CHAPTER 1 Measurement of pH, preparation of

CHAPTER 1 Measurement of pH, preparation of reagents, standard acids, alkali and buffers Suresh Kumar Division of Biochemistry, ICAR-Indian Agricultural Research Institute, New Delhi-110012, India. Introduction The pH of a solution is a measure of the molar concentration of hydrogen ions [H+] in the solution, and it is a measure of the acidity or basicity of the solution. The symbol pH stands for "power of hydrogen" and numerical value for pH is just the negative of the power of 10 of the molar concentration of hydrogen ions (H+). The ionic product of water, Kw, is the basis for the pH scale. The total hydrogen ion concentration from all sources is experimentally measurable and is expressed as the pH of the solution. The pH scale designates the concentration of H+ (and thus of OH) in any aqueous solution in the range between 1.0 M H+ and 1.0 M OH. The term pH is defined by the expression

where the symbol p denotes “negative logarithm of” Thus. pH can be defined as the –ve logarithm of the H+ concentration. For a neutral solution at 25C, in which the concentration of hydrogen ions is 1.0x10-7 M, the pH can be calculated as follows:

Thus, the ionic product of water makes it possible to calculate the concentration of H , with given concentration of OH, and vice versa. The value of 7 for pH of a neutral solution is not an arbitrarily chosen figure; it is derived from the absolute value of the ion product of water. A solution having pH greater than 7 is alkaline or basic, and the concentration of OH- is greater than that of H+. A solution having a pH less than 7 is acidic. Note that the pH scale is logarithmic, not arithmetic. For example: Two solutions differ in pH by 1 unit, means that one solution has 10 times the H+ concentration of the other; but it does not tell us the absolute magnitude of the difference. In other words, for every 10‐fold change in concentration of the H+ ion (e.g. 0.1 to 1.0), the pH changes by 1 unit. The pH of an aqueous solution can be approximately measured using various indicator dyes, including litmus, phenolphthalein, and phenol red, which undergo color changes as a proton dissociates from the dye molecule. Accurate determinations of pH in the laboratory can be made with a glass electrode that is selectively sensitive to H+ concentration but insensitive to Na+, K+, and other cations. In a pH meter, the signal from an electrode is amplified and compared with the signal generated by a solution of accurately known pH. Measurement of pH is one of the most important and frequently used procedures in biochemistry. The pH affects the structure and activity of biological macromolecules; such as +

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the catalytic activity of enzymes. Measurements of the pH of blood and urine are commonly used in medical diagnoses. Nearly everything around is an acid or a base, with the exception of water. Water is neither acidic nor basic, rather it is neutral. The pH scale was developed to measure how acidic or basic a substance is. This scale measures pH from 0  14. Acids have pH values between 0 and 7, and bases have pH values between 7 and 14. Pure water is neutral and has a pH of exactly 7. The pH of an aqueous solution can be measured in a variety of ways. The most common way used now-adays is using a pH‐sensitive glass electrode in a pH meter. However, alternative methods for measuring the pH of a solution include use of (i) litmus or pH papers, (ii) Indicator dyes, and (iii) a pH meter. Litmus is an extract from specific lichen which is formed into a powder and then used in a chemical solution to treat paper, cut into strips, and thus litmus paper is made. When litmus paper is dipped into a solution to check pH, it changes colour depending on the pH of the solution. The colour of wetted litmus paper is matched to a colour on a colour chart to infer a pH value. The pH paper is available commercially for narrow pH ranges (e.g., 3.0 to 5.5 pH, 4.5 to 7.5 pH and 6.0 to 8.0 pH), and fairly wide pH ranges of 1.0 to 11.0 pH. Also, three different types of litmus papers (viz. blue, red and neutral) are available in the market. Blue litmus paper turns red in acidic solutions, and remains blue in alkaline solutions. Red litmus paper turns blue in alkaline solutions, and remains red in acidic solutions. However, neutral litmus paper turns red in acidic solutions and turns blue in alkaline solutions. At neutral pH, none of the litmus papers change colour. The litmus or pH paper is typically used for preliminary and small volume measuring. It cannot be used for continuous monitoring of a process. Though pH paper is fairly inexpensive, it may be affected by certain solutions, which may interfere with the colour change. Indicators are materials that change color when exposed to solution of different pH values. They are used in acid-base titrations, and include Phenolphthalein, Methyl red, and Bromothymol blue. These are also known as Universal indicators. Phenolphthalein turns colorless in acidic solutions and pink in basic solutions. It adopts four different states in aqueous solution: under strong acidic conditions, it exists in protonated form, showing orange colour, under acidic conditions, its lactone form is colorless. Under basic conditions, it shows singly deprotonated phenolate form and gives pink colour. However, in strongly basic solutions (> 13.0 pH), it becomes completely colourless. Methyl red is a dark red crystalline powder and it turns red in acidic solutions having pH < 4.4, yellow in pH > 6.2, and orange in between. Bromothymol blue is mostly used in applications that require measuring substances that would have a relatively neutral pH. A common use is for measuring the presence of carbonic acid. It acts as a weak acid in solution; thus, it can be in protonated (< pH 6.0, yellow) or deprotonated (> pH 7.6, blue) form. A pH Meter is an instrument that measures the hydrogen-ion concentration or pH of a solution, indicating its acidity or alkalinity. It measures the difference in electrical potential between a pH electrode and a reference electrode, and display the result converted into the corresponding pH value. Electrodes are the key parts: rod-like structures usually made of 2

