- Email: [email protected]
pH and ORP
pH AND ORP by Michael Banhidi Consultant, Leonia, N,J.
pH The term "pit" is used to express the degree of acidity or alkalinity of a solution. Although many different ions may be formed in solution those that establish whether the solution is acid or alkaline are the hydrogen (H +) and the hydroxyl (OH) ions, respectively. When the hydrogen ions exceed the hydroxyl ions in number, the solution is acid. When they are equal, the solution is neutral. Alkalinity (or basicity) denotes that the hydroxyl ions exceed the hydrogen ions. Acids and alkalis vary in the degree to which they form ions in solution. Those that almost completely ionize are called strong acids or bases. Hydrochloric acid and sodium hydroxide are examples. Other acids and bases ionize to only a small degree and are called weak acids or bases. Acetic acid and ammonium hydroxide are typical examples. Acids and bases may also be classified according to the number of H + ions they can donate per molecule. If an acid can donate only one H + ion it is known as monofunctional. If the available H + ions are more than one a di- (two) or tri- (three) classification is used. Even pure water dissociates to a minute degree into hydrogen and hydroxyl ions. Pure water at 25°C always contains 0.0000007 or, in simple form, 1 × 10-7 gram-equivalents per liter of hydrogen ions and, since pure water is neutral, an equal concentration of hydroxyl ions. The product of these concentrations is a constant equal to 1.0 × 10-14 at 25°C. Thus, if the hydrogen-ion concentration of a solution were 1.0 X 10-4 gram-equivalents per liter, the hydroxyl-ion concentration would be equal to 1 × 10 io gram-equivalents per liter. Because, in this case, the hydrogen-ion concentration is greater than the l!ydroxyl-ion concentration, the solution is acidic. To express hydrogen-ion concentration more conveniently than by the use of decimals or by negative exponents, the term "pH" was adopted. Expressed mathematically, pH is equal to the negative logarithm to the base of 10 of the hydrogen-ion concentration. This is expressed as follows: hydrogen-ion concentration = 1 × 10-v gram-equiv/dents of hydrogen ion per liter or has a pH value of 7. Since acid solutions contain more than 1 × 10-7 gram-equivalents of hydrogen ion per liter, the pH values of these solutions are less than 7. Conversely, alkaline solutions have pH values greater than 7. It is also helpful to remember that a unit change in pH represents a tenfold change in acidity or alkalinity. Thus, compared with a solution of 5 pH, a solution of 4 pH is ten times as acidic--of 3 pH, a hundred times as acidic--and 2 pH, a thousand times as acidic. The commonly accepted range within which pH values are expressed covers the scale 0 to 14. In plating solutions pH determinations below 1 and above 12 are of little or no value. Outside these limits acidities or alkalinities are more conveniently expressed as percent concentrations. Nonaqueous solvents may provide a pH outside the 0 t o 14 range normally encountered in aqueous solutions. Many digital meters are capable of-19.99 to +19.99 pH. In other meters, such as the analog meters, the scale is limited. In this case the millivolt range of the pH meter provides greater range, such as ± 1,400 mV, than the normal 0 to 14 pH range, which is -+421 mV at 25°C; therefore, the pH value can be calculated by comparing the millivolt values obtained in a sample. Of course there is reduced readability with the larger millivolt range. Relating pH and millivolt values observed on the pH meter, the pH valuecan be calculated using the following formula: pH x = pH~ - (E x - Es)/S
599
Table I. Components of Common Buffers Components
Hydrochloric acid Potassium chloride Potassium hydrogen tartrate Potassium hydrogen phthalate Acetic acid Sodium acetate Borax
Concentration
0.365g/L 6.7 g/L 5.64 g/L 10.2 g/L 6.0 g/L 8.2 g/L 19.06 g/L
pH
1:1mixture
2.08
l:lmixture
3.57 4.00 6.9 9.2
where pH x is the pH of the sample (x); pH~ is the pH of the standard (s); E× is the millivolt value of the sample as read by the pH meter; E s is the millivolt value of the standard as read by the pH meter; and S is the slope, normally 59.16 mV/pH unit at 25°C. The pH can be determined colorimetrically or electrometrically.
