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LABORATORY 1: WATER, pH, BUFFERS, AND AMINO ACIDS

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INTRODUCTIONAll life takes place in water. Biochemical reactions occur in water catalyzed by

enzymes and contained within lipid bilayers whose structures depend on water. The unique

physical and chemical properties of water must be understood in order to comprehend these

structures and functions of proteins, carbohydrates, lipids and nucleic acids. To illustrate

water=s properties and its interactions with biological molecules, this lab will investigate the

dissociation of water with a simple buffer and a simple biomolecule, an amino acid.

The Dissociation of Water:

Water molecules dissociate into hydrogen ions (protons or H+) and hydroxyl ions (OH-).

H2O H+ + OH- (1)

As the dissociation of water is an equilibrium process, we can write, from the Law of Mass Action,the dissociation equation as follows:

[][]

[]

where: K = the dissociation constant (also called the equilibrium constant), and the bracketed

items represent concentrations in moles per liter.

For pure water, at standard temperature and pressure, the probability of a hydrogen atom

being ionized is P = 1.8 x lO-9. Protons (H+) as such are not present in water; the positive charges

exist in the form of hydronium ions (H3O+) which are hydrated, through hydrogen bonding, by

other water molecules to form the H9O4+ ion and other hydrates. However, for biochemical

purposes, the symbol H+ is used to designate the positively charged species.

Since one gram molecular weight of water is 18, the concentration of water molecules in aliter of water is 1000 grams/liter divided by 18 grams/mole to give 55.55 moles per liter. Theconcentration of hydrogen ions or hydroxyl ions can be obtained by multiplying the probability of

the ions times the molar concentrations of water as follows:

[H+] = [OH

-] = (1.8 x 10

-9) (55.55 moles/liter) (3)

[H+] = 10

-7 molar (at 25C) (4)

The ion product of water (Kw) is: Kw = [H2O]K = [H+] [OH

-] (5)

Kw = (55.55)K = (10-7) (10-7) (6)

Kw = 10-14

(at 25C) (7)

In 1909 Sorensen published two classical papers on hydrogen ion concentration. Since

the range of hydrogen ion concentration is so great in aqueous solutions, Sorensen initiated an

exponential notation employing the term pH. He defined pH as the negative log of the hydrogenion concentration.

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pH = -log[H+] (8)

For pure water then: pH = -log[10-7

] (9)

pH = 7 (at 25C) (10)

Since Kwis less than 10-14

at temperatures less than 25C, the pH of pure water will be greater

than 7 below 25C, while the opposite is true above 25C.

Definitions of Acids and Bases:

Although acids and bases were used in both industry and science for centuries, the first

theoretical treatment of acid behavior came at the end of the eighteenth century. Antoine-LaurentLavoisier proposed that all acids contained oxygen, but his hypothesis was destroyed by the fact

that HCl was found not to contain oxygen. Sir Humphrey Davy, in 1815, suggested that hydrogen

was the critical element in acids. Other chemists, like the German scientist, Liebig, classified the

role of acids, but bases were considered to be simple acid neutralizers. Ostwald and Arrheniusdeveloped their theories of electrolyte dissociation that considered acids as substances that coulddissociate to form protons and bases as substances that could dissociate to form hydroxyl ions.

A more generalized definition was proposed, essentially simultaneously, by two separate

scientists, Johannes Nicolaus Bronsted (in Denmark) and Thomas Martin Lowry (in England).

This definition became known as the Bronsted-Lowry formalism. Although even more general

definitions for acids and bases were later proposed by Lewis and others, the formalism introducedby Bronsted and Lowry in 1923 has become the standard for acid-base interactions in

biochemistry. According to the Bronsted-Lowry formalism, an acid is a proton donor, and a base

is a proton acceptor. Acid-base interactions in the Bronsted-Lowry formalism always involve aconjugate acid-base pair, formed by the proton donor and the proton acceptor. Thus, an acid-basereaction is one in which a proton transfer occurs.

