Lab.13. pH measurement

Key words:

pH, pOH, acids, bases, buffers, electrodes, buffer capacity.

Literature:

J. Crowe, T. Bradshaw, P. Monk, Chemistry for the Biosciences. The essential concepts,

Oxford University Press, 2006; Chapter 17, 515 – 545.

J. Brady, N. Jespersen, A. Hysop, Chemistry, International Student Version, 7thed. Wiley,

2015, Chapter 16, 762 – 798.

P. Monk, Physical Chemistry , Wiley 2004, Chapter 6, 233- 276.

Theoretical background

Introduction

In chemistry, pH is a measure of the acidity or basicity of an aqueous solution. Pure water is

said to be neutral, with a pH close to 7.0 at 25 C (77 F). Solutions with a pH less than 7 are

said to be acidic and solutions with a pH greater than 7 are basic or alkaline. pH

measurements are important in medicine, biology, chemistry, food science, environmental

science, oceanography, civil engineering and many other applications.

In a solution pH approximates but is not equal to p[H], the negative logarithm (base 10) of the

molar concentration of dissolved hydronium ions (H3O+); a low pH indicates a high

concentration of hydronium ions, while a high pH indicates a low concentration.

Crudely, this negative of the logarithm matches the number of places behind the decimal

point, so for example 0.1 molar hydrochloric acid should be near pH 1 and 0.0001 molar HCl

should be near pH 4 (the base 10 logarithms of 0.1 and 0.0001 being −1, and −4,

respectively).

Pure (de-ionised) water is neutral, and can be considered either a very weak acid or a very

weak base (center of the 0 to 14 pH scale), giving it a pH of 7 (at 25 C (77°F)), or 0.0000001

M H+. For an aqueous solution to have a higher pH, a base must be dissolved in it, which

binds away many of these rare hydrogen ions. Hydrogen ions in water can be written simply

as H+ or as hydronium (H3O+) or higher species (e.g. H9O4

+) to account for solvation, but all

describe the same entity.

Lab.13. pH measurement

Most of the Earth's freshwater surface bodies are slightly acidic due to the abundance and

absorption of carbon dioxide; in fact, for millennia in the past most fresh water bodies have

long existed at a slightly acidic pH level.

However, pH is not precisely p[H], but takes into account an activity factor. This represents

the tendency of hydrogen ions to interact with other components of the solution, which affects

among other things the electrical potential read using a pH meter. As a result, pH can be

affected by the ionic strength of a solution – for example, the pH of a 0.05 M potassium

hydrogen phthalate solution can vary by as much as 0.5 pH units as a function of added

potassium chloride, even though the added salt is neither acidic nor basic.

Hydrogen ion activity coefficients cannot be measured directly by any thermodynamically

sound method, so they are based on theoretical calculations. Therefore the pH scale is defined

in practice as traceable to a set of standard solutions whose pH is established by international

agreement. Primary pH standard values are determined by the Harned cell, a hydrogen gas

electrode, using the Bates–Guggenheim Convention.

History

The concept of p[H] was first introduced by Danish chemist Søren Peder Lauritz Sørensen at

the Carlsberg Laboratory in 1909 and revised to the modern pH in 1924 after it became

apparent that electromotive force (emf) in cells depended on activity rather than concentration

of hydrogen ions. In the first papers, the notation had the H as a subscript to the lower case p,

like so: pH.

It is unknown what the exact definition of 'p' in pH is. A common definition often used in

schools is "percentage". However some references suggest the p stands for “Power”, others

refer to the German word “Potenz” (meaning power in German), still others refer to

“potential”. Jens Norby published a paper in 2000 arguing that p is a constant and stands for

“negative logarithm”; H then stands for Hydrogen. According to the Carlsberg Foundation pH

stands for "power of hydrogen". Other suggestions that have surfaced over the years are that

the p stands for puissance (also meaning power but then the Carlsberg Laboratory was French

speaking) or that pH stands for the Latin terms pondus Hydrogenii or potentia hydrogenii. It is

also suggested that Sørensen used the letters p and q (commonly paired letters in

mathematics) simply to label the test solution (p) and the reference solution (q).

Lab.13. pH measurement

Mathematical definition

pH is defined as a negative decimal logarithm of the hydrogen ion activity in a solution.

(1)

where aH is the activity of hydrogen ions in units of mol/L (molar concentration).

Activity has a sense of concentration, however activity is always less than the concentration

and is defined as a concentration (mol/L) of an ion multiplied by activity coefficient. The

activity coefficient is a number between 0 and 1 and it depends on many parameters of a

solution, such as nature of ion, ion force, temperature etc.