Colorimetrie Method For a visual indication of the pH a so-called acid-base indicator or mixture of indicators is employed. Such an indicator is a weak organic acid (or base) having the special feature that it changes color within a definite pH range. The standard colorimetric method is carried out by placing a given quantity of solution to be tested in a glass tube. A set amount of indicator is added and compared with standards, in which a solution with known pH has been prepared with the same amount of indicator. This method is cumbersome and rarely used today. The derivative of the standard colorimetric method is the pH test paper. The indicator is impregnated on an absorbent paper strip. A color chart with the appropriate pH values is printed on the strip or a card. After immersion in the solution the strip is compared with the standard color. Reliable pH papers are accurate _+0.3 units from the electrometric value. In the case of a strong oxidizing or reducing solution the pH papers are not reliable.
Electrometrie Method Certain electrodes, such as hydrogen, quinhydron, antimony and others, when immersed in a solution, develop voltages (called electrical potential), which depend upon the pH of the solution. As a result a pH meter is basically a millivolt meter. The glass bulb electrode is the key to making a pH measurement. The special composition glass used is very selective and sensitive to hydrogen ions. The potential that is developed at the glass membrane can be related to the pH of the solution. To complete the circuit and provide a stable and reproducible referencing potential, another electrode is required. The "reference" electrode makes contact with the solution through a junction, which allows slow leakage of the filling solution into the sample. When the glass (pH) and reference electrodes are paired, or built together as a combination electrode, and connected to a pH meter, the voltage developed at the electrode pair is amplified and displayed on a meter or a digital readout. The meter is standardized with a known solution called a buffer. A buffer is a specially prepared solution that resists changes in pH and has a specific pH value at a specific temperature. In Table I the component of the most common buffers is shown. When highly accurate pH measurements are required, temperature, meter standardization, and electrode technique should be considered. Temperature has two effects on the accuracy of pH readings. The first effect is on the electrodes, which is corrected by either automatic temperature compensation (ATC) or by manually adjusting the temperature control on the pH meter. Temperature also effects the pH value of the buffer (which is used for standardization). Because samples do not change pH with temperature in the same way, the 600
Table IL Standard Potential (E °) Electrode Reaction
E ° (V)
Na + + e- = Na CNO- + It20 + 2e- = CN + 2OHZn2+ + 2e- = Zn 2H + + 2e- = H2 Ag/AgC1 electrode, 4N KC1 Calomel electrode, sat KC1 (Cr207)~ + 14H+ + 6e- = 2Cr3+ + 7H20
-2.714 -0.97 -0.763 0.000 +0.•99 +0.244 +1.33
relationship, between temperature and pH must be determined experimentally. Most buffer bottles have a chart on their labels that gives the pH as a function of temperature. The pH meter should be standardized with accurate buffers, which have pH values close to the pH values of the samples. Use of the slope control corrects for nonideal electrode behavior associated with aging of the electrodes. Proper storage of the electrodes increases their life and provides faster response and less drift when making pH measurements. The electrodes should be carefully rinsed between each measurement to prevent contamination. Do not wipe the glass electrode, which could transfer a charge to the glass bulb resulting in a long stabilization time. The rinsed glass electrode should be stored in distilled water. Finally, the consistency of the measuring technique is important, whether the samples are measured in static or stirred condition; however, the standardization buffer and the sample solution should be treated the same.