Acid Base + H+ (11)

To generalize, an acid is HA and its conjugate base is A- . The reaction for the dissociation of anacid to its conjugate base and a proton is:

HA A-+ H

+ (12)

and the equilibrium constant K can be obtained by:

[][]

[]

For example, acetic acid (CH3COOH) is a proton donor, while acetate (CH3COO-) is a proton

acceptor. The dissociation of acetic acid can be written as follows:

CH3COOH CH3COO- + H

+ (14)

The dissociation constant is given by:

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[

][]

[]

Strong Acids and Bases and the pH of Water

Depending on their dissociation constants, acids are classified as strong acids if they

dissociate almost completely in water or as weak acids if they only partially dissociate in water.For example, if a strong acid such as HCl is added to water, it produces hydrogen ions

equal to the molarity of the original acid added because, for all practical purposes, it completelydissociates. The total hydrogen ion concentration will, therefore, arise from two sources: (1) the

dissociation of the HCl added, and (2) the existing dissociation of water itself. For example, if 8 x

10-3

molar HCl were added to water, the total hydrogen ion concentration would be 8 x 10-3

from

the HCl plus 1 x 10-7

molar from water which equals 8.0001 x 10-3

molar. Therefore, the pH =

2.1. Often the contribution of ions by the water is so small it can be ignored, because it is lost

when rounding to the most significant figure.The calculation of pH for a basic solution is equally as simple. A strong base will produce

hydroxyl ions equal to the original concentration of base added. The total hydroxyl ion

concentration is, therefore, the sum of the contributions of the base added and the water. For

example, if 1 x 10-3

mole of NaOH were added to one liter of water, the total hydroxyl

concentration would be 1 x 10-3

molar from the base plus 1 x 10-7

molar from the dissociation of

water yielding 1.0001 x 10-3

molar hydroxyl ions. Because of the small contribution by water we

can round the figure to 1 x 10-3

molar. To calculate the pH of this solution, recall that:

Kw = 10-14

= [H+] [OH-] (16)

Now substituting the concentration of NaOH (1 x 10-3

molar):

10-14

= [H+] [10-3

] (17)

[H+] = 10-11

(18)

Therefore, the pH of the solution is:

pH = -log[H+] = 11 (19)

Weak Acids and Bases and the pH of Water

Most biological acids and bases are classified as weak. To understand how thesechemical groups effect the pH of a cell or the blood, we must relate the dissociation constant K to

the pH. The real thermodynamic equilibrium constant, K, is unchangeable for a set of physicalconditions (e.g., temperature and pressure), but the apparent equilibrium constant (K'), isdependent on the type of ionic species in solution, and their thermodynamic activity. The value of

K' is experimentally determinable and is called the apparent equilibrium constant.

[][]

[]

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Rearranging:

[][]

[]

Taking the negative log of both sides:

[] []

[]

The term pH is defined as -log [H+]. The term pK is defined as -log K. This is a convenientway of expressing a dissociation constant in logarythmic terms.

[]

[]

Isolating the pH term:

[]

[]

Restating:

[ ]

[ ]

Equation 24 is known as the Henderson-Hasselbalch equation. It relates the dissociation constant

of a weak acid with the pH and the concentration of the weak acid and its conjugate base, and isfundamental to any quantitative treatment of acid-base equilibria.

Titration of Weak Acids and Bases:

NaOH is a strong base, and its dissociation in an aqueous environment can be written as:

NaOH Na+ + OH

- (26)

Acetic acid, from (14) dissociates as:

CH3COOH CH3COO- + H

+ (27)

The OH- contributed by NaOH can react with the H

+ present and that coming from acetic acid as

well in the formation of water.

OH-+ H

+ H2O (28)

In a solution of acetic acid, as NaOH is added, the concentration of H+is reduced. In order

to maintain the equilibrium conditions, acetic acid dissociates to H+and acetate ion in order to

replace the lost H+. Therefore, as NaOH is added, the [CH3COOH] will decrease and the

[CH3COOH] will increase. If the pH is plotted against the amount of OH-ions (NaOH) added, a

titration curve is obtained as illustrated in Figure 1. The titration curve is the graphical expressionof the Henderson-Hasselbalch equation.

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FIGURE 1: Titration of one mole of acetic acid with NaOH. The forms and relative

concentrations of acetic acid and acetate ion are shown in the graph. At pK, the concentration of

proton donor and acceptor are equal.