For a strong electrolyte activity of an ion approaches it concentration in diluted solutions.

Activity can be measured experimentally by means of an ion-selective electrode which

responds, according to the Nernst equation, to hydrogen ion activity.

pH is commonly measured by means of a glass electrode connected to a milli-voltmeter with

very high input impedance which measures the potential difference, or electromotive force, E,

between an electrode sensitive to the hydrogen ion activity and a reference electrode, such as

a calomel electrode or a silver chloride electrode.

Quite often glass electrode is combined with the reference electrode and a temperature sensor

in one body. The glass electrode relatively good (95 - 99.9%) follows the Nernst equation:

(2,3)

where E is a measured potential , E0 is the standard electrode potential, that is, the electrode

potential for the standard state in which the activity is one. R is the gas constant, T is the

temperature in Kelvins, F is the Faraday constant and n is the number of electrons transferred

(ion charge), one in this instance. The electrode potential, E, is proportional to the logarithm

of the hydrogen ion activity (or concentration at first approximation).

Lab.13. pH measurement

This definition, by itself, is wholly impractical, because the hydrogen ion activity is the

product of the concentration and an activity coefficient. To get proper results, the electrode

must be calibrated using standard solutions of known activity.

Note that, the pH of a solution is temperature-dependent.

Measurement of extremely low pH values, such as some very acidic mine waters, requires

special procedures. Calibration of the electrode in such cases can be done with standard

solutions of concentrated sulfuric acid, whose pH values can be calculated with using Pitzer

parameters to calculate activity coefficients.

pH is an example of an acidity function. Hydrogen ion concentrations can be measured in

non-aqueous solvents, but this leads, in effect, to a different acidity function, because the

standard state for a non-aqueous solvent is different from the standard state for water.

Superacids are a class of non-aqueous acids for which the Hammett acidity function, H0, has

been developed.

pOH

pOH is sometimes used as a measure of the concentration of hydroxide ions, OH−, or

alkalinity. pOH is not measured independently, but is derived from pH. The concentration of

hydroxide ions in water is related to the concentration of hydrogen ions by

[OH−] = KW /[H+] (4)

where KW is the self-ionisation constant of water. Taking cologarithms

pOH = pKW − pH. (5)

So, at room temperature pOH ≈ 14 − pH. However this relationship is not strictly valid in

other circumstances, such as in measurements of soil alkalinity.

Application

Pure (neutral) water has a pH around 7 at 25 °C (77 °F); this value varies with temperature.

When an acid is dissolved in water the pH will be less than 7 (if at 25 °C (77 °F)) and when a

base, or alkali is dissolved in water the pH will be greater than 7 (if at 25 °C (77 °F)). A

Lab.13. pH measurement

solution of a strong acid, such as hydrochloric acid, at concentration 1 mol dm−3 has a pH of

0.

A solution of a strong alkali, such as sodium hydroxide, at concentration 1 mol dm−3 has a pH

of 14. Thus, measured pH values will mostly lie in the range 0 to 14. Since pH is a

logarithmic scale a difference of one pH unit is equivalent to a tenfold difference in hydrogen

ion concentration.

Because the glass electrode (and other ion selective electrodes) responds to activity, the

electrode should be calibrated in a medium similar to the one being investigated. For instance,

if one wishes to measure the pH of a seawater sample, the electrode should be calibrated in a

solution resembling seawater in its chemical composition.

An approximate measure of pH may be obtained by using a pH indicator. A pH indicator is a

substance that changes color around a particular pH value. It is a weak acid or weak base and

the color change occurs around 1 pH unit either side of its acid dissociation constant, or pKa,

value. For example, the naturally occurring indicator litmus is red in acidic solutions (pH<7 at

25 °C (77 °F)) and blue in alkaline (pH>7 at 25 °C (77 °F)) solutions.

Universal indicator consists of a mixture of indicators such that there is a continuous color

change from about pH 2 to pH 10. Universal indicator paper is simple paper that has been

impregnated with universal indicator.

A solution whose pH is 7 (at 25 °C (77 °F)) is said to be neutral, that is, it is neither acidic nor

basic. Water is subjected to a self-ionization process.

H2O H+ + OH−

The dissociation constant, KW, has a value of about 10−14, so in neutral solution of a salt both

the hydrogen ion concentration and hydroxide ion concentration are about 10−7 mol dm−3.

