KHC8H4O4(aq) + NaOH(aq) → H2O(l) + NaKC8H4O4(aq) (2)

A detailed step-by-step titration protocol is presented in the following section.

 PROCEDURE

1. Standardization of NaOH with Potassium Hydrogenphthalate (KHC 8H4O4)

1. To begin, the titrant, sodium hydroxide, must be standardized. Prepare a stock NaOH solution by
dissolving about 4 g of NaOH pellet in 100 mL of deionized water. Note that NaOH is a hazardous
chemical which is corrosive to skin and irritant to eye, be cautious and wear proper personal protection
equipment (PPE) to avoid skin or eye contact.
2. Make a 1:10 dilution of the sodium hydroxide solution by adding 25 mL of the stock sodium hydroxide
solution to a 500-mL bottle. Sodium hydroxide absorbs carbon dioxide. It is important to prevent this by
making sure to use boiled, deionized water, an oven-dried bottle, and to cap the bottle quickly. Make the
solution up to 250 mL with the deionized water and shake to mix.
3. Dry 4–5 g of the primary standard acid, KHC8H4O4 at 110 °C for 4 h in a drying oven and then cool the
solid in a desiccator for 1 h.
4. Dissolve about 4 g of dried KHC8H4O4 in 250 mL of deionized water. Record the mass accurately.
Calculate the molar concentration of the KHC 8H4O4 solution.
5. Pipette 25 mL of KHC8H4O4 into a clean and dry Erlenmeyer flask. Add 2 drops of phenolphthalein, and
swirl gently to mix well. Note that phenolphthalein is toxic and irritant, be cautious to avoid skin or eye
contact.
6. Clean a 50-mL burette and a funnel thoroughly with detergent and water. Flush the burette with water and
rinse 3x with deionized water. Rinse the burette with the diluted NaOH solution 3x, making sure that the
NaOH wets the entire inner surface and drain the waste through the tip. Mount the washed burette on a
ringstand with a clamp and ensure that it stands vertically.
7. Fill the clean burette with the diluted NaOH solution. It should be noted that the amount of the diluted
NaOH needs not be exactly at the zero mark but should be within the scale and sufficient for at least one
titration. Air bubbles may affect the accuracy of volume reading. Carefully check the burette for air
bubbles, and gently tap the burette to free them and open the stopcock to let a few mL of titrant to flow
through and at the same time releasing any trapped air. Read the volume by viewing the bottom of the
meniscus after 10 s. Record this initial volume. Pay attention to the significant figures of the reading.
Record the value to two decimal places in mL.
8. Place the Erlenmeyer flask containing potassium hydrogen phthalate (KHC8H4O4) under the burette and
adjust the height of the burette properly. Titrate the KHC8H4O4 solution by slowly adding NaOH solution in
1–2 mL increments using one hand to control the flow rate by adjusting the stopcock, and the other
swirling the flask.
9. When close to the endpoint, begin adding the titrant drop by drop. The endpoint is reached when the
solution turns a faint, persistent pink color. Record the final volume of the diluted NaOH in the burette.
10. Repeat the titration at least twice more to obtain consistent data. Calculate the molar concentration of the
diluted NaOH solution.

2. Titration of Vinegar with Standardized Sodium Hydroxide Solution

a) The sodium hydroxide solution is now standardized and can be used as a titrant to analyze vinegar. To
reduce the pungent aroma of vinegar, dilute 10 mL of the vinegar solution to be tested in a 1:10 ration to
a total volume of 100 mL.
b) Pipette 25 mL of analyte, to a clean and dry Erlenmeyer flask (noted as VA). Add 2 drops of
phenolphthalein.
c) Fill the burette with the standardized NaOH solution from the first part of the Procedure. Record the initial
volume of titrant (V1).
d) Progressively add the standardized NaOH solution to the vinegar. When the volume of titrant approaches
the expected value, adjust the stopcock to add the titrant drop by drop. Continue to swirl the flask with
one hand and keep the other hand ready to close the stopcock. Once the analyte solution changes to
light pink color, swirl for a few seconds to see whether the color will fade. If the color persists, the titration
reaches the end point. Record the final volume of titrant (V1'). If the solution color fades, add one more
drop of titrant. Wash the bottom tip of the burette using the wash bottle. Collect the washed mixture and
watch the color change of the analyte solution. Continue the titration till the end point. Record the amount
of titrant needed (Vt1 = V1' V1).
e) Repeat the titration at least twice until three concordant values that are within 0.1 mL of one another is
obtained (Vt2 and Vt3).
f) Calculate the mean value of titrant volume using the three values obtained in three different titrations: V t=
(Vt1 + Vt2 + Vt3)/3. The molar concentration of acetic acid in vinegar can be thus calculated using Equation
3.

https://www.jove.com/science-education/5699/introduction-to-titration

2. Define the following terms in titration:

Titrand

- The titrand is the solution whose concentration is not known. The titrand is therefore the analyte
in a chemical titration.