The titration curve for acetic acid has an inflection point, the pH value of whichcorresponds to the pK'of the acid. At that pH, the concentration of acetate and that of acetic acid

are equivalent. If the concentration of acetate (A- ) is made equal to the concentration of aceticacid (HA) in Henderson-Hasselbalch equation, the equation reduces to pH = pK.

Buffers:

A nearly flat region in the titration curve extends to both sides of the point where pH = pK.

Within that zone the conjugate acid-base pair is called a buffer because the addition of OH- or H+

ions will have little effect on the pH of the solution. As a rule of thumb, the buffer zone of a

conjugate acid-base pair extends 1 pH unit to either side of the point where pH = pK.

The buffering capacity, or the capacity to absorb excess H+or OH

- ions, depends upon the

concentration of the acid-base pair, the pH and the pK.

Two important buffering systems for vertebrates are the carbonic acid/bicarbonate buffersystem for extracellular fluids and the phosphate buffer system for intracellular components.

Polyprotic acids such as phosphoric acid (H3PO4) or carbonic acid (H2CO3) are capable of

dissociating more than one proton. Phosphoric acid, for example, can dissociate a total of 3protons according to the following dissociation schemes:

H3PO4 H++ H2PO4

-1 pK = 2.14

H2PO4-1

H++ HPO4

-2 pK = 7.20

HPO4-2

H++ PO4

-3 pK = 12.40

The most biologically important conjugate acid-base pair of the phosphate system is the middle

one because the pK'is near the normal pH of the biological systems. Each dissociation step for

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the titration of phosphate is defined by the Henderson-Hasselbalch equation. The dissociationscheme of phosphoric acid as indicated above and the titration curve of phosphoric acid (Figure 2)

illustrate some important facts concerning the dissociation of polyprotic acids. First, a particular

ionic species (e.g. H2PO4-1

can act as proton acceptor in one titration step while in the following

titration step the same ionic species is the proton donor. Secondly, even though at a given pH,two ionic species will dominate (e.g. H2PO4

-1and HPO4

-2at pH, 7.2) trace amounts of other ionic

species are also present since there is an equilibrium between the three steps. Carbonic acid

(H2CO3) follows a dissociation pattern similar to that of phosphoric acid, however, there are onlytwo dissociation steps.

Fig. 2: Titration of one mole of phosphoric acid with NaOH. Note the three pKvalues and thenumber of equivalents needed to reach the three pKs.

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The Dissociation of Amino Acids:

During the dissociation of phosphoric acid (H3PO4) one ionic species, e.g. H2PO4-1

, can actas a proton acceptor as well as a donor. Compounds which have this unusual property are knownas amphoteric substances, or amphoteric electrolytes often abbreviated as ampholytes. A group

of natural ampholytes found is the amino acids and their polymeric forms, polypeptides and

proteins.

The simplest amino acid, glycine, can exist in several ionic forms:

Fully protonated: H3N+-CH2-COOH (HAH

+)

Unprotonated: H2N-CH2-COO- (A-)

Intermediate forms: H3N+-CH2-COO

- and H2N-CH2-COOH

(AHa) (AHb)

The intermediate forms are electrically neutral. In form AHa, glycine exists as a charged state inwhich the charges neutralize each other, while in form AHb the amino acid is uncharged. Both

forms can act as proton donors or acceptors, and hence qualify as ampholytes. Bjerrum (1923)discovered that glycine, as well as other amino acids, crystallize in the ampholytic form. Amino

acids in crystalline form have relatively high melting and decomposition points (e.g., melting point

for glycine: 233C) which can be ascribed to the fact that the molecules in the crystalline latticeinteract by electrostatic forces between opposite charges.

If glycine in its fully protonated state is titrated with base, the titration curve shown in Fig.

3 is obtained. The titration curve consists of two buffering regions, one representing thedissociation of the carboxyl group, the other the dissociation of the amino group. Each region of

the curve is mathematically defined by the Henderson-Hasselbalch equation. The point where the

two regions of the titration curve for glycine join is defined as the isoelectric pH, the pI. At thispH, primarily the ampholytic form exists in solution. Glycine has two dissociable groups, (the

-carboxyl group and the -amino group), two pK's (pK1and pK2), and hence two buffering

ranges, from pH 1.2 to 3.2 and from pH 8.5 to 10.5. At the pI, no buffering power exists asindicated in the titration curve (Fig. 3) by the steep slope that exists around pI. The value for pI

can be calculated by taking the arithmetic mean of pK1 and pK2. The dissociation of glycine is

the simplest case for an amino acid.

