pH Simplified

by Robert J. Joyce

What is pH ?

The pH notation is an index of Hydrogen's chemical activity in a solution.

pH is a Log Scale Unit of Measure, and is used to express the degree of acidity of a substance.

Values will range from pH 1 to pH 2 for strong acids, through pH 7 for neutral solutions such as ultra pure water, on up to values of pH 11 and higher for very strong bases like lye.

The pH notation is an index of Hydrogen's 
 
chemical activity in a solution.

The centimeter is a unit measure of length. The gram is a unit measure of weight. Similarly, pH is the unit measure we use to say how much free or active acid is in a substance. The pH scale goes from 0 to 14. A pH of 0 means a very high acid activity; a pH of 14 means a very low acid activity. In between these two extremes is a pH of 7. This is the pH of pure water.

Addition of a strong acid, such as sulfuric acid ( H2SO4 ) to water makes the resulting solution very high in active acid concentration. This is called an acidic solution. The addition of a strong base or alkali material, such as sodium hydroxide ( NaOH ), to water makes the resulting solution very low in active acid concentration. This is called a very basic or alkali solution. Water, which is neither very acidic nor very alkali, is said to be neutral. The pH scale is a quantitative way of expressing the active acid or alkali concentration of a solution.

Why pH is Important ?


The pH or acidity of a solution is important throughout all phases of chemistry and biochemistry.

In the Chemical Industry ...
The efficient production of nylon, as well as other modern fibers depends on rigid pH control.

In Biochemistry ...
The pH of our blood is normally controlled to within a few tenths of a pH unit by our body chemistry. If our blood pH changes as much as half a pH unit, serious illness will result. Proper skin pH is essential for a healthy complexion. The pH of one's stomach directly affects the digestive process.

In Agronomy ...
The pH of the soil regulates the availability of nutrients for plant growth, as well as the activity of soil bacteria. In alkaline soils ( pH 8 and above ) the amount of nitrogen, phosphorus, iron and other nutrients in solution become so low that special treatment is necessary to insure proper growth.

In Food Science ...
The efficient production of food products depends upon careful pH control. The proper curd size, uniformity, and structure of cottage cheese is directly related to the pH at cutting time. Yeast can ferment and leaven a dough only within certain pH limits. Jelly will not gel properly unless the pH is in the 3.5 region.

In the Pulp and Paper Industry ...
pH control is essential to the proper operation of bleaching plants and wet-end processes. Also, in order to conform with environmental protection regulations, the pH of wastewater from these plants must be controlled.

In Chemical Research and Engineering ...
Accurate pH measurement is necessary to the study of many chemical processes. The researcher needs to know the pH at which a chemical reaction proceeds at its fastest in order. to understand the reaction. The engineer uses the information to develop practical commercial processes.

In Environmental Research and Pollution Control ...
The pH of a river or lake is important in maintaining a proper ecological balance. The pH of the water directly affects the physiological functions and nutrient utilization by plant and animal life. Extremes in pH can reduce a lake to a lifeless, smelly bog.

Protecting our waterways requires constant monitoring of industrial effluent. Plating and metal finishing plants tend to produce acidic wastewater, as do mining operations, Chemical plants often have very alkaline wastewater. pH measurements are used as a guide to the proper neutralization of these plant wastes, as well as to monitor the final effluent quality. Occasionally, an acidic stream can be combined with an alkaline stream to produce a final stream, which is close to neutral. pH measurements assure the proper management of this cost saving technique.

More About pH ... ( for Those Who Really Want to Know ! )

To understand more about pH, we need to know more about the chemistry of water. A molecule of water is composed of one oxygen atom and two hydrogen atoms and looks something like this.

H = Hydrogen: O = Oxygen:

Water Molecule ( H2O ):

In pure water, most of the water molecules remain intact. However, a very small amount of them react with each other in the following manner.

H2O + H2O ===> H3O+ + OH

Water + Water ===> Hydronium Ion+ ( an Acid ) + Hydroxyl Ion ( a Base )

The Hydronium ion ( H3O+) is the chemical unit which accounts. for the acidic properties of a solution. The hydroxyl ion ( OH ) is the chemical which accounts for the basic or alkaline properties of a solution.

As you can see, when pure water reacts as described in Figure 2, it produces an equal amount of H3O+ and OH. Thus, it does not have an excess of either ion. It is therefore called a neutral solution.

If a strong acid, such as hydrochloric acid ( HCl ) is added to water, it reacts with some of the water molecules as follows:

HCl + H2O <=====> H3O+ + Cl

Thus, the addition of HCl to water increases the H3O+ or acid concentration of the resulting solution.

