

Titration of acids and bases

Experiment 16: Titrations of Acids and Bases Titration of Acids and Bases EXPERIMENT
To become familiar with the techniques of titration, a volumetric method of analysis; to determine the amount of acid in an unknown. APPARATUS AND CHEMICALS:
One of the most common and familiar reactions in chemistry is the reaction of an acid with a base. This reaction is termed neutralization, and the essential feature of
this process in aqueous solution is the combination of hydronium ions with hydroxide
ions to form water.
In this experiment you will use this reaction to determine accurately the concentration of a sodium hydroxide solution that you have prepared. The process of
determining the concentration of a solution is called standardization. Next you will
measure the amount of acid in an unknown. To do this, you will accurately measure,
with a buret, the volume of your standard base that is required to exactly neutralize the
acid present in the unknown. The technique of accurately measuring the volume of a
solution required to react with another reagent is termed titration.
An indicator solution is used to determine when an acid has exactly neutralized a
base, or vice versa. A suitable indicator changes colors when equivalent amounts of
acid and base are present. The color change is termed the end point of the titration.
Indicators change colors at different pH values. Phenolphthalein, for example, changes
color from colorless to pink at a pH of about 9; in slightly more acidic solutions it is
colorless, whereas, in more alkaline solutions it is pink.
The point at which stoichiometrically equivalent quantities are brought together is known as the equivalence point of the titration. It should be noted that the
equivalence point in a titration is a theoretical point.
Experiment 16: Titrations of Acids and Bases The most common way of quantifying concentrations is molarity (symbol M) that is defined as the number of moles of solute per liter of solution, or the number of millimoles of solute per milliliter of solution: From Equation [1] the moles of solute (or mmol solute) are related to the molarity and the volume of the solution as follows: M x Liters = moles solute and M x mL = mmol solute 1. PH concept
The concentration of H+ in natural and biological environments, as well as in all chemically, oriented problems, is highly important. where brackets denote concentrations in mole/L. As it is seen from the value of Kw, ion product of pure water, [H+] = [OH-] = 1.0 x 10-7M. If an acid or a base other than
H2O present. [H+] deviates from the value of 1.0 x 10-7M; however, [H+]x[OH-] always
equals to Kw at constant temperature. The concentration of [H+] may change over a
wide range of values and these values are frequently expressed in terms of exponential
numbers. For this reason, a simpler form of representation for [H+] is provided as
follows: pH = -Iog[H+]
What are the pH values for a) 0. 1 M HCl and b) 0. 1 M NaOH?
a) HCI is a strong acid, it dissociates almost completely in aqueous solution. Therefore 0.1 M HCl
gives 0.1 M H+ and 0. 1 M Cl-.
b) NaOH is a strong base; it dissociates almost completely in aqueous solution. 0.1 M NaOH gives