The pH of pure water decreases with increasing temperatures. For example, the pH of pure

water at 50 °C is 6.55. Note, however, that water that has been exposed to air is mildly acidic.

This is because water absorbs carbon dioxide from the air, which is then slowly converted

into carbonic acid, which dissociates to liberate hydrogen ions:

Lab.13. pH measurement

CO2 + H2O H2CO3 HCO3− + H+

Calculation of pH for weak and strong acids

In the case of a strong acid, there is complete dissociation, so the pH is simply equal to minus

the logarithm of the acid concentration.

For example, a 0.01 molar solution of hydrochloric acid has a pH = −log(0.01), that is, pH =

2.

The pH of a solution of a weak acid may be calculated by means of an ICE table.

For acids with a pKa value greater than about 2,

pH = ½ ( pKa − log c0) (6)

where c0 is the concentration of the acid. This is equivalent to Burrows' weak acid pH

equation

(7)

A more general method is as follows. Consider the case of dissolving a weak acid, HA, in

water. First write down the equilibrium expression.

HA A− + H+

The equilibrium constant for this reaction is specified by

(8)

The analytical concentration of the two reagents, CA for [A−] and CH for [H+] must be equal to

the sum of concentrations of those species that contain the reagents. CH is the concentration of

added mineral acid.

CA = [A−] + Ka[A−][H+] (9)

Lab.13. pH measurement

CH = [H+] + Ka[A−][H+] (10)

From the equation (8)

(11)

Substitution of this expression into the equation (9) gives

(12)

This simplifies to a quadratic equation in the hydrogen ion concentration

(13)

Solution of this equation gives [H+] and hence pH.

This method can also be used for polyprotic acids. For example, for the diprotic acid oxalic

acid, writing A2− for the oxalate ion,

CA = [A2−] + β1[A2−][H+] + β2[A

2−][H+]2 (14)

CH = [H+] + β1[A2−][H+] + 2β2[A

2−][H+]2 (15)

where β1 and β2 are cumulative protonation constants.

Following the same procedure of substituting from the first equation into the second, a cubic

equation in [H+] results. In general, the degree of the equation is one more than the number of

ionisable protons. The solution of these equations can be obtained relatively easily with the

aid of a spreadsheet such as EXCEL or Origin. The pH always has an amount of fractional

figures equal to the amount of significant figures of the concentration.

Lab.13. pH measurement

pH in nature

pH-dependent plant pigments that can be used as pH indicators occur in many plants,

including hibiscus, marigold, red cabbage (anthocyanin), and red wine.

Living systems

The pH of different cellular compartments, body fluids, and organs is usually tightly regulated

in a process called acid-base homeostasis.

Tab. pH in living systems

Compartment pH pH

Gastric acid 1

Lysosomes 4.5

Granules of chromaffin cells 5.5

Human skin 5.5

Urine 6.0

Neutral H2O at 37 °C 6.81

Cytosol 7.2

Cerebrospinal fluid (CSF) 7.3

Blood 7.34–7.45

Mitochondrial matrix 7.5

Pancreas secretions 8.1

Lab.13. pH measurement

The pH of blood is usually slightly basic with a value of pH 7.365. This value is often referred

to as physiological pH in biology and medicine.

Plaque can create a local acidic environment that can result in tooth decay by

demineralisation.

Enzymes and other proteins have an optimum pH range and can become inactivated or

denatured outside this range.

The most common disorder in acid-base homeostasis is acidosis, which means an acid

overload in the body, generally defined by pH falling below 7.35

In the blood, pH can be estimated from known base excess (be) and bicarbonate concentration

(HCO3) by the following equation:

(16)

pH meter

A pH meter is an electronic instrument measuring the pH (acidity or alkalinity) of a liquid

(though special probes are sometimes used to measure the pH of semi-solid substances). A

typical pH meter consists of a special measuring probe (a glass electrode) connected to an

electronic meter that measures and displays the pH reading.

The probe

The pH probe measures pH as the activity of hydrogen ions surrounding a thin-walled glass

bulb at its tip. The probe produces a small voltage (about 0.06 volt per pH unit) that is

measured and displayed as pH units by the meter. For more information about pH probes, see

glass electrode.

pH meter Calibration and use

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 e.m.f. 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 and pH 10 are

acceptable. The pH meter has one control (calibrate) to set the meter reading equal to the

Lab.13. pH measurement

value of the first standard buffer and a second control (slope) 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, which can be obtained from a variety of suppliers, usually state how

the buffer value changes with temperature.