glass, with a bulb containing the sensor at the bottom. Frequent calibration with solutions of known pH, generally before each use, ensures the accuracy of the instrument. The pH meters range from simple and inexpensive pen-like devices to complex and expensive laboratory instruments with computer interfaces and several inputs for indicator and temperature measurements to be entered to adjust for the variation in pH caused by temperature. Specialty meters and probes are available for use in special applications, harsh environments, etc. For very precise work the pH meter should be calibrated before each measurement. For normal use calibration should be performed at the beginning of each day. The reason for this is that the glass electrode does not give a reproducible measurement over longer periods of time. Calibration should be performed with at least two standard buffer solutions that span the range of pH values to be measured. For general purposes buffers at pH 4.00 and pH 10.00 are acceptable. The pH meter has one control (calibrate) to set the meter reading equal to the value of the first standard buffer and a second control which is used to adjust the meter reading to the value of the second buffer. A third control allows the temperature to be set. Standard buffer sachets/solution can be obtained from commercial suppliers. For more precise measurements, a three buffer solution calibration is preferred. As pH 7 is essentially, a "zero point" calibration (zeroing or taring a scale or balance), calibrating at pH 7 first, calibrating at the pH closest to the point of interest (either 4 or 10) second, and checking the third point will provide a more linear accuracy. Some meters allow a three-point calibration which is the preferred scheme for the most accurate work. Higher quality meters will have a provision to account for temperature coefficient correction, and high-end pH probes have temperature probes built in. After each single measurement, the probe must be rinsed with distilled water to remove traces of the solution adhering to the electrode, blotted with a clean wipe to absorb any remaining water.

Materials, Reagents and Solutions     

pH meter Wash bottle Distilled water pH solutions (pH 4.0, 7.0, 10.0) Kim-wipes (Lint-free wipes)

Procedure Calibration of pH meter: 1. Before you begin to calibrate and use your pH meter you first need to turn the pH meter on to warm up for an adequate time which generally takes around 30 minutes, but check your pH meter’s operating manual for the exact time. 2. Take the electrode out of its storage solution and rinse it with distilled water under an empty waste beaker. Once rinsed, blot dry with Kim-wipes. Avoid rubbing the electrode while wiping. 3. Since pH readings are temperature dependent, allow the standard buffers to reach to the room temperature. Pour the buffers into individual beakers for calibration. Buffers should be kept in a beaker for no longer than two hours. Do not pour used buffer back into its original container.