If a strong base, such as sodium hydroxide, is added to water, it ionizes as follows:

NaOH <=====> Na+ + OH

Thus, the addition of NaOH to water increases the OH or alkali concentration of the resulting solutions.

Another interesting aspect of water is that the concentration of H3O+ and OH remain in balance with each other. An increase in the concentration of H3O+ causes a proportional decrease in the concentration of OH.

Accordingly, a table can be constructed which shows the relationship of the pH's   H3O+ concentration, and OH concentration.

Ion Activity ( Moles / liter )

         pH     H3O+ (Acid)               OH- (Base)
 
         0      1.0                       0.00000000000001
         1      0.1                       0.0000000000001
 
         |      |                         |
         |      |                         |
         |      |                         |
 
        13      0.0000000000001           0.1
        14      0.00000000000001          1.0

Note five things about this chart.

1. As the acid ( H3O+) concentration decreases, the pH increases.

2. As the acid ( H3O+) concentration decreases, the base ( OH ) concentration increases proportionately.

3. At pH 7 the acid ( H3O+) and base ( OH ) concentrations are equal. This is called the neutral point.

4. The pH scale represents the number of places the decimal point is moved to the left of one in expressing the acid ( H3O+) concentration.

5. Each pH unit represents a tenfold change in H3O+ or OH concentration. For example, solution at pH 6 is 10 times more concentrated in H3O+ ions than a solution at pH 7.

Thus, you can see from this chart that the pH scale is a far more concise way of quantitively expressing the acidity of a solution.

How pH is Measured

Today, the pH of a solution is measured either by an indicator dye or by a pH meter and an electrode system whose voltage output is proportional to the active acid ( H3O+) concentration in solution.

Certain organic dye solutions change color over a relatively small pH range. These are called indicator solutions. They can be used to indicate the approximate pH of a solution. By adding a few drops of a phenolphthalein indicator to a solution one can tell if the pH of the solution has a pH greater than 9 by the red color present, or a pH less than 9 by the lack of color. Other dye materials can be chosen whose color changes indicate other pH ranges. For example, phenol red changes at pH 8, bromthymol blue at pH 7, and bromphenol blue at pH 4.

For convenience, these dyes are often deposited on a strip of paper. When a drop of solution to be tested is placed on the paper, the resulting color change is indicative of the approximate pH of the test solution. Dye indicator solutions or paper have the advantage of being quite inexpensive, very portable, and often suitable where only an approximate pH measurement is needed. On the other hand, where precise measurements are needed and / or the solution to be measured is colored, a pH meter is required. Accordingly, pH meter and electrode systems have been developed which respond in a precise manner to the pH of a solution.

To measure pH one can use any number of readily available ($25 – $100 ) pH probes. A pH probe acts like a battery that proportionately generates positive DC voltage for low pH, nothing for pH 7, and negative voltages for high pH values. So, all we have to do is measure this voltage and convert it to pH units.

But there are two problems involved. One problem is that pH is temperature sensitive, with the output voltage ranging from 54 millivolts per pH unit at zero degrees centigrade up to 74 millivolts per pH unit at 100° C.   This means that we have to manually vary the gain or conversion constant of our pH measurement to be able to correct for temperature of the solution being measured.

The second problem is a bit more complex and explains the previously high cost of pH instruments. The source impedance of our pH probe is 15 megohms for the "low–impedance" probes and ranges upwards into hundreds of megohms for special units. In order to measure pH, our voltage amplifier must have an input impedance that is very high compared with 15 megohms. Here is where CMOS electronics has come to the rescue, producing accurate and inexpensive pH meters.

The pH Electrode System

pH electrode systems are always composed of two electrodes, a sensing electrode and a reference electrode. For convenience, these two electrodes can be constructed in one common body which is called a combination electrode. This is the most popular form of the pH electrode system. The sensing electrode contains the specially designed surface whose voltage changes with the pH of the test solution. The reference electrode is used to complete the electrical measuring circuit. Its only function is to give a stable (unchanging) voltage to which the sensing electrode voltage can be compared.

The pH Sensing Electrode

In 1901 a German chemist named Fritz Haber discovered that the voltage at certain glass surfaces changed in a regular manner with the acidity of a solution. Modern pH sensing electrodes are a refinement of this fundamental discovery.

The essential features of a pH sensing electrode are shown in this figure.


Electrode Lead
Electrode Cap
Electrode-Body
Internal Reference
E¹ Electrode

Internal Solution with Constant pH and Reference Ion Activity
pH Sensitive Glass Membrane

TEST SOLUTION


The important requirements of this electrode are that ...