The calibration process correlates the voltage produced by the probe (approximately 0.06

volts per pH unit) with the pH scale. After each single measurement, the probe is rinsed with

distilled water or deionized water to remove any traces of the solution being measured, blotted

with a clean tissue to absorb any remaining water which could dilute the sample and thus alter

the reading, and then quickly immersed in another solution.

Storage conditions of the glass probes

When not in use, the glass probe tip must be kept wet at all times to avoid the pH sensing

membrane dehydration and the subsequent dysfunction of the electrode.

A glass electrode alone (i.e., without combined reference electrode) is typically stored

immersed in an acidic solution of around pH 3.0. In an emergency, acidified tap water can be

used, but distilled or deionised water must never be used for longer-term probe storage as the

relatively ionless water "sucks" ions out of the probe membrane through diffusion, which

degrades it.

Combined electrodes (glass membrane + reference electrode) are better stored immersed in

the bridge electrolyte (often KCl 3 M) to avoid the diffusion of the electrolyte (KCl) out of

the liquid junction.

Types of pH meters

pH meters range from simple and inexpensive pen-like devices to complex and expensive

laboratory instruments with computer interfaces and several inputs for indicatorerature

measurements be entered to adjust for the slight variation in pH caused by temperature.

Specialty meters and probes are available for use in special applications, harsh environments,

etc.

Lab.13. pH measurement

Glass electrode

A glass electrode is a type of ion-selective electrode made of a doped glass membrane that is

sensitive to a specific ion. It is an important part of the instrumentation for chemical analysis

and physico-chemical studies. In modern practice, widely used membranous ion-selective

electrodes (ISE, including glasses) are part of a galvanic cell. The electric potential of the

electrode system in solution is sensitive to changes in the content of a certain type of ions,

which is reflected in the dependence of the electromotive force (EMF) of galvanic element

concentrations of these ions.

History

F. Haber and Z. Klemensiewicz [Polish scientist] publicized on January 28, 1909 results of

their research on the glass electrode in The Society of Chemistry in Karlsruhe (first

publication — The Journal of Physical Chemistry by W. Ostwald and J. H. van 't Hoff) —

1909).

Legend:

1. a sensing part of electrode, a bulb made from

specific glass

2. internal electrode, usually silver chloride

electrode or calomel electrode

3. internal solution, usually 0.1M HCl for pH

electrodes or 0.1M MeCl for pMe electrodes

4. sometimes electrode contain small amount of

AgCl precipitate inside the glass electrode

5. reference electrode, usually the same type as 4

6. junction with studied solution, usually made from

ceramics or capillary with asbestos or quartz

fiber.

7. reference junction

8. body of electrode, made from non-conductive

glass or plastics.

Lab.13. pH measurement

Applications

Glass electrodes are commonly used for pH measurements. There are also specialized ion

sensitive glass electrodes used for determination of concentration of lithium, sodium,

ammonium, and other ions. Glass electrodes have been utilized in a wide range of

applications — from pure research, control of industrial processes, to analyze foods,

cosmetics and comparison of indicators of the environment and environmental regulations: a

microelectrode measurements of membrane electrical potential of a biological cell, analysis of

soil acidity, etc.

Saturated calomel electrode

The Saturated calomel electrode (SCE) is a reference electrode based on the reaction between

elemental mercury and mercury(I) chloride. The aqueous phase in contact with the mercury

and the mercury(I) chloride (Hg2Cl2, "calomel") is a saturated solution of potassium chloride

in water. The electrode is normally linked via a porous frit to the solution in which the other

electrode is immersed. This porous frit is a salt bridge.

In cell notation the electrode is written as:

Theory of operation

The electrode is based on the redox reaction

The Nernst equation for this reaction is

where E0 is the standard electrode potential for the reaction and aHg is the activity for the

mercury cation (the activity for a liquid of 1 Molar is 1). This activity can be found from the

solubility product of the reaction

Lab.13. pH measurement

By replacing the activity in the Nernst equation with the value in the solubility equation, we

get

The only variable in this equation is the activity (or concentration) of the chloride anion. But

since the inner solution is saturated with potassium chloride, this activity is fixed by the

solubility of potassium chloride. When saturated the redox potential of the calomel electrode

is +0.2444 V vs. SHE at 25 °C, but slightly higher when the chloride solution is less than

saturated. For example, a 3.5M KCl electrolyte solution increases the reference potential to

+0.250 V vs. SHE at 25 °C, and a 0.1 M solution to +0.3356 V at the same temperature.

