Chapter 12  pH and Alkalinity              return to Table of Contents                            


                 Pure water ionizes slightly into hydronium and hydroxide ions.  It has been found by experiment that one liter of pure water contains only one ten millionth of a mole of hydronium ions and one ten millionth of a mole of hydroxide ions.  A substance which contains more hydronium ions than hydroxide ions is acidic; a substance which contains more hydroxide ions than hydronium ions is basic; and a substance such as pure water which contains an equal number of hydronium ions and hydroxide ions is neutral.

                 The pH scale indicates the hydronium ion concentration.  Another related scale (rarely seen or used), the pOH scale, is used to indicate the hydroxide ion (OH-) concentration1.  If you know the pH, simply subtract the pH value from 14 to change to the pOH scale.

             pH is a number that exactly describes the concentration of acidity of a solution.  The pH scale was developed in 1909 by a scientist named Sorensen and at the same time devised the symbol pH.  An acid is defined as a material having ionized (or free) hydrogen ions while a material which has ionized (or free) hydroxyl ions is described as basic.  pH is directly related to the ratio of  H+ to OH- ions.  If the H+  is greater than the OH- then the material is acid.  If the OH- is more numerous than the H+, then the material is a base.  If equal amounts of ions are present, the material is neutral.

                 Probably the most fundamental concept in the approach to understanding water chemistry is the acidity-alkalinity relationship.  The first step in grasping this is an understanding of the dissociation of the water molecule into hydrogen ions and hydroxyl ions.  Though general discussions use the hydrogen ion (H+), one should remember that it actually exists in a hydrated form, H3O+.

                 Since the dissociation constant is so very small, 10 - 14, at neutrality where there are the same number of hydrogen and hydroxyl ions there are only 10 - 7 moles per liter of each.  This is equal to only 10 - 4 millimoles per liter, corresponding to an actual concentration of only 0.0001 mg/L H+ ion, equivalent to 0.005 mg/L as CaCO3.  Because we are dealing with such small numbers in the dissociation of water into its ions, it is more convenient to substitute an expression involving the power of 10.  This expression is defined as pH.

                 The dissociation constant, K, changes with temperature, and this must be taken into account in interpreting data involving  H+ and OH- ions.  For example, many water treatment operations are carried out at high temperatures, and samples from the system are usually cooled prior to analysis.  The H+ and OH- concentrations measured on the cooled sample, even though different from those in the hot system, are usually used for control purposes.  But a physical chemist needing to know the conditions prevailing in the hot system must use the dissociation constant for the temperature in that system.

                 The hydrogen ion concentration can be measured with a pH meter.  It can also be titrated when the concentration becomes large enough to be detectable by chemical analysis.  Since pH is a logarithmic function, the hydrogen ion concentration increases by a factor of 10 for each unit of pH reduction.  The pH scale is temperature dependent with the neutral point falling in pH as the temperature rises.  Newer pH probes have a temperature sensor and electronics to take care of the temperature factor.  These probes are more complex and require more care and the special membrane can withstand temperatures only within a range.

                 When the pH drops below approximately 5, the hydrogen ion begins to reach mg/L levels, concentrated enough to be determined by titration, using the correct organic dye indicator.  The chemical indicator originally selected by the water chemist for this purpose was methyl orange, changing color at pH 4.2 to 4.4 SU (standard units).  The color change of this indicator was so subtle - - orange on the alkaline side to salmon pink on the acid side - - that researchers looked for a substitute to give a more pronounced color change.  The one they developed produces a blue color on the alkaline side and red on the acid side, with gray at the endpoint.  Even though this special indicator has replaced methyl orange, the water chemist still defines alkalinity as methyl orange alkalinity (“M” alkalinity) which exists above the approximate pH range of 4.2 to 4.4 SU.  “M” acidity is strong mineral acidity which exists below this pH range.

                 To the theoretical chemist, a pH of 7 is considered neutral; to the water chemist, a pH of 7 means very little.  He must also know how much total alkalinity and how much free or combined CO2 is present.  For the water chemist the dividing point between acidity and alkalinity is not pH 7.0, but rather the “M” alkalinity endpoint, corresponding to a pH of approximately 4.4.