OXIDATION-REDUCTION POTENTIAL Oxidation-reduction potential (ORP) measurements are used to monitor chemical reactions, quantify ion activity, or determine the oxidizing or reducing properties of solutions. Although ORP measurements are somewhat similar to those of pH, the potential value must be interpreted carefully for useful results. An ORP measurement is made using the millivolt mode of the pH meter. Consequently, by substituting a metallic electrode for the pH glass electrode, many other ions besides the hydrogen ion can be detected with the same pH meter. In many chemical reactions electrons are transferred from one substance to another. By definition a substance gains electrons in a reduction reaction and loses electrons in an oxidation reaction. Oxidation and reduction reactions occur together. The available electrons from an oxidized substance are taken up by the reduced substance until an equilibrium condition is reached. Since it is impossible to measure absolute potentials, an arbitrary standard, the hydrogen electrode, is chosen. ORPs are defined relative to this standard. The electrode reaction 2H + + 2e- = H 2 is assigned a potential of 0.000 V when the hydrogen activity is 1 M (concentration of 1 mole/L), and the partial pressure of hydrogen is 1 arm. When reactions are written as oxidation (e.g., Na = Na + + e-), potentials have the opposite polarity. The standard potential E o of any oxidation-reduction reaction is referenced to the standard hydrogen electrode and refers to the condition of the oxidation-reduction reaction where temperature is 25°C, ion activity is unity, and gases are at 1 arm pressure. Table II shows the standard potential E ° associated with various reactions. The ORP is characteristic of reactions involving both oxidation and reduction and varies as a function of (1) the standard potential, (2) relative ion concentration, (3) temperature, and (4) the number of electrons transferred in the reactions.
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ORPs are usually displayed as millivolts (mV). When measured with a pH meter (set to read in mV), this ORP is generally the electromotive force (EMF) difference developed between the ORP electrode and a constant voltage reference electrode (saturated calomel, instead of a normal hydrogen electrode) that is immersed in the solution. Any one of three different types of metallic electrodes may be used. There are three types of metallic electrodes used in ORP measurements that differ in constmction but are based on the same principle that an oxidized and a reduced state must be present. The first type of metallic electrode to be considered consists of a metal contact with a solution of its own ion. The metallic electrode is in a reduced state and its ions are in an oxidized state. An example of this type is silver in a silver nitrate solution. It is used mainly on the analytical field. The second type of metallic electrode consists of a metal coated with a sparingly soluble salt of this metal in a solution of soluble salt with the same anion (e.g., silver-silver chloride in a solution of potassium chloride). The third type of metallic electrode consists of an inert metal in contact with a solution containing both the reduced and oxidized state of an oxidation-reduction system. An example would be platinum in contact with ferric-ferrous ions. Platinum and gold are the most common ORP electrodes. The nature of the test solution and the method to be used will determine the choice of the electrode. The reference electrodes can be identical, but a noble metal electrode replaces the glass pH electrode. The signal from the ORP electrodes must be fed into an amplifier with high-input resistance. If one or both reactions pair hydrogen ions the ORP measurement becomes pH dependentl Consider the following reaction, which occurs in the reduction of hexavalent chromium: Cr2072- + 14H + + 6e
= 2Cr 3+ + 7H20
The reaction depends on solution pH. Potential changes measured by the ORP electrode will continue to vary with the redox ratio, but the absolute potential will also vary with pH. In the first step the pH is lowered to 2 to 2.5. Sulfur dioxide or sodium sulfite solution is used as the reductant. The overall reaction is Cr2072- + 2H + + 3H2SO 3 = 2Cr 3+ + 4H20 + 3S042 In the second part of the process the waste liquors are neutralized to a pH level of 7 to 8. At this pH the chromic ion precipitates as a sludge and is sent to clarifiers for ultimate disposal or recovery. All these reactions will take place in a definite pH and in a specific millivolt range. Although applications for ORP measurement are not as widespread as for pH one of the most important applications in the metal-finishing industry is the oxidation of cyanide wastes. Oxidation converts the toxic cyanides to harmless compounds. Typically, chlorine gas or sodium hypoehlorite are the oxidants. In the first step caustic or lime is added to make the cyanide-bearing waste alkaline to a pH of 9 to 10. An acid solution would release deadly cyanide gas; therefore, the system generally incorporates pH control. The first stage reactions are C12 + C N - + 2OH7 = 2C1- + OCN
+ H20
The second step in cyanide oxidation takes place at a controlled pH in the 7 to 8 range. The reaction is: 3C12 + 2OCN
+ 4OH
= 6C1- + 2CO 2 + N z + 2H20
In industrial plants ORP is rarely applied to nice clean reactions where potentials can be estimated easily. In waste treatment or sewage plants, for instance, solutions contain a host of constituents that the reagent oxidizes and reduces simultaneously. ORP relates to the concentrations and activities-of all participating reactions. It frequently becomes necessary to determine the control points experimentally.