All amino acids which do not have other ionizable groups show titration patterns similar

to glycine. Of the 20 amino acids commonly occurring in proteins, five have ionizable groups

in addition to the-amino and -carboxyl groups. They include: 1) the second carboxyl group

found in aspartic acid and glutamic acid, 2) the amino group on the aliphatic chain of lysine, 3) the

imidazole ring in histidine and, 4) the guanidinium group in arginine. All of these amino acids

have three pK's, pK1 and pK2,for the -carboxyl and -amino group and pKRfor the additionalionizable group. In general, the same consideration as discussed for glycine will apply to these

amino acids. Typical titration curves for aspartic acid, lysine and histidine can be found in Figs.

4-6.

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Moles of KOH added

Fig. 3: Titration of one mole of glycine, free acid, with KOH.

Moles of KOH added

Fig. 4: Titration of one mole of aspartic acid with KOH. Note that the pKs of the primarycarboxyl and the R-group carboxyl are so close that there is no inflection point between them.

One can determine that there are two different carboxyl groups within the one continuous slope

by two observations. The entire slope required two moles of KOH / mole of aspartic acid, i.e.two acid groups per molecule were titrated. Also, the slope extended from pH 1 to pH 5. A

single group usually buffers over a range of only 2 pH units.

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Moles of KOH added

Fig. 5: Titration of one mole of lysine, free acid, with NaOH. Note here again that the pKs of the

primary amino and R-group amino functions are so close that there is no inflection point betweenthem. Note also that the single slope spans pH 8 to pH 12 and requires two moles of KOH / moleof lysine.

Moles of KOH added

Fig. 6: Titration of one mole of histidine, free acid with NaOH. Note that the first pK is due to

the carboxyl group; the second pKis due to the imidazole R-group; and the third pKis due to the

amino group.

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Materials

Equipment:

1. pH meter & electrode immersed in a beaker of deionized water

2. Standard buffers, pH 4.00, 7.00 & 10.00, in 50 ml tubes with stirring fleas3. One empty 50 ml tube with stirring flea

4. magnetic stirrer

5. deionized water in squeeze bottle6. 1-10 ml pipets7. Kimwipes

Solutions:1. 0.100 M acetic acid

2. 0.100 M unknown amnio acid

3. 0.500 M potassium hydroxide (KOH)

Methods

Daily Laboratory Clean-up:1. Wash your used glassware and return it to your lab glassware kit.2. Carefully place dirty pipets (tips down) in pipet containers.

3. Cover all solutions with lids.

4. Make sure electrode tips are submerged in the water beaker.5. Turn battery-operated pH meters off and others to standby.

6. Turn off magnetic stirrers.

Use of the pH meterRead these instructions before attempting to use the pH meter. If your pH meter develops

trouble or the readings are variable, call your instructor -- do not attempt to repair the meter.

Adhere to the following guidelines:

1. Do not bump your electrode on the bottom of the beaker.

2. The function switch should always be on standby when the electrodes are out of the

solution.3. Keep the filling hole of the electrode covered with the loose rubber sleeve during

measurements.

4. When rinsing the electrode, rinse only the lower half. Avoid squirting water into thefilling hole of the electrode.

5. After rinsing, blot the excess water by touching bottom of the electrode with a folded

Kimwipe. Do not wipe the electrode, or you will generate a static charge on the glass and

throw off your measurements. Since you are blotting clean water, the Kimwipe may beused several times.

In this laboratory you will be using a single glass-calomel combination electrode. This

combination is designed to produce zero millivolts at pH 7.00.

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Dual Standardization of a pH Meter:

1. Rinse electrode, blot, and immerse in pH 7.00 standard buffer while stirring.

2. Set meter to 7.00 with the

Calibrate knob (on Corning pH meters)Intercept knot (on Sargeant-Welch pH meters)

Assymetry knob (on Beckman pH meters)

3. Rinse the electrode, blot, and immerse it inpH 4.00 standard buffer (for the range pH 1-7)

or

pH 10.00 standard buffer (for the range, pH 7-12)

while stirring.4. Set the meter to 4.00 or 10.00 using the Temp.C knob.