Application

The SCE is used in pH measurement, cyclic voltammetry and general aqueous

electrochemistry.

This electrode and the silver/silver chloride reference electrode work in the same way. In both

electrodes, the activity of the metal ion is fixed by the solubility of the metal salt.

The calomel electrode contains mercury, which poses much greater health hazards than the

silver metal used in the Ag/AgCl electrode.

Standard hydrogen electrode

The standard hydrogen electrode (abbreviated SHE), also called normal hydrogen electrode

(NHE), is a redox electrode which forms the basis of the thermodynamic scale of oxidation-

reduction potentials. Its absolute electrode potential is estimated to be 4.44 ± 0.02 V at 25 °C,

but to form a basis for comparison with all other electrode reactions, hydrogen's standard

electrode potential (E0) is declared to be zero at all temperatures. Potentials of any other

electrodes are compared with that of the standard hydrogen electrode at the same temperature.

Hydrogen electrode is based on the redox half cell:

Lab.13. pH measurement

2H+(aq) + 2e- → H2(g)

This redox reaction occurs at platinized platinum electrode. The electrode is dipped in an

acidic solution and pure hydrogen gas is bubbled through it. The concentration of both the

reduced form and oxidized form is maintained at unity. That implies that the pressure of

hydrogen gas is 1 bar and the activity of hydrogen ions in the solution is 1 molar. The activity

of hydrogen ions is their effective concentration, which is equal to the formal concentration

times the activity coefficient. Activity coefficients are close to 1.00 for very dilute water

solutions, but are usually lower for more concentrated solutions. The Nernst equation should

be written as:

or

where:

aH+ is the activity of the hydrogen ions, aH+=fH

+ CH+ /C0

pH2 is the partial pressure of the hydrogen gas, in pascals, Pa

R is the universal gas constant

T is the temperature, in kelvins

F is the Faraday constant (the charge per a mole of electrons), equal to 9.6485309*104

C mol-1

p0 is the standard pressure 105 in Pa

Lab.13. pH measurement

On line pH calculator

http://www.webqc.org/phsolver.php

For solutions with ionic strengths of 0,1 M or less, the electrolyte effect is independent

of the kind of ions and dependent only on the ionic strength.

Standard hydrogen electrode scheme:

1. platinized platinum electrode

2. hydrogen gas

3. acid solution with an activity of

H+=1 mol/l

4. hydroseal for prevention of oxygen

interference

5. reservoir via which the second half-

element of the galvanic cell should be

attached

Lab.13. pH measurement

EXPERIMENTAL PART

Equipment

pH meter;

solutions of different acids and bases

solutions of acetic acid, ammonia

solution different salts solution

Procedure

1.Measurement of pH of strong acids

Step 1.

Prepare 5 solutions of H2SO4 (concentration from 1.0 to 0.0001 M, according to data in Table

1) in the beakers of capacity 25 mL.

Step 2.

Insert the clean and dry electrode of the pH meter to selected solution (e.g. 0.0001M).

Step 3.

Record the values of pH in Table 1.

Step 4.

Clean the electrode of the pH meter with distilled water and dry it with tissue.

Step 5.

Repeat measurement procedure (steps 2-4) to remaining acid solution(e.g.0.001, 0.01, 0.1 and

1.0M)

Table 1. Determined and calculated data for H2SO4.

Concentration of H2SO4 [mol/L] 1.0 0.1 0.01 0.001 0.0001

pH calculated from equation: pH = - log [H+]

pH measured experimentally

Lab.13. pH measurement

Step 6.

Prepare 5 solutions of HCl (concentration from 0.0001 to 1.0 M, according to data in Table 2)

in the beakers of capacity 25 mL.

Step 7

Repeat steps 2-6. Report data in Table 2.

Table 2.Determined and calculated data for HCl.

Concentration of HCl [mol/L] 1.0 0.1 0.01 0.001 0.0001

pH calculated from equation: pH = - log [H+]

pH measured experimentally

2.Measurement pH of weak acid.

Step 8.

Prepare 5 solutions of acetic acid in the beakers of capacity 25 mL(concentration from 0.0001

to 1.0 M according to data in Table 3).

Step 9.

Repeat steps from 2-6. Record the values of pH in Table 3.

Table 3. Determined and calculated data for acetic acid.

Concentration of acetic acid[mol/L] 1.0 0.1 0.01 0.001 0.0001

pH calculated from equation: pH = - log [H+]

acida CKH ][

pH measured experimentally

For calculation pH of weak acid apply the formula:

acida CKH ][

where Ka = 1.7 ∙10-5

Lab.13. pH measurement

3.Measurement pH of weak base.