1.) the voltage at the internal reference / filling solution surface ( E ) remain constant,

2.) the voltage at the internal solution / glass membrane surface ( E² ) remain constant, and ...

3.) the voltage at the glass membrane / test solution surface ( E³ )changes proportional to the pH of the test solution.

It should be noted that the electrical resistance of the glass membrane is extremely high. Thus, a specialized voltmeter is required to measure the, voltage from a pH sensing electrode.

The Reference Electrode

When using a voltmeter to measure the voltage at the pH sensing electrode, the electrical circuit must be completed. The reference electrode performs this function. Just a piece of bare wire could be used to complete the circuit. However, the voltage at its surface would change in an unpredictable fashion with time and test sample composition. Accordingly, a reference electrode is a wire which has been terminated with the proper choice of metal and surrounded by the proper metal ion solution, so as to give a constant voltage independent of time and test sample composition.
The essential features of a reference electrode are shown in this figure.

Electrode Lead
Electrode Cap
Reference Metal Wire

Reference Metal Ion Solution

Salt Bridge Solution
Liquid Junction

TEST SOLUTION


The important requirements of this electrode are that the voltages E¹, E², and E³, remain constant with time and test sample composition.

The Combination Electrode

The combination electrode is a version of the pH electrode system in which the pH sensing electrode and the reference electrode are combined into one common body. All comments applicable to the individual electrodes are also applicable to their combination.

The advantages of this form of the electrode system include handling convenience and rugged construction. The single body construction also allows one, to measure the pH of small sample volumes, as well as the pH of surfaces, such as soil and skin.

The pH Meter

A pH meter is a specialized voltmeter which has two fundamental requirements. First, it must be able to function accurately when measuring the voltage of extremely high resistance electrodes. Second, one must be able to change its sensitivity as a voltmeter to correspond to the pH / voltage characteristics of the electrode system.

Most modern pH meters use all solid-state electronics with very high input resistance or impedance characteristics. These meters measure the voltage of the pH electrode system while drawing extremely little current. Fortunately, the voltage change of a pH electrode varies linearly with pH units. At room temperature, a change of 1 pH unit causes a voltage change of about 60 millivolts (mV) or 0.060 volts. At O° centigrade (temperature at which water freezes) 1 pH unit change causes a 54 mV change. At 100° C. a 1 pH unit change causes a 70 mV change. Thus, a properly designed pH meter will have a temperature dial which varies the sensitivity of the meter to match the voltage from the electrodes.

Occasionally, specialized sensing electrodes fall short of delivering the full voltage which theory would predict. Accordingly, very versatile pH meters will also have an additional sensitivity control, called a slope control. This control, like the temperature dial, allows the analyst to vary the sensitivity of the meter to match the voltage from the electrodes.

The pH Standard

The voltage from the pH electrodes at any given pH value can be predicted approximately. However, for highest accuracy, the pH electrode system can be dipped into a solution of known pH and then the meter adjusted to correspond to this pH value. This adjustment is called standardizing the pH system. The solution used is called a pH standard buffer solution. The chemical composition of pH standard buffer solutions have been defined by the U.S. National Bureau of Standards. Such solutions may be prepared by a competent chemist or technician. They are also available from most pH meter manufacturers.
The following table lists the more popular pH standard solutions.

pH Value 25° C. Composition

  1.68 – Potassium Tetroxalate ( 0.05M )
  3.56 – Potassium Hydrogen Tartrate ( Saturated )
  4.01 – Potassium Hydrogen Phthaiate ( 0.05M )
  6.86 – Potassium Dihydrogen Phosphate ( 0.025M )
  9.18 – Borax ( 0.01M )
12.45 – Calcium Hydroxide ( Saturated )


For best accuracy, a pH meter should be standardized using a standard solution whose value is near that of the test solution. However, standardizing with the pH = 6.86 standard constitutes a good compromise when the test solutions cover a broad range of pH values.

A pH meter may be standardized as follows:

1. Rinse electrodes with distilled (or deionized) water and blot dry.
2. Place electrodes in pH standard buffer solution.
3. Adjust pH meter temperature dial to the temperature of the standard solution.
4. Turn pH meter to "operate."
5. Adjust pH meter to the pH value of the standard, using the "standardize" control. THE pH METER IS NOW STANDARDIZED.
6. Turn pH meter to "standby".
7. Remove electrodes from standard and rinse with water.


