X. USING A pH METER AND PREPARING BUFFER SOLUTIONS

BACKGROUND

Acids and bases For purposes of getting started, an acid can be defined as a chemical compound that yields hydrogen ions (H+) when it is dissolved in water, and a base can be defined as a chemical compound that yields hydroxyl ions (OH-) when it is dissolved in water.  Examples of commonly encountered acids are hydrochloric acid (HCl), sulfuric acid (H2SO4), and acetic acid (C2H4O2), and some common bases are sodium hydroxide (NaOH) and potassium hydroxide (KOH).  Acids and bases react with each other, and when they do so the resulting products are water (H2O) and an ionic compound, i.e., a "salt".  This type of reaction is described overall as "neutralization".

EXAMPLE:  sodium hydroxide plus hydrochloric acid --> sodium chloride plus water.

As a chemical equation: NaOH + HCl --> NaCl + H2O.

Expressed as amounts and concentrations, the reaction means that one mole of NaOH will react with one mole of HCl.  Therefore, 1 liter of 1M NaOH (which contains one mole of OH-) will react with and neutralize 1 liter of 1M HCl (which contains one mole of H+).  Likewise, 2 liters of 0.5M NaOH (which also contains one mole of OH-) will react with and neutralize 1/2 liter (500ml) of 2M HCl (which contains one mole of H+).

Gram equivalent weights and normality

The gram equivalent weight (sometimes called simply the "equivalent weight") of an acid is the weight in grams that will provide one mole (6.02 X 1023) of H+ ions.  The gram equivalent weight of a base is the weight in grams that will provide one mole (6.02 X 1023) of OH- ions.

EXAMPLE:  potassium hydroxide (KOH) has a formula weight = 56.1. Since 1 mole of KOH represents 1 mole of OH-, the gram equivalent weight of KOH is 56.1.

EXAMPLE:  hydrochloric acid (HCl) has a formula weight = 36.45. 1 mole of HCl represents 1 mole of H+, so the gram equivalent weight of HCl is 36.45.

Some acids and bases can yield more than one H+ or OH-.

EXAMPLE:  sodium phosphate (NaH2PO4) has a formula weight = 120. Since 1 mole of NaH2PO4 represents 2 moles of H+, the gram equivalent weight of NaH2PO4 is 120/2 = 60.

EXAMPLE:  magnesium hydroxide (Mg(OH)2) has a formula weight = 58.34. Since 1 mole of Mg(OH)2 represents 2 moles of OH-, the gram equivalent weight of Mg(OH)2 is 58.34/2 = 29.17.

A solution that contains one gram equivalent weight of an acid or a base is a one normal (1 N) solution, and the number of gram equivalent weights per liter is called the normality of the solution.  A specific volume of any 1N acid solution (say, 10ml of 1N) will exactly neutralize the same volume of any 1N base solution (again, 10ml of 1N).

In practice, for any acid that will yield one H+ in water, according to its chemical formula (such as HCl), a 1N solution = 1M = 1 gram formula weight per liter.  Correspondingly, for any base yielding one OH- according to its chemical formula (such as NaOH), a 1N solution = 1M = 1 gram formula weight per liter.  For any acid that will yield two H+ (such as sulfuric acid, H2SO4), a 2N solution = 1M = 1 gram formula weight per liter.  For any base yielding two OH- (such as Mg(OH)2, a 2N solution = 1M = 1 gram formula weight per liter.

NOTE:  Not all hydrogen atoms (H) or hydroxyl groups (OH) of a compound will necessarily be yielded as ions; covalently-bonded atoms will not.  Thus, you must know the properties of a given compound in addition to its chemical formula.

EXERCISE #1:

Examine the information sheet entitled "ACIDS AND BASES - NORMALITY (N), posted on the wall above the pH meter.  It contains information from the CRC Handbook of Chemistry and Physics.

What is the starting molarity of 36% (concentrated) hydrochloric acid (HCl)?  What is the normality of 36% HCl?

What molarity is concentrated acetic acid (99-100% CH3COOH = glacial acetic acid)?  What is the normality of glacial acetic acid?

What is the starting molarity of 95% (concentrated) sulfuric acid (H2SO4)?  What is the normality of 95% H2SO4?

What molarity is concentrated phosphoric acid (85% H3PO4)?  What is the normality of 85% H3PO4?

List molarities and normalities of these acids in a Table in your Tech. Facility notebook.