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4. Place the electrode in the buffer with a pH value of 7 and begin reading. Press the calibrate button to begin reading the pH, allow the pH to stabilize by letting it sit for approximately 1-2 minutes. Once the reading is stabilized, set the pH meter 7.0 by pressing the measure button. If you stir your buffer before measuring, be sure to stir all other buffers and samples in the same way. 5. Rinse and dry the electrode with Kim-wipes. 6. Place the electrode in the buffer with a pH value 4.0 and begin reading by pressing the measure button. Once the reading has stabilized, set the pH meter to pH 4.0. 7. Rinse and dry the electrode with Kim-wipes. 8. Use distilled water to rinse the electrode and dry with Kim-wipes. 9. If needed, calibrate the pH meter with standard pH buffer for 10.0. 10. Now, the pH meter is ready for taking measurement of pH of unknown samples. Measurement of pH of a sample solution: 1. Rinse the electrode with distilled water under and blot dry with Kimwipes Place the electrode in the sample, press the measure button and leave the electrode in the sample for approximately 1-2 minutes. 2. Once the reading has stabilized, press the measure button and note down the pH nalue of the sample. 3. Rinse the electrode with distilled water and blot dry with a lint-free Kimwipes. You can take reading of other samples or store the pH meter for using next time.

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Preparation of reagents, standard acid and alkali Introduction Commercially produced chemical reagents such as acids and ammonia are highly concentrated solutions. For example, commercially available concentrated sulfuric acid (H2SO4) is approximately 18 M. To prepare working solution of lower concentration for qualitative analysis or titration, a calculated volume of the concentrated solution is taken from the stock solution and then added to a specified volume of distilled water. However the volume of the concentrated solution to be taken depends on the information provided by the manufacturer on the label pasted on the bottle. For example, the label on the bottle of concentrated H2SO4 the specific gravity is mentioned to be 1.80. This means, the provided chemical is 1.80 times heavier than an equal volume of H2O. One ml of the acid solution weighs 1.80 g or 1000 ml of the acid solution weighs 1.80 x 1000 = 1800 g. Similarly, 98% by mass mentioned on the bottle means: 98 g of the acid (solute) is present in 100 g of the solution. In other words, 100 g of the acid solution contains 98 g of H2SO4. One liter (1000 ml) or 1800 g of the acid solution contains (98/100) x 1800 = 1764 g pure H2SO4. Therefore, its mass concentration = 1764 g/ L. Hence, its molar concentration, C = 1764 g/98 g = 18.0 Molar/L. How to calculate molar concentration: Generally, the original molar concentration (Co) of a chemical can be calculated using the formula Co = (10 x P x d) / M, where P = Percent purity (% purity), d = density or specific gravity of the chemical, and M = Molar mass. Exercise 1: Calculate the molar concentration (Molarity) of commercial trioxonitric acid ((HNO3) having specific gravity 1.42, purity 70.0%, and molar mass is 63.0. Answer:

Co = ?,

Given that P% = 70, d = 1.42,

Co = 10 x P x d M

=

M = 63

10 x 70 x 1.42 =15.8 mole/l 63

Exercise 2: How to prepare 40 mM hydrogen peroxide from 30% H2O2 solution? Answer: The chemical dataset indicates 30% (w/w) = weight per weight, which means that the solution contains 30 g H2O2 per 100 g of the solution. The density of the solution is 1.11 g/ml; therefore, 100 g solution have a volume of 90.09 ml. So 1 liter will contain 333 g of H2O2, which is a molar concentration of 9.79 M per liter. Therefore, to prepare 40 mM solution take 4.09 ml of 30% H2O2 solution and make up the volume to 1 liter with distilled water. How to calculate volume of solution to be diluted: For example, to prepare 500ml of 0.1 M H2SO4 from concentrated H2SO4 having Specific gravity of 1.82, Purity = 97% and Molar mass = 98. First of all calculate the concentration (Molarity) using the above formula. Molarity = 10 x P x d = 10 x 1.82 x 97 = 18.01 M M 98 5

Then, to calculate the amount of H2SO4 required to prepare 500 ml, 0.1 M H2SO4, the formula C1 x V1 = C2 x V2 can be used, where C1 = concentration of stock solution, V1 = volume of the of stock solution required, C2 = concentration of the working solution to be made, and V2 = volume of the of working solution to be made. C1 = 18.01 M, V1 = ?, V1 = C2 x V2 = 0.1 × 500 C1 18.01

C2 = 0.1 M,

V2 = 500 ml

= 2.78 ml

Therefore, 2.78 ml of concentrated H2SO4 will be required to prepare 500 ml, 0.1 M H2SO4. CAUTION: When diluting any concentrated acid, always add Acid to Water – never add water to acid.