Titrant

- The titrant in a titration is prepared from a standard solution. Standard solutions are solutions
whose concentration is known because they are prepared by weighing precise amounts of the
chemical being dissovled.

Titer

-Titer (or titre) is a way of expressing concentration. Titer testing employs serial dilution to obtain
approximate quantitative information from an analytical procedure that inherently only evaluates as
positive or negative. The titer corresponds to the highest dilution factor that still yields a positive
reading. For example, positive readings in the first 8 serial twofold dilutions translate into a titer
of 1:256 (i.e., 2−8). Titers are sometimes expressed by the denominator only, for example 1:256 is
written 256.[

The term has also two other, conflicting meanings. In titration, the titer is the ratio of actual to nominal
concentration of a titrant, e.g. a titer of 0.5 would require 1/0.5 = 2 times more titrant than nominal.
This is to compensate for possible degradation of the titrant solution. Second, in textile engineering,
titer is a synonym for linear density.
Equivalence point

- The equivalence point, or stoichiometric point, of a chemical reaction is the point at which chemically
equivalent quantities of bases and acids have been mixed. In other words, the moles of acid are equivalent
to the moles of base, according to the equation (this does not necessarily imply a 1:1 molar ratio of
acid:base, merely that the ratio is the same as in the equation). It can be found by means of an indicator, for
example phenolphthalein or methyl orange.

https://en.wikipedia.org/wiki/Equivalence_point

Endpoint

- The endpoint (related to, but not the same as the equivalence point) refers to the point at which the
indicator changes color in a colorimetric titration.

Indicator

- Indicator is any substance that gives a visible sign, usually by a colour change, of the presence or
absence of a threshold concentration of a chemical species, such as an acid or an alkali in a
solution. An example is the substance called methyl yellow, which imparts a yellow colour to an
alkaline solution. If acid is slowly added, the solution remains yellow until all the alkali has been
neutralized, whereupon the colour suddenly changes to red.

https://www.britannica.com/science/chemical-indicator

3. Define and discuss the principles of the different types of titration:

a. Direct Titration
A direct titration is the basic titration method that involves the reaction between the unknown compound and the
compound with known concentration. Here, the addition of excess reagents is not done as in back titrations. The
unknown compound is directly reacted with the known compound. Therefore, the end point of the titration indicates
the end of the reaction. By using that endpoint, the amount of unknown compound present in the sample solution
can be determined.

Most importantly, the end point of the direct titration should be obtained carefully since the endpoint is directly
taken for the further calculations. However, the end point of the direct titration does not often give the
exact equivalence point of the reaction. This is because the endpoint is given when the indicator used in the titration
give its color change. This color change is given a moment after the completion of the reaction. Therefore, it is very
important to determine the exact point where the reaction ends.

Acid-base titrations are good examples for direct titrations. Here, an acid is reacted with a base. An indicator is used
to determine the end point of the reaction since almost all acids and bases are colorless compounds. With the
progression of the reaction, the pH of the solution is changed. At a certain pH, the indicator gives its color change.
The point of the color changes is taken as the endpoint of the reaction. Then, we can determine the concentration of
the unknown (acid or base) according to the stoichiometric relationship between the acid and the base.

b. Indirect Titration
 Indirect titrations , Sometimes it turns out to be impossible to determine the concentration of a
certain substance by titration; for example when such a substance is unstable of not soluble.
In such cases we can execute a so called indirect titration.
The substance to be determined must react beforehand with another reactant (en intermediate or a
substitute). The original substance finally does not participate in the actual titration.
The substitute is titrated and its number of mols calculated.
after that, by calculating, you can determine how much original substance was present.

example:
Marble has a high percentage of Calcium carbonate, and you want to determine that percentage.
An option is to add a substance that reacts with the Calcium carbonate, so to remove all carbonate.
Imagine that you would use Nitric acid to do so; make sure that uyou know exactly how much nictric
acid you added at the start. Part of that acid will react with the carbonate, a rest of it remains.
So, if you know how much acid you added in the beginning, and you know how much acid is left,
then the difference is exactly the amount of acid that reacted with the carbonate.
The acid that remained can be determined by titration with a base, could be: NaOH(aq).

Often we use, in the case of redoxtitrations, the couple: I3/S2O32-(Iodine with tiosulfate).
This titration mostly is applied at indirect titration: the substance to investigate is replaces by Iodine
or Tio, dependent on the question: was the unknown substance an oxydator or a reductor?
At these so called Iodometric titrations we use the cheap and well working indicator: starch(aq).
Fresh starch has spiral shaped molecules (see also module 12, biochemistry), wherein the tri-Iodide
ions exactly fit and cause a dark blue color.
the tri-iodide itself also has a color (yellow brownish), but the (dis)appearance of that color is not as
clear to observe as is the dark blue color

If you want to execute an acid base titration without any indicator, then you could use a pH meter.
On the meter you must continuously check the (changing) pH values: In the equivalence point, the
pH change is at its maximum; a sudden big pH change shows up.
You can plot the course of the pH change in a graph (by hand or automatically).