5. Rinse the electrode, blot, and perform your titration. The meter is now accurate in the pH

1-7 range or in the pH 7-12 range. (NOT IN BOTH)

To Read the pH of an Unknown Solution:

1. With the function switch on standby, the electrode is washed with distilled water andthen gently blotted with a tissue (carefully -- an electrode does not respond favorably torough treatment!!).

2. Place your clean electrode into the unknown solution and turn the function switch to pH

(or read and pH in some cases).3. Record the pH while the solution is stirring. If you are titrating an unknown solution, add

single aliquots of acid or base while the solution continues to stir, and record the pH aftereach addition when the pH stabilizes.

4. At the end of the experiment, take your electrode out of the solution. With the functionswitch on standby, the electrode is washed with distilled water and then gently blotted with

a tissue.

5. Always store your pH electrode with the glass bulb in a beaker of deionized water.Never allow it to dry in the air.

EXPERIMENTS

Experiment 1: Titration of Pure Water

1. Standardize your pH meter for the pH 7-12 range.

2. Rinse your titration vessel thoroughly with deionized H2O (dH2O). Do not allow the small

magnetic stirring bar, the Aflea@, to drop into the drain.

3. Fill with 20.0 mls of dH2O; insert the pH electrode; and record the pH. The measured pH of

water will drift between 4 and 9 because the response time of the electrode is very long in theabsence of other ions. Just record the initial pH and begin your additions. Stir continuouslythroughout the experiment.

4. Add 0.100 ml portions of 0.500 M KOH, and record the pH after each addition. Stop whenthe pH reaches 12.

5. Plot pH vs. volume of KOH added.

Experiment 2: Titration of a Buffer, Acetic Acid

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1. Standardize your pH meter for the pH 1-7 range.

2. Rinse your titration vessel thoroughly with dH2O. Do not allow the small magnetic stirring

bar, the Aflea@, to drop into the drain.

3. Place 20.0 ml of 0.100 M acetic acid in your titration vessel; insert the pH electrode; begin

stirring; and determine the pH of your solution. Make sure the solution stirs continuously.4. Add 0.400 ml aliquots of 0.500 M KOH from a P1000 pipettor or a 5 ml pipette. Determine

the pH of your solution after each addition of base. Record your data in a table as pH versus the

total volume of KOH added.5. As you approach the complete neutralization of the acetic acid, the pH/vol of KOH will

begin to increase (see Fig. 1). Decrease the aliquot size to 0.100 ml of KOH. The point ofcomplete neutralization occurs at the maximum pH/vol of KOH. In this experiment, you do

not have to restandardize when the pH passes 7. Stop when the pH passes 10.6. Plot the pH of the solution versus the volume of KOH added.

pKIDENTIFICATION CRITERIA

1) For a single weak acid, the pK=is the midpoint of the buffering range, the point where the

change in pH vs. the change in volume of titrant is a minimum. A clearer way to find the

midpoint is to find the volume at which all the weak acid has been neutralized by the KOH, thepoint of complete neutralization, where the pH shoots upward. Divide that volume by half, and

using that value on the x axis, locate the corresponding pH on the titration curve. That pH will be

the pK=.

2) If there are two chemical groups on the same amino acid with pK=s close to each other,

such as the 2 carboxyls on aspartic acid or the two aminos on lysine, the buffering region of the two

groups should span 2 moles of KOH per mole of amino acid. To find the pK=s, divide the

combined buffering region in two, and assign the pK=s at the half-way points in each of the two

regions.

Experiment 3: Titration of Unknown Amino Acid

You will be given a 0.100 M solution of an unknown amino acid which has been brought to

pH 1.0 during dissolving. The unknown could be glycine, aspartic acid, histidine or lysine. You

are to determine the unknown=s identity from its titration curve, the graph of pH vs. vol. of KOH

added.

1. Standardize your pH meter for the pH 1-7 range.2. Rinse your titration vessel thoroughly with dH2O. Do not allow the small magnetic stirring

bar, the Aflea@, to drop into the drain.