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Preparation of buffers 1. Phosphate Buffer (0.1 M) A: 0.2 M solution of monobasic sodium phosphate (27.6 g NaH 2PO4.H2O or 31.2 g NaH 2 PO 4 .2H 2 O in 1000 ml) B: 0.2 M solution of dibasic sodium phosphate (35.6 g Na 2HPO4.2H2O or 53.6 g Na2HPO4.7H2O or 71.6 g of Na 2HPO4.12H2O in 1000 ml). M i x X m l of A and Y m l of B and m ak e up t he fi nal vol um e t o 200 m l wi t h di s t i l l ed wat er t o t he desi red pH (as pe r t he t abl e gi ven bel o w) . pH

X

Y

5.7

93.5

6.5

5.8 5.9 6.0 6.1 6.2 6.3 6.4 6.5 6.6 6.7 6.8 6.9 7.0 7.1 7.2 7.3 7.4 7.5 7.6 7.7 7.8 7.9 8.0

91.4 90.0 87.7 85.0 81.5 77.5 73.5 68.5 62.5 56.5 51.0 45.0 39.0 33.0 28.0 23.0 19.0 16.0 13.0 10.5 8.5 7.0 5.3

8.6 10.0 12.3 15.0 18.5 22.5 26.5 31.5 37.5 43.5 49.0 55.0 61.0 67.0 72.0 77.0 81.0 84.0 87.0 89.5 91.5 93.0 94.7

A l t e r n a t i v e l y, a s m a l l v o l um e o f A a n d a d d B s l o w l y w h i l e m e a s u r i n g t h e p H o f t h e m i x t u r e u n t i l t h e d e s i r e d pH i s o b t a i n e d . Th e n m a k e u p t h e f i n a l v o l u m e t o 2 0 0 m l w i t h di s t i l l e d w a t e r .

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2 . T r i s - H C I b u f f e r ( 0. 0 5 M )

A : 0.2 M Tris (hydroxy methyl) amino methane (24.2 g in 1000 ml) B: 0.2 M HCI Take 50 ml of solution A and add X ml of sol ution B. Make up t he fi nal vol um e t o 200 m l wi t h di st i l l ed wat er t o t he desi red pH (as per t he t abl e gi ven b el ow) . X

pH

5.0 8.1 12.2 16.5 21.9 26.8 32.5 38.4 41.5 44.2

9.0 8.8 8.6 8.4 8.2 8.0 7.8 7.6 7.4 7.2

A l t e r n a t i v e l y, t ake 50 ml of solution A and a d d B s l o wl y w h i l e m e a s u r i n g t h e p H o f t h e m i x t ur e u n t i l t h e d e s i r e d p H i s o b t a i n e d . T h en m a k e u p t h e f i n a l v o l u m e t o 2 0 0 m l w i t h di s t i l l e d w a t e r . 3. Acetate Buffer (0.1 M)

A: 0.2 M solution of glacial acetic acid (11.51 ml in 1000 ml) B: 0.2 M solution of sodium acetate (16.4 g of C 2 H 3 O 2 Na or 27.2 g C 2 H 3 O 2 Na.3H 2 O in 1000 ml). M i x X m l of A and Y m l of B and m ake up t he fi nal vol um e t o 1 00 m l wi t h di s t i l l ed wat er t o t he desi red pH (as pe r t he t abl e gi ven bel o w) . pH

X

Y

3.6 3.8 4.0 4.2 4.4 4.6 4.8 5.0 5.2 5.4 5.6

46.3 44.0 41.0 36.8 30.5 25.5 20.0 14.8 10.5 8.8 4.8

3.7 6.0 9.0 13.2 19.5 24.5 30.0 35.2 39.5 41.2 45.2

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A l t e r n a t i v e l y, a s m a l l v o l u m e o f A a n d a d d B sl o w l y w h i l e m e a s u r i n g t h e p H o f t h e m i x t u r e un t i l t h e d e si r e d p H i s o b t a i n e d. T h e n m ak e u p t h e f i n a l v o l u m e t o 1 0 0 m l w i t h di s t i l l e d w a t e r .