Every time you add, for example 0.5 ml titrant, you read the pH value, or the potential (in the case
of redox reaction) and you put that value on paper. That's how you create a titration curve. In such a
curve you can easily derive the equivalence point.

c. Back Titration

A back titration is a titration method used to determine the concentration of an unknown using an excess
amount of a compound with a known concentration. There is a chemical reaction between these
compounds. Since the amount of the compound with a known concentration added is known already, we
can determine the amount of the compound that has reacted with the unknown compound by doing a back
titration.

In a titration, there are two components involved: a titrant and a titrand. The titrant is the solution with a
known concentration. Titrand is the analyte or the sample. This sample is composed of a compound with
an unknown concentration and this compound should react with the titrant solution. A back titration does
not involve a direct reaction between the titrant and the titrand. First, we add a compound in excess to the
sample solution which can cause a chemical reaction. Then we measure the amount of remaining
compound. Therefore, the titrand here is also a known compound.
Example
Let us consider an example in order to understand this concept. We are given a metal ion solution with an
unknown metal ion having an unknown concentration. We can use EDTA back titration method, a
common titration method, for the analysis of this solution. Here, an excess amount of EDTA should be
first added to the sample solution. The concentration of the EDTA solution should be determined later
using a primary standard. The addition of EDTA will cause the formation of metal ion-EDTA complex.
Then the remaining amount of EDTA present in the sample is determined using a Mg+2 solution in the
presence of EBT indicator. Metal ions always form complexes with EDTA in 1: 1 ratio. Since the
previously added amount of EDTA is known, we can calculate the amount of EDTA that reacted with
unknown metal.
http://pediaa.com/difference-between-back-titration-and-direct-titration/

http://www.daanvanalten.nl/chemistry/trefwoorden/indirectetitratie_140504.html

4. Define and discuss the principles of the different titrimetric methods

a. Acid-base titration-An acid–base titration is a method of quantitative

analysis for determining the concentration of an acid or base by exactly neutralizing it
with a standard solution of base or acid having known concentration. A pH
indicator is used to monitor the progress of the acid–base reaction. If the acid
dissociation constant (pKa) of the acid or base dissociation constant (PKb) of base in
the analyte solution is known, its solution concentration (molarity) can be
determined.
https://en.wikipedia.org/wiki/Acid%E2%80%93base_titration
b. Redox titration- A redox titration is a type of titration based on a redox reaction
between the analyte and titrant. Redox titration may involve the use of a redox
indicator and/or a potentiometer. A common example of a redox titration is treating a
solution of iodine with a reducing agent to produce iodide using a starch indicator to
help detect the endpoint. Iodine (I2) can be reduced to iodide (I−) by
e.g. thiosulfate (S2O32−), and when all iodine is spent the blue colour disappears.

https://en.wikipedia.org/wiki/Redox_titration
 c. Complexometric titration- Complexometric titrations are particularly useful for
the determination of a mixture of different metal ions in solution. An indicator capable
of producing an unambiguous color change is usually used to detect the end-point of
the titration. This method represents the analytical application of a complexation
reaction. In this method, a simple ion is transformed into a complex ion and the
equivalence point is determined by using metal indicators or electrometrically.
Various other names such as chilometric titrations, chilometry, chilatometric titrations
and EDTA titrations have been used to describe this method.
https://www.scribd.com/doc/25357952/Types-of-Titration
d. Precipitation titration- Precipitation titrations are a form of titration useful in
the determination of halides such as chlorides, bromides and iodides. These
titrations involve the use of a precipitating agent such as silver nitrate, and are
therefore also known as argentimetric titrations. Depending on the method of
detecting the end point of the titration, there are three methods in precipitation
titrations: Mohr’s method, Volhard’s method and Fajan’s method.

https://sciencing.com/precipitation-titration-techniques-8665193.html

e. Zeta Potential Titration- is a titration of heterogeneous systems, for

example colloids and emulsions. Solids in such systems have very high surface area.
This type of titration is used to study the zeta potential of these surfaces under
different conditions.Is a titration of heterogeneous systems, for
example colloids and emulsions. Solids in such systems have very high surface
area. This type of titration is used to study the zeta potential of
these surfaces under different conditions.

https://wikivisually.com/wiki/Zeta_potential_titration

f. Miscellaneous Titration- a form of titration can also be used to determine the

concentration of a virus or bacterium. The original sample is diluted until the last
dilution does not give a positive test for the presence of virus. This value, the
titre, may be used on TCID50, EID50, ELD50, LD50 or pfu. This procedure is
more commonly known as assay.
https://www.scribd.com/doc/25357952/Types-of-Titration

5.