3. Pipet a 20 ml vol. of 0.100 M unknown amino acid in your titration vessel. Insert the pH

electrode, and begin stirring. Record the pH of your solution. Make sure the solution stirs

continuously.4. Add 0.4 ml aliquots of 0.500 M KOH, recording the pH after each addition of base, until pH 7 ispassed. Remove the electrode and set aside your sample. Restandardize your pH meter for the pH

7-12 range. Reinsert the pH electrode into your sample, and continue to add base until pH 12 is

reached. Plot pH vs. volume of KOH added.

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Lab 1 Water, pH, Buffers & Amino Acids Lab Report

1. Graph Experiment 1.

2. Calculate the pH of 20 mls of water plus 0.100 ml of 0.500 M KOH. Compare this with thepH observed in Exp.1. Does pure water have any buffer capacity?

3. Graph Experiment 2. Identify the pKof acetic acid graphically using the pK=identificationcriteria.

4. Graph Experiment 3. Identify the unknown amino acid. On the graph, indicate the pK=s of

the - carboxyl group, the - amino group, and the R-group(if any).

5. a. In Experiment 2, from the volume of KOH added and the pH observed, calculate the pK=

using the Henderson-Hasselbach equation using the following table:

Assume initially that [HA] = 0.1 M, and that HA + KOH A- + H2O, i.e. [KOH] = [A- ].

KOH pH [HA] [A- ] pK=

vol. Observed calculated

_______ _______ _______ _______ _______

0.4 ml0.8

1.2

1.62.0

2.4

2.83.2

3.6

5.b. What is the mean of the calculated pK=values? What is the standard deviation?

5.c. How closely does the mean agree with the value determined graphically?

5.d. Comparing all calculated pK=values, does the variation appear to be due to errors of

accuracy and precision, gross accidental errors or systematic errors?Errors of accuracy and precision relate to how close a measured value on a particular instrument is

to the true value.

Gross accidental errors relate to incorrect technique, incorrect calculation or mismatching units.

Systematic errors relate to errors in zero adjustment, calibration, faulty primary standards or errorsin reading the instrument.

6. What weights of the zwitterionic Good buffer HEPES, N-[2-hydroxethyl]piperazine-N-[2-ethanesulfonic acid], MW 238.30 g/mole) and NaOH (MW 40.01 g/mole) are required to

make 1.00 L of 0.100 M HEPES/NaOH buffer, pH 7.00? The pKa of the HEPES buffer is 7.50.

The buffer concentration of 0.100 M refers by convention to the concentration of the buffering ion,the HEPES, both protonated and deprotonated together.

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GOALS1. Definitions

ion product of water hydrogen ion

hydroxyl ion acid, base

neutralization molaritystrong acid/base weak acid/base

pH K'

pK' bufferHenderson-Hasselbach Eq. protonated/deprotonated

2. Operation of a pH meter.

3. Role of weak acids and bases as buffers in controlling pH. Calculation of buffer capacity.

4. Calculation of a buffer concentration from the concentration and volume of acid or baserequired to neutralize it completely.

5. Identification of number and pK' of titratable groups by inflection point,

moles of amino acid / mole of KOH, and span of buffer zone.

6. Identification of glycine, aspartic acid, lysine, or histidine from number and pK'of

titratable groups.

7. Accurate measurement of volumes in pipets and pH from analog meters.

REFERENCES

1. Chapter 2, Mathews, C. K. And Van Holde, K. E., Biochemistry , 2nd Ed.,

Benjamin/Cummings, Menlo Park, CA (1995).

2. Bates, R.G. Electrometric pH Determination, Theory and Practice. John Wiley and

Sons, Inc., New York (1954).

3. Christensen, H.N. pH and Dissociation. W.B. Saunders Co., Philadelphia (1964).

4. Davenport, H.W. ABC or Acid-Base Chemistry, 5th Ed. University of Chicago Press,

Chicago (1969).

5. Eisenman, G., Butes, R., Muttock, G. and Friedman, S.M. The Glass Electrode. John

Wiley and Sons, New York (1966).

6. Segel, I.H. Biochemical Calculations, 2nd. Ed. Wiley and Sons, Inc., New York (1976).