4. Glycine-HCI buffer (0.05 M) A: 0.2M solution of gl ycine (15.01 g in 1000 ml) B: 0.2 M HCl Take 50 ml of solution A and add X ml of sol ution B. Make up t he fi nal vol um e t o 200 m l wi t h di st i l l ed wat er t o t he d esi red pH (as pe r t h e t abl e gi v en bel ow) . X

pH

5.0 6.4 8.2 11.2 16.8 24.2 32.4 44.0

3.6 3.4 3.2 3.0 2.8 2.6 2.4 2.2

A l t e r n a t i v e l y, t ake 50 ml of solution A and a d d B s l o wl y w h i l e m e a s u r i n g t h e p H o f t h e m i x t ur e u n t i l t h e d e s i r e d p H i s o b t a i n e d . T h en m a k e u p t h e f i n a l v o l u m e t o 2 0 0 m l w i t h di s t i l l e d w a t e r . 5. Citrate Buffer (0.1 M) A: 0.1 M solution of citric acid (21.1 g in 1000 ml) B: 0.1 M solution of sodium citrate (29.41 g C 6HSO7.Na3.2H2O in 1000 ml) Note: Use of the salt with 5H 2 0 is not recommended. M i x X m l of A and Y m l of B and m ak e up t he fi nal vol um e t o 1 00 m l wi t h di s t i l l ed wat er t o t he desi red pH (as pe r t he t abl e gi ven bel o w) . X

Y

pH

46.5 43.7 40.0 37.0 35.0 33.0

3.5 6.3 10.0 13.0 15.0 17.0

3.0 3.2 3.4 3.6 3.8 4.0

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31.5 28.0 25.5 23.5 20.5 18.0 16.0 13.7 11.8 9.5 7.2

18.5 22.0 24.5 27.0 29.5 32.0 34.0 36.3 38.2 41.5 42.8

4.2 4.4 4.6 4.8 5.0 5.2 5.4 5.6 5.8 6.0 6.2


6. Citrate-Phosphate Buffer A: 0.1 M solution of citric acid (19.2 g in 1000 ml) B: 0.2 M solution of dibasic sodium phosphate (53.6 g of Na2HPO4.7H20 or 71.6 g of Na2HPO4.12 H2O in 1000 ml) M i x X m l of A and Y m l of B and m ak e up t he fi nal vol um e t o 1 00 m l wi t h di s t i l l ed wat er t o t he desi red pH (as pe r t he t abl e gi ven bel o w) . X

Y

44.6

5.4

42.2 39.8 37.7 35.9 33.9 32.3 30.7 29.4 27.8 26.7

7.8 10.2 12.3 14.1 16.1 17.7 19.3 20.6 22.2 23.3

25.2

24.8

pH

X

Y

pH

2.6

24.3

25.7

5.0

2.8 3.0 3.2 3.4 3.6 3.8 4.0 4.2 4.4 4.6 4.8

23.3 22.2 21.0 19.7 17.9 16.9 15.4 13.6 9.1 6.4

26.7 27.8 29.0 30.3 32.1 33.1 34.6 36.4 40.9 43.6

5.2 5.4 5.6 5.8 6.0 6.2 6.4 6.6 6.8 7.0


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7. Carbonate-Bicarbonate Buffer, 0.05 M A: 0.2 M solution of anhydrous sodium carbonate (21.2 g in 1000 ml) B: 0.2 M solution of sodium bicarboante (16.8 g in 1000 ml) M i x X m l of A and Y m l of B and m a ke up t he fi n al vol um e t o 2 00 m l wi t h di st i l l ed wat er t o t he desi red pH ( a s per t he t abl e gi ven bel ow) . X