603
Table III. O R P Values of Quinhydrone-Buffer Standards pH4
Reference Ag/AgC1 Calomel
pH7
20 °C
25 °C
30 °C
20 °C
25 °C
30 °C
+268 +233
+263 +218
+258 +213
+92 +47
+86 +41
+79 +34
If day-to-day relative potential values are to be compared, the pH meter must be standardized to the same starting point. Short the meter glass and reference inputs and adjust the standardization control until zero millivolt is displayed using the "absolute rnillivolt" mode to set the potential to some arbitrary value when the electrodes are reading the potential in a repeatable standard solution. Because ORP is a characteristic measure of redox equilibrium it should not require standardization or calibration. The measured potential is absolute in a sense. Yet, frequently, it is desirable to check systems for proper operation and electrode poisoning. Solutions of known potential can be developed b y saturating buffer solutions with quinhydrone. The reaction is such that the measured potential will vary only along with the solution pH and temperature. The procedure is as follows: (1) Saturate the buffer with quinhydrone, made up fresh for each test. Quinhydroue is not readily soluble so a few crystals stirred into the buffer are sufficient. The solution will he amber colored. (2) Clean the platinum electrode. (3) Place the platinum and reference electrodes in a quinhydrone-buffer solution and measure the potential and temperature. Measured potential will generally be within ± 10 mV of theoretical value. The ORP values of quinhydrone-huffer solution can be seen in Table III.
j/, there is a solution for pH monitoring in the ;e of acidic fluorides, Sensorex's N E W HF version 'face self cleaning pH electrodes resist coatings, n, breakage and now F'(fluoride) too! Available in four convenient styles: • In-Line • Side S t r e a m ace.
• Submersion • I n s e r t i o n / W e t Tap
ther details, fax, phone, o r email Sensorex and ~, our F o r m HFFS and Bulletin S C 6 0 0 . 11661 Seaboard Circle, Stanton, CA 90680 • USA 714,895,4344 • Fax714,894.4839 [email protected] • ht~p://www.sensorex.com
604
pH The term "pit" is used to express the degree of acidity or alkalinity of a solution. Although many different ions may be formed in solution those that establish whether the solution is acid or alkaline are the hydrogen (H +) and the hydroxyl (OH) ions, respectively. When the hydrogen ions exceed the hydroxyl ions in number, the solution is acid. When they are equal, the solution is neutral. Alkalinity (or basicity) denotes that the hydroxyl ions exceed the hydrogen ions. Acids and alkalis vary in the degree to which they form ions in solution. Those that almost completely ionize are called strong acids or bases. Hydrochloric acid and sodium hydroxide are examples. Other acids and bases ionize to only a small degree and are called weak acids or bases. Acetic acid and ammonium hydroxide are typical examples. Acids and bases may also be classified according to the number of H + ions they can donate per molecule. If an acid can donate only one H + ion it is known as monofunctional. If the available H + ions are more than one a di- (two) or tri- (three) classification is used. Even pure water dissociates to a minute degree into hydrogen and hydroxyl ions. Pure water at 25°C always contains 0.0000007 or, in simple form, 1 × 10-7 gram-equivalents per liter of hydrogen ions and, since pure water is neutral, an equal concentration of hydroxyl ions. The product of these concentrations is a constant equal to 1.0 × 10-14 at 25°C. Thus, if the hydrogen-ion concentration of a solution were 1.0 X 10-4 gram-equivalents per liter, the hydroxyl-ion concentration would be equal to 1 × 10 io gram-equivalents per liter. Because, in this case, the hydrogen-ion concentration is greater than the l!ydroxyl-ion concentration, the solution is acidic. To express hydrogen-ion concentration more conveniently than by the use of decimals or by negative exponents, the term "pH" was adopted. Expressed mathematically, pH is equal to the negative logarithm to the base of 10 of the hydrogen-ion concentration. This is expressed as