The pH Measurement Procedure

Once the pH meter is standardized, the measurement procedure is as simple as this:

1. Rinse electrodes with distilled (or deionized) water and blot dry.
2. Place electrodes in test solution.
3. Adjust pH meter temperature dial to the temperature of the test solution.
4. Turn pH meter to "operate."
5. Read the pH of the test solution on the pH meter directly.
6. Turn pH meter to "standby."
7. Remove electrodes from test solution and rinse with water.

Redox Measurements

The measurement of the oxidation-reduction potential of a solution is commonly called a redox measurement. This measurement gives an indication of oxidizing or reducing power of a solution. Since a pH meter is also a very good voltmeter, it can be used in making redox measurements. The sensing electrode used in this measurement is usually platinum, although gold and silver have been used for special purposes. The reference electrode is the same as that used in pH measurements. The electrode potential is usually expressed in millivolts ( mV ). Thus, most pH meters have a ( mV ) scale, as well as a pH scale. Also, since the temperature coefficient varies with the particular redox couple being measured, the temperature control is deactivated during the ( mV ) measurement.

plon Measurements

In recent years electrodes similar to the pH electrode but specific for other ions have been developed. These include electrodes for ammonia, chloride, cyanide, nitrate and sulfide, to name a few. These electrodes may be used in combination with a reference electrode with any modern pH meter. The meter should be standardized in a solution of known plon of the ion of interest, just as in the pH standardization. The plon of a test solution can then be read directly on the meter as in the pH measurement.
 
 
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203 Oak St. Del Mar, CA 92014


As we see, pH is a way of talking about the electrical state of the chemical solution. Very pure water will not conduct electricity and the measure of the resistivity of water is often used as a indicator of water's purity. The higher the resistance the purer the water. It is the ionic charge of the atoms and minerals dissolved in the water that is responsible for electron flow.

"Both Ph and Specific Conductance can be significantly affected by the presence of minute amounts of "Impurities" such as carbonates and oxides. The solubility of many of these is in the range of only 5 to 25 ppm. However, upon conversion to bicarbonates by atmospheric carbon dioxide, solubility may be greatly increased. Example: the conversion of calcium oxide — to calcium carbonate — to calcium bicarbonate." --- Thomas M. Riddick



pH of The Blood   —   Acid–Base Balance
Michael J. Bookallil – Senior Lecturer in Anaesthetics, Royal Prince Alfred Hospital – The University of Sydney

pH and Body Temperature
M J Bookallil


How Our Bodies Regulate pH


BIOlogical TRANSmutations

by Professor C. Louis Kervan

Member of the New York Academy of Science
Director of Conferences of Paris University
Member of Conseil d'Hygiene de la Seine

Translation and Adaptation by Michel Abehsera
Copyright 1989

The movement of life stems from the constant change of one element into another.


If there is abundant literature showing that the presence of K is dependent on the availability of oxygen, there are also several experiments showing its relation to hydrogen, for according to our reaction, K + H = Ca. In other words, if K is too abundant in the presence of H, it will give Ca.

The presence of H is linked to acidity ( low pH ). An excess of H ions signifies an acidity that might become dangerous for the cell. However, in that case K can join an H nucleus to produce Ca, thereby establishing alkalinity and an optimum Ca/K ratio. The agent of equilibrium is thus K. The effects between K and Ca are opposite in appearance only; they are in fact complementary.

Hoagland writes that there is a clear tendency toward acidification of the of the cellular medium, freezing H+ ions; the addition of K+ ions leads to the alkalinization of the cellular liquids.

Reinberg notices that "the alkalinization of the cellular liquids with K is well known by arboriculturists, who use potassium nitrate to speed up fruit maturation."

It is of interest to point out that the proportions of K and Ca are of the same order in animal life — in the plasma as well as in seawater, where life began.

...

Darrow pointed out that a K increase in the cell decreases the cell's acidity because it causes a decrease in H. Thus the alkalinization takes place when K takes H to give Ca. Ca is taken back by the outside liquid and excreted, producing a negative Ca balance sheet. More Ca is excreted than ingested, but the main source of Ca is Mg.

The internal equilibrium of the animal cell postulates a large K content and a small Ca content. The reactions with H help to reduce acidification, since H is taken away.

It has been found that micro-organisms in the soil excrete H ions which acidify the soil; however, K neutralizes this acidity when it comes in contact with the roots.

If the calcium concentration in the nutritive medium is increased, there is a smaller absorption of K. This can be explained by the fact of reversibility:

...

This specific reaction allows a biological equilibrium to be maintained.

( I left out the tables and some equations for now. I think this is another book I will have to scan. )


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