Step 10.

Prepare 5 solutions of ammonia in the beakers of capacity 25 mL (concentration from 0.0001

to 1 M, according to data in Table 4)

Step. 11.

Repeat steps 2-6. Record the values of pH in Table 4.

Table 4. Determined and calculated data for ammonia.

Concentration of ammonia [mol/L] 1.0 0.1 0.01 0.001 0.0001

pH calculated from equation: pH = - log [H+]

pH = 14-pOH; pOH = -log [OH-];

baseb CKOH ][

pH measured experimentally

For calculation pH of weak base apply the formula:

baseb CKOH ][

where Kb = 1.7 ∙10-5

4.Measurement of pH of different salts

Step 12.

Prepare 0.1M solutions of salts from Table 5 in the beakers of capacity 25 mL.

Step 13.

Repeat steps 2-6. Record the values of pH in Table 5.

Table 5. Determined and calculated data for salts

Salt Concentration pHexp pHcalc

NH4Cl

pH = 1 /2 pKw – 1 /2 pKb – 1/2 lg C

CH3COONa

pH = 1/2 pKw + 1/2 pKA +1/2 log C

Calculate pH of the salts solutions.

Lab.13. pH measurement

For salts with one strong and one weak component the best way of pH calculation it to treat

conjugate acid (or base) as the only source of H+ (or OH-) ions. For example if we have

solution of salt of weak acid (with dissociation constant Ka) and strong base, reaction of

hydrolysis is

and the equilibrium is described by conjugate base dissociation constant

where

Starting from these equations we can calculate pOH and pH of the solution using method and

assumptions shown for weak acid and base. Exactly the same approach can be used for salt of

strong acid and weak base - just using the Ka constant for acid conjugate with weak base.

In the case of salt of weak acid and weak base situation is more complicated, but sometimes

we can get pretty good results assuming similar degree of both hydrolysis processes (as seen

above).

Let's say we have a solution of AB salt of weak acid and weak base of concentration C and

dissociation constants Ka and Kb. Our first assumption is that the hydrolysis is not too strong,

so that in the equilibrium [A-]=[B+]=C. If so, equations for Ka and Kb take forms

We will solve them for [H+] and [OH-]:

Lab.13. pH measurement

Multiplying:

and rearranging:

Please note that these Ka and Kb values are not related by KaKb=Kw, as they describe different

substances.

Now it is time for the second assumption - that degree of both hydrolysis reactions is similar,

so [HA]=[BOH]:

It gives concentration of HA which can be used for [H+] calculation. If equation above is

insert into

It is obtain:

or

Lab.13. pH measurement

5.Determination of Buffer Capacity

A. Determination of Buffer Capacity with HCl

Step 15.

Transfer 5 mL of the test solution [ BUFFER pH 4 AND BUFFER pH 10] to another beaker

(use a graduated cylinder) and set aside.

Step 16.

Using the automatic dispenser add 0.25 mL of 0.5 M HCl to the test solution. (One pump of

the dispenser).

Step 17.

Mix well with the glass stirring rod and then measure the pH of the solution. Record data in

Table 6.

Step 18.

Discard the solution from the beaker into the waste bottle. Thoroughly rinse the beaker with

DI water and then dry the beaker.

B. Determination of the Buffer Capacity with NaOH

Step 19.

Transfer 5 mL of the test solution [ BUFFER pH 4 AND BUFFER pH 10] to another beaker

(use a graduated cylinder) and set aside.

Step 20.

To the remaining 5 mL portion of the test solution add 0.25 mL of 0.5 M NaOH (one pump

of the automatic dispenser).

Step 21

Mix well with the glass stirring rod and then measure the pH. Record data in Table 6

Lab.13. pH measurement

Step 22.

Discard the solution in the waste bottle. Thoroughly rinse the beaker with water and then dry

the beaker.

Table 6. Buffer capacity data

Solution Initial pH pH after H+

Addition

pH after OH-

Addition

Buffer

Capacity

(mmol

H+/∆pH)

Buffer

Capacity

(mmol OH-

/∆pH)

Buffer pH = 4.0

Buffer pH = 4.0

Buffer pH = 10.0

Buffer pH = 10.0

Calculations to be shown in your lab notebook:

Calculate the buffer capacity (mmoles of acid or base added/change in pH) Buffer solutions as

mmoles H+/∆pH and as mmoles OH-/∆pH.