Acids and bases - strong versus weak

Some acids, such as HCl, ionize essentially completely when dissolved in water, and some bases, such as NaOH and KOH do likewise.  These are known as strong acids and bases.  Other acids and bases ionize only partially in water, and these are known as weak acids and bases.  Acetic acid, for example, forms a mixture of hydrogen ions (H+), acetate ions (C2H3O2-), and intact non-ionized acetic acid molecules (C2H4O2) in water.  Nevertheless, if the H+ in the initial mix were to react with something else and be removed from the solution, more H+ would be produced, eventually using up all of the H+ in the "donor" acetic acid molecules.  Another way of saying this is that the reaction would continue to go toward equilibrium.

pH:

Pure water consists of H2O molecules plus very small numbers of H+ and OH- ions produced by ionization of water molecules.  These ions are responsible for the slight conductivity of pure water (see module on water quality).  Pure water is "neutral"; it contains equal concentrations of H+ and OH- ions, each 10-7M.  These concentrations shift when pure water is mixed with other chemicals.  When the concentration of H+ (i.e., [H+]; brackets [] denote concentration) is greater than that of OH- (i.e., [H+] > [OH-]), the solution is "acidic", and when [H+] < [OH-], the solution is "basic".  The product of [H+] X [OH-] is essentially constant, and measurements have shown that it equals 10-14.  Therefore, if [H+] = 10-2, then [OH-] = 10-12 (the exponents added together = -14).  If [H+] = 10-11, then [OH-] = 10-3 (again, the exponents added together = -14).

pH is defined in words as "minus the log to the base 10 of the hydrogen ion concentration", or, in proper notation, -log10 [H+].  The reason that pH is used to designate the H+ level rather than the actual concentration of H+ is that the latter varies over an enormous numerical range.  pH, on the other hand, is a logarithmic scale yielding whole positive numbers to work with.  It is simply a device introduced for convenience (by Sorenson, in 1909), so that we don't have to deal with wide-ranging exponents.

In a neutral solution, or in absolutely pure water, [H+] = 10-7, and [OH-] = 10-7.  Since the log10 [H+] = log10 [10-7] = -7, the pH of a neutral solution = -(-7) = 7.  In an acidic solution, say one with [H+] = 10-2 and [OH-] = 10-12, the pH = -log1010-2 = -(-2) = 2.  In a basic solution, such as one with [H+] = 10-12 and [OH-] = 10-2., the pH = -log1010-12 = -(-12) = 12.  Thus, the more acid a solution is, the lower its pH; conversely, the more basic the solution, the higher its pH.

EXAMPLE:  What is the pH of a 0.01M solution of HCl?  [H+] =

10-2, so that the pH = 2 (very acidic). EXAMPLE:  What is the pH of a 0.000001M solution of HCl?  [H+] = 10-6, so that the pH = 6 (mildly acidic).

EXAMPLE:  What is the pH of a 0.01M solution of NaOH?  Here, [OH-] = 10-2, so [H+] must = 10-12 (remember that [H+] X [OH-] =

10-14), and the pH = 12 (very basic).

MEASURING pH - SOME PRACTICAL INFORMATION: Measurements of pH are made in the laboratory using a pH meter equipped with a "glass" electrode (also called an "indicating" electrode) and a "reference" electrode.  Sometimes the two electrodes are separate, and sometimes they are combined into one unit (called a "combination electrode").  You will be using a combination electrode.  Some pH meters have a digital readout, such as the one you will use, and some have a visible scale of numbers with a moving pointer (analog readout).  You will be using a digital pH meter representative of newer instruments, but an older analog meter is located near the digital meter for demonstration purposes.

Glass electrode:  This is the pH-sensing electrode, with components as shown in Fig. 1, A.  During measurements, an electrical potential is produced across the thin "glass membrane" tip of the glass electrode, which is a bulb of special glass that is electrically conductive.  The rest of the glass body of the electrode is non-conductive.  The potential is proportional to the H+ concentration, and the pH meter converts the potential differences directly into pH units.  THE GLASS TIP OF THE ELECTRODE IS EXTREMELY FRAGILE!  GREAT CARE MUST BE TAKEN NOT TO TOUCH IT TO ANY HARD SURFACE (SUCH AS HITTING THE BOTTOM OR EDGE OF A TUBE OR OF A BEAKER)!!  THESE ELECTRODES ARE VERY EXPENSIVE; ONE BROKEN ELECTRODE = - \$100 OR MORE!!