Y

4.0

46.0

7.5 9.5 13.0 16.0 19.5 22.0 25.0

pH

X

Y

pH

9.2

27.5

22.5

10.0

42.5 40.5 37.0 34.0 30.5 28.0

9.3 9.4 9.5 9.6 9.7 9.8

30.0 33.0 35.5 38.5 40.5 42.5

20.0 17.0 14.5 11.5 9.5 7.5

10.1 10.2 10.3 10.4 10.5 10.6

25.0

9.9

45.0

5.0

10.7

A l t e r n a t i v e l y, a s m a l l v o l um e o f A a n d a d d B s l o w l y w h i l e m e a s u r i n g t h e p H o f t h e m i x t u r e u n t i l t h e d e s i r e d pH i s o b t a i n e d . Th e n m a k e u p t h e f i n a l v o l u m e t o 2 0 0 m l w i t h di s t i l l e d w a t e r . 8. N-2-Hydroxyethylpiperazine-N'-ethanesulphonic acid (HEPES)-NaOH buffer

0.05 M, pH 7.0  8.2 at 21°C Take 25 ml of 0.1 M HEPES (23.83 g/1), add X ml of 0.1 M NaOH and make up the final volume to 50 ml with double distilled H 2 O X ml (0.1 M NaOH) pH 7.0 7.2 7.4 7.6 7.8 8.0 8.2

21°C 0 6.6 8.7 11.2 13.7 16.3 18.8

37°C 7.4 9.9 12.3 14.6 17.1 19.5 0

A l t e r n a t i v e l y, t a k e 2 5 m l o f H EP ES ( 0. 1 M ) a n d a d d sl ow l y 0.1 M NaOH w h i l e m e a s u r i n g t h e p H o f t h e m i x t u r e a t d e s i r e d t e m p e r a t u r e ( e . g. 2 1 °C) u n t i l t h e d e s i r e d p H i s ob t a i n e d . T h e n m a k e u p t h e f i n a l vo l u m e t o 5 0 m l w i t h d ou b l e di s t i l l ed w a t e r . 11

Molarity, Molality and Normality Although molarity and molality are homophones, they cannot be interchanged. While molarity is a measurement of the moles in the total volume of the solution, molality is a measurement of the moles in relationship to the mass of the solvent. Like molarity, normality relates the amount of solute to the total volume of solution. However, normality is specifically used for acids and bases. When water is used as solvent and the concentration of the solution is low, these differences can be negligible (d = 1.00 g/ml). However, when the density of the solvent is significantly different than 1 or the concentration of the solution is high, these changes become much more evident. The mole-equivalents of an acid or base is calculated by determining the number of H+ or OH ions per molecule. Therefore, normality (N) = n × M (where n is an integer). For an acid solution, n is the number of H+ ions provided by a formula unit of the acid, and for a basic solution, n is the number of OH ions provided by a formula unit of the base. Molarity (M) is the number of moles of Molarity = moles of solute/liter of solution.

solute

per liter of solution.

Molality (m) is the number of moles of solute per kilogram of solvent. Molality = moles of solute/kilograms of solvent. Normality

(N) is

the

number

of mole-equivalents per liter of solution.

Normality = number of mole-equivalents/liter of solution. Example 1. Compare the Molar and Molal volumes of 1 mol of a solute dissolved in CCl4 (d = 1.59/ml). For a 1 Molar solution, 1 mol of solute is dissolved in CCl4 until the final volume of solution is 1 liter. For a 1 Molal solution, 1 mol of solute is dissolved in 1 kg of CCl4. 1 kg of CCl4 × (1,000 g/1 kg) × (ml/1.59 g) = 629 mL CCl4 Example 2. A 3 M H2SO4 solution is the same as a 6 N H2SO4 solution. As for an acid solution, n is the number of H+ ions provided by the formula unit of the acid. Example 3. A 1 M Ca(OH)2 solution is the same as a 2 N Ca(OH)2 solution. As for a basic solution, n is the number of OH ions provided by a formula unit of the base. Note: The normality of a solution is never less than its molarity !

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