0.1 M Na+ and 0. 1 M OH-.
[H+] = Kw/ [OH-] = 1 x 10-14 /0.1 = 1.0 x 10-13 M. EXAMPLE 2:
What is the pH of 0.1 M formic acid, HCOOH, solution? Experiment 16: Titrations of Acids and Bases SOLUTION :
Formic acid, HCOOH, is a weak acid. It does not dissociate completely in aqueous solution. Its dissociation depends on the following relations, H xHCOO Ka is the dissociation constant for formic acid. At equilibrium, x = HCOOH molecules dissociated mmol/L. If x moles of HCOOH dissociated, it would give x moles of H+ and x moles of HCOO- ions. If the initial concentration of formic acid is 0.1 M at equilibrium, it remains 0.1-x M. x in the denominator can be neglected, since 0.1 >>2.0x10-4 2. Indicators
Acid-base indicators, in general, are organic molecules that are weak acids or weak bases. They exhibit different colors in their protonated form, HIn, and unprotonated form, In-. For an indicator of weak acid type, the ratio of [In-] to [HIn] depends on [H+]. H xIn where Ki is the dissociation constant for indicator HIn. When HIn is the predominant species in the solution, the acid-color of the indicator is observed. On the other hand, the base-color is observed when the predominant species is In-. In general, depending on the sight ability of human eye the acid or base-color can be observed when the following conditions are met. log ([In-] MIn]) = - 1 acid-color is observed log ( [In-] / [HIn] = 1 base-color is observed Experiment 16: Titrations of Acids and Bases Every indicator has a specific color transition range, in terms of pH units, depending on its Ki value. In practice, this range may equal to or less than 2 pH units. EXAMPLE 3:
Phenolphthalein, an acid-base indicator, has a color transition range of pH = 8.0 - 10.0. Its acid-color is colorless, base-color is pink. If one drop of phenolphthalein solution is added to 5x 10-3 M NaOH solution, what would the color of this solution become? SOLUTION:
11.70 >10.0 base-color, solution becomes pink 3. Strong acid - base titration
The reaction between a strong acid (e.g. HCl) and a strong base (e.g. NaOH) can be HC1(aq) + NaOH(aq)  NaCl(aq) + H2O(l) Thus for a monoprotonic acid and base at the end point VacidMacid = VbaseMbase
A standard solution is the one whose concentration is accurately known. In this experiment a sodium hydroxide solution of unknown concentration will be standardized against a standard hydrochloric acid solution. The standardized NaOH solution will be used to determine the molecular weight of a monoprotonic acid by titrating a known weight of the acid against the NaOH solution. Aspirin is a monoprotonic acid C6H4(OCOCH3)COOH, acetylsalicylic acid, Ka=1.05x10-3, with the structure shown below. Aspirin tablets contain acetylsalicylic acid and inert "fillers". From its chemical formula you can calculate the weight of acid actually present in the sample, hence the percentage of acetylsalicylic acid in the Experiment 16: Titrations of Acids and Bases aspirin. The pH of the solution can then be calculated using dissociation constant of the acetylsalicylic acid. In this experiment we use an acid-base indicator, phenolphthalein, to signal the end
point in the titration. We choose an indicator such that its color change occurs as
closely as possible to the equivalence point.
Before beginning this experiment in the laboratory. you should be able to answer the following questions: 1. Define standardization and state how you would go about doing it. 5. What are equivalence points and end points and how do they differ? 8. What is the molarity of a solution that contains 1.89 g of H 9. If 50.0 mL of NaOH solution is required to react completely with 1.24 g KHP (monoprotic acid), what is the molarity of the NaOH solution? 10. In the titration of an impure sample of KHP, it was found that 29.4 mL of 0.100 M NaOH was required to react completely with 0.745 g of sample. What is the percentage of KHP in this sample? PROCEDURE:
Preparation of a Buret for Use: Clean a 50-mL, buret with soap solution and a buret
brush and thoroughly rinse with tap water. Then rinse with at least five 10-mL portions
of distilled water. The water must run freely from the buret without leaving any drops
adhering to the sides. Make sure that the buret does not leak and that the stopcock turns
Reading a Buret: All liquids, when placed in a buret, form a curved meniscus at their
upper surfaces. In the case of water or water solutions, this meniscus is concave and the
most accurate buret readings are obtained by observing the position of the lowest point
on the meniscus on the graduated scales. To avoid parallax errors when taking readings,
the eye must be on a level with the meniscus.