                 The water chemist is also concerned with “P” alkalinity (phenolphthalein alkalinity) which exists when the pH is over a range of 8.2 to 8.4, corresponding to the change in phenolphthalein indicator from a colorless condition below 8.2 to pink or red above 8.4.  In most natural water supplies, the pH is less than 8.2, so there is no “P” alkalinity.  Very few natural waters have a pH below about 5.0, so it is seldom that strong mineral acids are found in fresh water.  The pH range between the “M” endpoint and the “P” endpoint defines the alkaline range where bicarbonate alkalinity exists and weak acids may be present, the most prominent of which is carbonic acid -- carbon dioxide in solution.

                 Alkalinity is the acid neutralizing ability of the sample.  If a sample has no alkalinity, such as reagent water, addition of one drop of concentrated hydrochloric acid to a 100 mL portion of the water will drop the pH from 7.0 to about 4 SU.  If the reagent water is replaced with a solution containing alkalinity, adding that drop of acid may have little noticeable effect on the pH.  Solutions can be prepared with a known amount of alkalinity from pure chemicals.  These are called Buffer  solutions.  The pH of the buffer solution depends on the chemicals chosen to make it. 

                 Total alkalinity analysis is performed by measuring the amount of acid needed to reduce the sample’s pH to a certain value.  Alkalinity is expressed in terms of the equivalent amount of calcium carbonate neutralized, rather than attempt to determine the actual cause of alkalinity.  The endpoint of the titration is indicated by either pH electrode or by a visual color change.  The alkalinity of a water sample is a measure of the water’s capacity to neutralize acids.  In natural and treated waters, alkalinity is the result of bicarbonates, carbonates, and hydroxides of the metals of calcium, magnesium, and sodium.

                 Many of the chemicals used in water treatment, such as alum, chlorine, or lime, cause changes in alkalinity.  The alkalinity determination is a useful tool in calculating chemical dosages needed in coagulation and water softening.  Alkalinity must also be known to calculate corrosivity and to estimate the carbonate hardness of water.  Alkalinity is usually expressed in terms of Calcium Carbonate (CaCO)3 Equivalence.

                 There are five alkalinity conditions possible in a water sample: (1) bicarbonate alone, (2) bicarbonate and carbonate, (3) carbonate alone, (4) carbonate and hydroxide, and (5) hydroxide alone.


Alkalinity, mg/L as CaCO3

Titration Result




P = 0




P < T/2

T - 2P



P = T/2 




P > T/2


2T - 2P

2P - T

P = T 




where P = phenolphthalein alkalinity and T = total alkalinity.

Care needs to be taken to insure that the pH measurement of a sample is a true measure of the pH. If the sample is allowed to stand exposed to the atmosphere for much time, the Carbon Dioxide will escape, causing the pH to rise. This emphasizes the need for in situ measuring.

                 Standard Methods  in section 4500 - H+ B has the recipes for preparing many buffer solutions of well defined pH standard solutions.  These are used as pH standards and are also commercially available as both prepared solutions and as powdered mixtures, sometimes called ‘pillows’, to be diluted to a specified volume with reagent water.  Calibrating a pH meter is usually accomplished with these solutions.

                A pH can be measured colorimetrically or electrometrically.  The colorimetric methods are susceptible to interferences from turbidity, salinity, oxidants and colloidal matter.  The color standards fade over time.  The indicator solutions cover narrow portions of the pH scale requiring several sets of standards and indicators.

                 For quick, general determination of a solution’s pH, strips of specially impregnated paper are often used.  These indicator  strips change color when a certain pH level is reached.  Phenolphthalein is a common indicator that is colorless in solutions below 8.3 SU and bright pink-red when greater than 8.3 SU pH.  These strips are available in wide variety of formulations to provide useful versatility aids for the analyst.

                 Modern determination of pH relies upon pH meters.  The electrometric method is quite accurate and is the most commonly used.  The electrode response must be calibrated each day and sometimes for each sample.  The manufacturer’s directions should always be consulted for specific calibration instructions.  The most correct calibration procedure is to choose buffer standards which bracket the suspected pH of the sample.  The pH strips are quite useful in this situation.  Corroboration and reliability checks validate the analysis when using completely different ways to arrive at the same number.  Or, show a problem exists in the test equipment.  Read what is there, not what you think is there or want to see.