follows: hydrogen-ion concentration = 1 × 10-v gram-equiv/dents of hydrogen ion per liter or has a pH value of 7. Since acid solutions contain more than 1 × 10-7 gram-equivalents of hydrogen ion per liter, the pH values of these solutions are less than 7. Conversely, alkaline solutions have pH values greater than 7. It is also helpful to remember that a unit change in pH represents a tenfold change in acidity or alkalinity. Thus, compared with a solution of 5 pH, a solution of 4 pH is ten times as acidic--of 3 pH, a hundred times as acidic--and 2 pH, a thousand times as acidic. The commonly accepted range within which pH values are expressed covers the scale 0 to 14. In plating solutions pH determinations below 1 and above 12 are of little or no value. Outside these limits acidities or alkalinities are more conveniently expressed as percent concentrations. Nonaqueous solvents may provide a pH outside the 0 t o 14 range normally encountered in aqueous solutions. Many digital meters are capable of-19.99 to +19.99 pH. In other meters, such as the analog meters, the scale is limited. In this case the millivolt range of the pH meter provides greater range, such as ± 1,400 mV, than the normal 0 to 14 pH range, which is -+421 mV at 25°C; therefore, the pH value can be calculated by comparing the millivolt values obtained in a sample. Of course there is reduced readability with the larger millivolt range. Relating pH and millivolt values observed on the pH meter, the pH valuecan be calculated using the following formula: pH x = pH~ - (E x - Es)/S
599
Table I. Components of Common Buffers Components
Hydrochloric acid Potassium chloride Potassium hydrogen tartrate Potassium hydrogen phthalate Acetic acid Sodium acetate Borax
Concentration
0.365g/L 6.7 g/L 5.64 g/L 10.2 g/L 6.0 g/L 8.2 g/L 19.06 g/L
pH
1:1mixture
2.08
l:lmixture
3.57 4.00 6.9 9.2
where pH x is the pH of the sample (x); pH~ is the pH of the standard (s); E× is the millivolt value of the sample as read by the pH meter; E s is the millivolt value of the standard as read by the pH meter; and S is the slope, normally 59.16 mV/pH unit at 25°C. The pH can be determined colorimetrically or electrometrically.
Colorimetrie Method For a visual indication of the pH a so-called acid-base indicator or mixture of indicators is employed. Such an indicator is a weak organic acid (or base) having the special feature that it changes color within a definite pH range. The standard colorimetric method is carried out by placing a given quantity of solution to be tested in a glass tube. A set amount of indicator is added and compared with standards, in which a solution with known pH has been prepared with the same amount of indicator. This method is cumbersome and rarely used today. The derivative of the standard colorimetric method is the pH test paper. The indicator is impregnated on an absorbent paper strip. A color chart with the appropriate pH values is printed on the strip or a card. After immersion in the solution the strip is compared with the standard color. Reliable pH papers are accurate _+0.3 units from the electrometric value. In the case of a strong oxidizing or reducing solution the pH papers are not reliable.
Electrometrie Method Certain electrodes, such as hydrogen, quinhydron, antimony and others, when immersed in a solution, develop voltages (called electrical potential), which depend upon the pH of the solution. As a result a pH meter is basically a millivolt meter. The glass bulb electrode is the key to making a pH measurement. The special composition glass used is very selective and sensitive to hydrogen ions. The potential that is developed at the glass membrane can be related to the pH of the solution. To complete the circuit and provide a stable and reproducible referencing potential, another electrode is required. The "reference" electrode makes contact with the solution through a junction, which allows slow leakage of the filling solution into the sample. When the glass (pH) and reference electrodes are paired, or built together as a combination electrode, and connected to a pH meter, the voltage