A. Standardization of Sodium Hydroxide Solution
1. Completely fill the buret with the NaOH solution and remove the air from the tip by running out some of the liquid into an empty beaker. Make sure that the lower part of the meniscus is at the zero mark or slightly lower. Allow the buret to stand for at least 30 s before reading the exact position of the meniscus. Remove any hanging drop from the buret tip by touching it to the side of the beaker used Experiment 16: Titrations of Acids and Bases for the washings. (Caution solutions are corrosive. In case of accidental skin contact wash with plenty of water.) Record the initial buret reading. 2. Withdraw 15.00 mL of standard HCl solution into a clean Erlenmeyer flask and add a few drops of phenolphthalein solution. 3. Slowly add the sodium hydroxide solution to one of your flasks of HCl solution while gently swirling the contents of the flask, as illustrated in Figure 16.1. As the sodium hydroxide solution is added, a pink color appears where the drops of the base come in contact with the solution. This coloration disappears with swirling. As the end point is approached, the color disappears more slowly, at which time the sodium hydroxide should be added drop by drop. It is most important that the flask be swirled constantly throughout the entire titration. The end point is reached when one drop of the sodium hydroxide solution turns the entire solution in the flask from colorless to pink. The solution should remain pink when it is swirled. Allow the titrated solution to stand for at least 1 min so the buret will drain properly. Remove any hanging drop from the buret tip by touching it to the side of the flask and wash down the sides of the flask with water from the wash bottle. 4. Repeat this procedure with the other two samples. 5. From the data you obtain in the three titrations, calculate the molarity of the sodium hydroxide solution to four significant figures. 6. The three determinations should agree within 1.0 percent. If they do not, the standardization should be repeated until agreement is reached. The average of the three acceptable determinations is taken as the molarity of the sodium hydroxide. Calculate the standard deviation of your results. Save your standardized solution for the unknown determination. B. Molecular Weight of an Unknown Acid
1. As carefully as possible, weigh two 1 g samples of unknown acid. 2. Place the sample in a clean Erlenmeyer flask, add 50 mL of deionized water and shake until the solid has all dissolved. Add two drops of phenolphthalein solution. 3. Titrate with the NaOH solution until the pink color of the indicator persists for at least one minute. Record the volume of the NaOH solution used. 4. Titrate the second sample of the unknown acid against the NaOH solution, as C. Analysis of an Unknown Acid
1. Obtain an aspirin tablet from your assistant and weigh it. Experiment 16: Titrations of Acids and Bases 2. Dissolve the tablet in 100 mL of water, add a few drops of phenolphthalein solution and titrate against the standard NaOH solution. (If the tablet does not dissolve completely, begin the titration with shaking. The solid will dissolve as the titration is performed. It will be necessary to shake vigorously between additions of NaOH solution.) Flush all waste solutions down the sink drain with plenty of water. QUESTIONS:
1. Write the balanced chemical equation for the reaction of HCl with NaOH. 2. A solution of malonic acid, H2C3H2O4, was standardized by titration with 0.100 M NaOH solution. If 21.82 mL of the NaOH solution were required to neutralize completely 12.12 mL of the malonic acid solution, what is the molarity of the malonic acid solution? 3. Sodium carbonate is a reagent that may be used to standardize acids in the same way. In such standardization it was found that a 0.432-g sample of sodium carbonate required 22.3 mL of a sulfuric acid solution to reach the end point for the reaction. Na2CO3(aq) + H2SO4(aq)  H2O(l) + CO2(g) + Na2SO4(aq) 4. A solution contains 0.252 g of oxalic acid, H Experiment 16: Titrations of Acids and Bases Titration of Acids and Bases
A. Standardization of Sodium Hydroxide Solution
1. Molarity of HCI solution : …………… 2. Volume of NaOH solution: …………… 3. Volume of HCI solution used for the titrations 1st: …….2nd: …….3rd: ……. 4. Molarity of NaOH solution: ……………. B. Molecular Weight of an Unknown Acid
1. Weight of unknown acid samples: 1st: ……… 2nd: ………… 2. Volumeof NaOH solution used for the titrations: 1st:……….2nd: ………… 3. Molecular weight of the unknown acid: ………………. 4. M.Wt. of the unknown acid is 1st sample:.…. 2nd sample: ……. average: …… C. Analysis of an Unknown Acid
3. Volume of NaOH solution used for the titration : ……………. 4. Molarity of acetylsalicylic acid : ……………. 5. Weight of acetylsalicylic acid: ………………. 6. pH of acetylsalicylic acid: ……………………


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