Calibration of the probe and the meter is a system operation. Single point, two point, grab sample are types of procedures

                The electrometric pH probe should be thought of as a battery whose voltage changes as the pH of the solution in which it is inserted changes.  Electrodes, as batteries, have limited shelf lives. The probe consists of two parts (in fact, some pH measurements are made with two separate probes):  first, the hydrogen sensitive glass bulb and second, the reference electrode.  The special glass of the bulb has the ability to absorb H+ ions on the outside of the thin glass in proportion to the concentration in the bulk solution, and the charge on the inside influences the charges on the inside surface, which determines the electrical potential on the internal wire.  This bulb is a half-cell and requires a companion to function.  A reference electrode surrounded by a solution provides the other half-cell.  Manufacturers are continually competing to provide better, more functional equipment, so the actual physical description of a pH probe will differ somewhat, but the principle remains the same.

Vigorous stirring brings a sample, buffer or rinse solution to the measuring surface more quickly and will improve speed of response. Care must be taken to keep the electrode's relatively thin, delicate surface from striking a surface and getting scratched, gouged or broken. As a rinse solution, use a part of the next sample or buffer which will be measured. This action aids minimizing contaminating the next sample from the previous sample. After exposure to a sample, buffer or rinse solution, carryover/contamination may be reduced to a minimum by blotting -never wiping- the electrode with a clean, non-abrasive paper or a clean cloth towel.

Electrodes age. You'll notice aging by the longer time to get the reading or span errors or both. Using electrodes in hot and/or hi/lo pH liquids will degrade the unit quicker.

The respone time of solutions differ. A reading with a new electrode in buffer with settle down in 10 to 15 seconds. Samples may take longer to stablize to a reading.

         Temperature is important. A sample at one temperature will have a different pH when the temperature is different. Probably the main reason pH readings are always taken in situ (right then, on site, not later). The output of the probe also varies with temperature.

        pH is the scale used to express the quantity of acid present in a water solution.  The scale ranges from 0 to 14.0 is very acidic, equivalent to a 1 M solution of hydrochloric acid.  14 is very basic, equivalent to a 1 M solution of sodium hydroxide.  Each unit difference in the scale is a concentration change by a factor of 10.  Thus, a pH 6 SU has one-tenth the concentration of hydrogen ions as a pH 5 SU (Standard Units).

                 Always pour acid into water.  Always pour bases into water.  Why?  The other way will likely splash back.  Why?  Pouring water into the acid provides many more opportunities for water to be encapsulated, surrounded by the acid.  As the reaction is exothermic and the steam generated is also contained for a while (things happen very quickly in this scenario) the pressure increases until a release /fracture/weakpoint in the boundary is found.  Acid poured into water, on the other hand, disperses while the reaction and the steam generated is not contained by acid.  Either way should be done with caution.

                 Two tests done at most wastewater treatment plant labs, pH and TSS, deserve to have their bias and precision monitored, just as you monitor precision and bias for the BOD test by doing the glucose-glutamic acid (G/GA) test.  For TSS, the lab can prepare a suspension of 30 mg/L by suspending 30 mg of 20 micrometer cellulose, or diatomaceous earth, in a liter of distilled water.  Because the suspension is biodegradable, it should be kept in the 4º C refrigerator.  Likewise, the lab should analyze a pH standard every week or so.  If the lab calibrates at 4.0 and 7.0, the check standard could be a 5.0 or 6.0, and likewise if they calibrate at 7.0 and 10.0.  If they calibrate at 4.0, 7.0, and 10.0, the check standard should be close to the samples normally analyzed in the lab.  Very important to monitoring performance in the lab is the use of control charts for ALL STANDARDS THAT ARE ANALYZED ON A REGULAR BASIS (if a standard is analyzed infrequently ... e.g., a potassium permanganate standard for the chlorine test ... control charts won't do much good).


pH Procedure Checklist (4500 - H+ B,  SM18)



          ___  1.  Allow meter to warm up, then follow manufacturer’s

instructions to calibrate meter at pH 4.00 and 7.00 SU.

(Meter should be calibrated to a range no greater than 3 SU [2 point method] and bracketing the expected test values)

         ___  2.  Place a portion of the sample to be measured in a disposable

plastic cup.

          ___  3.  Measure pH with the meter and record value on lab sheet.

          ___  4.  Record the temperature of the sample on the benchsheet.

          ___  5.  Perform a duplicate analysis, calculate RPD  and update the

control chart.