developed at the electrode pair is amplified and displayed on a meter or a digital readout. The meter is standardized with a known solution called a buffer. A buffer is a specially prepared solution that resists changes in pH and has a specific pH value at a specific temperature. In Table I the component of the most common buffers is shown. When highly accurate pH measurements are required, temperature, meter standardization, and electrode technique should be considered. Temperature has two effects on the accuracy of pH readings. The first effect is on the electrodes, which is corrected by either automatic temperature compensation (ATC) or by manually adjusting the temperature control on the pH meter. Temperature also effects the pH value of the buffer (which is used for standardization). Because samples do not change pH with temperature in the same way, the 600
Table IL Standard Potential (E °) Electrode Reaction
E ° (V)
Na + + e- = Na CNO- + It20 + 2e- = CN + 2OHZn2+ + 2e- = Zn 2H + + 2e- = H2 Ag/AgC1 electrode, 4N KC1 Calomel electrode, sat KC1 (Cr207)~ + 14H+ + 6e- = 2Cr3+ + 7H20
-2.714 -0.97 -0.763 0.000 +0.•99 +0.244 +1.33
relationship, between temperature and pH must be determined experimentally. Most buffer bottles have a chart on their labels that gives the pH as a function of temperature. The pH meter should be standardized with accurate buffers, which have pH values close to the pH values of the samples. Use of the slope control corrects for nonideal electrode behavior associated with aging of the electrodes. Proper storage of the electrodes increases their life and provides faster response and less drift when making pH measurements. The electrodes should be carefully rinsed between each measurement to prevent contamination. Do not wipe the glass electrode, which could transfer a charge to the glass bulb resulting in a long stabilization time. The rinsed glass electrode should be stored in distilled water. Finally, the consistency of the measuring technique is important, whether the samples are measured in static or stirred condition; however, the standardization buffer and the sample solution should be treated the same.
OXIDATION-REDUCTION POTENTIAL Oxidation-reduction potential (ORP) measurements are used to monitor chemical reactions, quantify ion activity, or determine the oxidizing or reducing properties of solutions. Although ORP measurements are somewhat similar to those of pH, the potential value must be interpreted carefully for useful results. An ORP measurement is made using the millivolt mode of the pH meter. Consequently, by substituting a metallic electrode for the pH glass electrode, many other ions besides the hydrogen ion can be detected with the same pH meter. In many chemical reactions electrons are transferred from one substance to another. By definition a substance gains electrons in a reduction reaction and loses electrons in an oxidation reaction. Oxidation and reduction reactions occur together. The available electrons from an oxidized substance are taken up by the reduced substance until an equilibrium condition is reached. Since it is impossible to measure absolute potentials, an arbitrary standard, the hydrogen electrode, is chosen. ORPs are defined relative to this standard. The electrode reaction 2H + + 2e- = H 2 is assigned a potential of 0.000 V when the hydrogen activity is 1 M (concentration of 1 mole/L), and the partial pressure of hydrogen is 1 arm. When reactions are written as oxidation (e.g., Na = Na + + e-), potentials have the opposite polarity. The standard potential E o of any oxidation-reduction reaction is referenced to the standard hydrogen electrode and refers to the condition of the oxidation-reduction reaction where temperature is 25°C, ion activity is unity, and gases are at 1 arm pressure. Table II shows the standard potential E ° associated with various reactions. The ORP is characteristic of reactions involving both oxidation and reduction and varies as a function of (1) the standard potential, (2) relative ion concentration, (3) temperature, and (4) the number of electrons transferred in the reactions.
602
ORPs are usually displayed as millivolts (mV). When measured with a pH meter (set to read in mV), this ORP is generally the electromotive force (EMF) difference developed between the ORP electrode and a constant voltage reference electrode (saturated calomel, instead of a normal hydrogen electrode) that is immersed in the solution. Any one of three different types of metallic electrodes may be used. There are three types of metallic electrodes used in ORP measurements that differ in constmction but are based on the same principle that an oxidized and a reduced state must be present. The first type of metallic electrode to be considered consists of a metal contact with a solution of its own ion. The metallic electrode is in a reduced state and its ions are in an oxidized state. An example of this type is silver in a silver nitrate solution. It is used mainly on the analytical field. The second type of metallic electrode consists of a metal coated with a sparingly soluble salt of this metal in a solution of soluble salt with the same anion (e.g., silver-silver chloride in a solution of potassium chloride). The third type of metallic electrode consists of an inert metal in contact with a solution containing both the reduced and oxidized state of an oxidation-reduction system. An example would be platinum in contact with ferric-ferrous ions. Platinum and gold are the most common ORP electrodes. The nature of the test solution and the method to be used will determine the choice of the electrode. The reference electrodes can be identical, but a noble metal electrode replaces the glass pH electrode. The signal from the ORP electrodes must be fed into an amplifier with high-input resistance. If one or both reactions pair hydrogen ions the ORP measurement becomes pH dependentl Consider the following reaction, which occurs in the reduction of hexavalent chromium: Cr2072- + 14H + + 6e
= 2Cr 3+ + 7H20
The reaction depends on solution pH. Potential changes measured by the ORP electrode will continue to vary with the redox ratio, but the absolute potential will also vary with pH. In the first step the pH is lowered to 2 to 2.5. Sulfur dioxide or sodium sulfite solution is used as the reductant. The overall reaction is Cr2072- + 2H + + 3H2SO 3 = 2Cr 3+ + 4H20 + 3S042 In the second part of the process the waste liquors are neutralized to a pH level of 7 to 8. At this pH the chromic ion precipitates as a sludge and is sent to clarifiers for ultimate disposal or recovery. All these reactions will take place in a definite pH and in a specific millivolt range. Although applications for ORP measurement are not as widespread as for pH one of the most important applications in the metal-finishing industry is the oxidation of cyanide wastes. Oxidation converts the toxic cyanides to harmless compounds. Typically, chlorine gas or sodium hypoehlorite are the oxidants. In the first step caustic or lime is added to make the cyanide-bearing waste alkaline to a pH of 9 to 10. An acid solution would release deadly cyanide gas; therefore, the system generally incorporates pH control. The first stage reactions are C12 + C N - + 2OH7 = 2C1- + OCN
+ H20
The second step in cyanide oxidation takes place at a controlled pH in the 7 to 8 range. The reaction is: 3C12 + 2OCN
+ 4OH
= 6C1- + 2CO 2 + N z + 2H20
In industrial plants ORP is rarely applied to nice clean reactions where potentials can be estimated easily. In waste treatment or sewage plants, for instance, solutions contain a host of constituents that the reagent oxidizes and reduces simultaneously. ORP relates to the concentrations and activities-of all participating reactions. It frequently becomes necessary to determine the control points experimentally.
603
Table III. O R P Values of Quinhydrone-Buffer Standards pH4
Reference Ag/AgC1 Calomel
pH7
20 °C
25 °C
30 °C
20 °C
25 °C
30 °C
+268 +233
+263 +218
+258 +213
+92 +47
+86 +41
+79 +34
If day-to-day relative potential values are to be compared, the pH meter must be standardized to the same starting point. Short the meter glass and reference inputs and adjust the standardization control until zero millivolt is displayed using the "absolute rnillivolt" mode to set the potential to some arbitrary value when the electrodes are reading the potential in a repeatable standard solution. Because ORP is a characteristic measure of redox equilibrium it should not require standardization or calibration. The measured potential is absolute in a sense. Yet, frequently, it is desirable to check systems for proper operation and electrode poisoning. Solutions of known potential can be developed b y saturating buffer solutions with quinhydrone. The reaction is such that the measured potential will vary only along with the solution pH and temperature. The procedure is as follows: (1) Saturate the buffer with quinhydrone, made up fresh for each test. Quinhydroue is not readily soluble so a few crystals stirred into the buffer are sufficient. The solution will he amber colored. (2) Clean the platinum electrode. (3) Place the platinum and reference electrodes in a quinhydrone-buffer solution and measure the potential and temperature. Measured potential will generally be within ± 10 mV of theoretical value. The ORP values of quinhydrone-huffer solution can be seen in Table III.
j/, there is a solution for pH monitoring in the ;e of acidic fluorides, Sensorex's N E W HF version 'face self cleaning pH electrodes resist coatings, n, breakage and now F'(fluoride) too! Available in four convenient styles: • In-Line • Side S t r e a m ace.
• Submersion • I n s e r t i o n / W e t Tap
ther details, fax, phone, o r email Sensorex and ~, our F o r m HFFS and Bulletin S C 6 0 0 . 11661 Seaboard Circle, Stanton, CA 90680 • USA 714,895,4344 • Fax714,894.4839 [email protected] • ht~p://www.sensorex.com
604