Total Alkalinity Checklist (2320 B, SM18)



          ___  1.  1 N  Hydrochloric acid:   Add 83 mL of concentrated

hydrochloric acid to about 100 mL reagent water, then dilute to 1000 mL.  This solution must be standardized.

         ___  2.   0.1 N Sulfuric acid:  add 2.8 mL concentrated sulfuric acid to

about 100  mL reagent water.  Allow to cool, then dilute to

1000 mL.  This solution must be standardized.

          ___  3.  0.02 N acid:  Dilute exactly 200 mL of the standardized acid

solution to exactly 1000 mL reagent water.  (Divide the actual normality of the acid solution by  5,  .

          ___  4.   0.1 N sodium hydroxide:  Dissolve 4.0 g sodium hydroxide

(NaOH) pellets in 50 mL reagent water.  After cooling, dilute to 1000 mL with reagent water.  This solution must be standardized immediately before use.

          ___  5.   Potassium hydrogen phthalate (KHP), acidimetric primary

standard grade:  dry in oven at 104º C for at least an hour then store in desiccator prior to use.

         ___  6.   Sodium carbonate:  dry in a 250º C oven for at least 4 hours

and cool in a desiccator immediately prior to use.

         ___  7.   Bromocresol green indicator solution, pH 4.5 SU:  dissolve

100 mg bromocresol green sodium salt in 100 mL reagent


         ___  8.   Phenolphthalein indicator solution:  dissolve 0.5 g

phenolphthalein disodium salt in water and dilute to 100 mL.2


                           Standardization with KHP (Potassium hydrogen phthalate)


          ___  1.   Dissolve 408 mg potassium hydrogen phthalate in reagent

water and dilute to about 100 mL.  Record step.

          ___  2.   Add 2 to 3 drops of phenolphthalein indicator solution.

          ___  3.   Fill a 25 mL buret with the 0.1 N sodium hydroxide solution,

then titrate the KHP solution until it turns permanent red-pink.  Record the amount of sodium hydroxide solution used to 2 decimal places, ex. 21.00 mL.

         ___  4.   Calculate the normality of the sodium hydroxide solution to 3

decimal places using the following equation:

         ___  5.   Repeat the standardization and calculation from step 1.  If the

values agree within 1 %, use the average of the values for

the sodium hydroxide normality.  If not, repeat the

standardization until two values within 1 % are obtained.

          ___  6.   Using a volumetric pipet, add 20.00 mL of the acid solution to

a beaker with a magnetic stir bar.

          ___  7   Add 2 to 3 drops of phenolphthalein indicator solution.

          ___  8.  Titrate the acid solution with the standardized sodium

hydroxide solution until it turns permanent red-pink.  Record the volume of sodium hydroxide solution used to 2 decimal places, ex. 21.03 mL.

         ___  9.   Calculate the normality of the acid solution to 3 decimal

 places using the following equation:

         ___ 10.   Repeat the standardization and calculation of the normality of

the acid solution, beginning at step 7.    If the values agree within 1 %, use the average of the values for the sodium hydroxide normality.  If not, repeat the standardization until two values within 1 % are obtained.



                              Procedure for Alkalinity

         ___  1.   Pour 200 ml of sample into a beaker containing a magnetic

stir bar.

         ___  2.   Add 2 to 3 drops of bromocresol green indicator, or, insert pH

meter electrode into sample.

         ___  3.   Titrate with 0.1 N acid until the indicator changes color

permanently, or, until pH 4.5 is reached.   If more than 25 mL of acid is required, start over at step 1 with a smaller sample size.

         ___  4.   Record the volume of acid used.

         ___  5.   Calculate the alkalinity of the sample using the following


         ___  6.   Calculate the RPD on a duplicate sample and update the

control chart.

  On to Chapter 13  
 My question sets for self exam    
Top of this Document       a pH quiz errors?


1Trivia:  Human blood has a pH of 7.4. The gastric juices in your stomach have a pH of approximately 0.9 to aid in the digestion of food.

2If only phenolphthalein is available instead of the disodium salt, add 0.5 g of the phenolphthalein to 10 mL of water in an Erlenmeyer flask stirred on a magnetic stirrer.  Add 0.1 N sodium hydroxide dropwise slowly until a permanent pink color is obtained and everything is in solution.  Then just depolarize by dropwise addition of 0.1 N sulfuric acid.