Reference electrode (Fig. 1, B):  The reference electrode serves as a stable potential standard for comparison with the glass electrode.  It also completes the electrical measuring circuit, making use of a porous "liquid junction" to do so.  It contains an electrolyte solution, typically saturated KCl, and, to complete the circuit, a very small amount of this concentrated KCl solution must pass through a very small hole (at bottom of electrode) continuously while pH measurements are being made.  This KCl actually contaminates the solution being measured, but it is normally negligible unless the volume of solution being measured is itself very small.  Proper functioning of the liquid junction (i.e., flow of KCl) requires that an air hole near the top of the electrode be opened up prior to making measurements.  If contamination with traces of KCl cannot be tolerated for a particular experimental solution, pH measurements can be made on small samples of the solution that are then discarded.  If the pH of a large volume of solution needs to be adjusted, the required volumes of acid or base needed for the adjustment can be determined on small samples (discarded), and extrapolated for addition to the large volume without using the electrode.

IN PRACTICE, IN MOST CASES THE KCl CONTAMINATION IS IGNORED.  BEFORE USING THE pH METER IN A RESEARCH LAB, IT IS A GOOD IDEA TO CHECK WITH YOUR RESEARCH SPONSOR OR OTHER EXPERIENCED WORKERS, JUST TO MAKE SURE THAT IT IS ALRIGHT TO DO SO. In a combination electrode, both a glass electrode and a reference electrode are contained within a single glass body (Fig. 1, C).

Calibration of the pH meter: The pH meter, together with its electrodes, is calibrated by using solutions of known pH.  These pH "standard" solutions are prepared and sold commercially by various chemical companies.  We will have three pH standard solutions available in the Tech Facility: pH 4, pH 7, and pH 10.

EXERCISE #2:  TYPES OF pH METERS

Look for the pH meter station in the Techniques Facility.  Examine the two pH meters visually (visually only - do not twiddle any dials yet!!), comparing their controls and readout.

ANALOG VS. DIGITAL

Note that one of the pH meters has a digital readout, i.e., the pH will be shown directly in numbers.  The other pH meter has an analog readout, in which numerical data are expressed         indirectly in terms of a measured quantity (here, the position of a pointer controlled by voltages).  Either, or possibly even both pH meter types can be found in a given research laboratory.

The analog pH meter is the "Accumet Model 5".  You will use this one first.  However, before operating it, you must locate and familiarize yourself with the functions of the various front panel control knobs and the analog display.  The latter indicates pH over the 0 - 14 range, in 0.2 pH unit increments.

EXERCISE #3:  GETTING FAMILIAR WITH THE "ACCUMET MODEL 5"

pH METER

(1) Find the various front panel controls:

(a) function selector -  allows the pH meter to be placed on 3 different modes:  "Standby" mode when no measurements are being taken; "pH" mode for pH measurements; "mV" mode for millivolt    measurements (which we won't do).

(b) slope control -  used only when in the "pH" mode in order to adjust the electrode response.

(c) temperature control -  measures temperature over the range of 0o-100oC.  It is used only for manual compensation of   temperature when making measurements in the "pH" mode.  For our purposes, during operation of the pH meter, the temperature

control will be set at 25oC, room temperature.

(d) standardize control -  when used for determining pH, this control will calibrate the pH meter at the pH value of the  Standard Buffer Solution used in the Standardization procedure.

(2) Look at the combination pH electrode on this meter:

(a) If necessary, review important features and use of the combination electrode - refer to pages 4-6 of this Module.

(b) Locate a white circular cap ring at the top of the electrode.  In this particular combination electrode model, the  fill hole is located on this cap ring.

pH METER STANDARDIZATION

The first step in doing pH measurements on unknown solutions, or in adjusting the pH of a solution to desired value, is to test whether the pH meter is reading the pH of known solutions correctly.  If it is not, its settings must be adjusted so that the readings are accurate.  This is called pH meter "standardization".

The pH meter is calibrated using Standard pH Buffer Solutions that are usually purchased from chemical supply companies.

EXERCISE #4:  DOING A ONE-POINT STANDARDIZATION ON "THE ACCUMET MODEL 5" pH METER

(1)  Turn on pH meter using the "ON" switch of power strip into which meter is plugged.  Allow ~10 min. for the meter to warm up.

(2)  Check to see that pH meter "function" selector is on "Standby" mode before starting.

(3)  Lift the electrode carefully from its arm, just enough to be able to turn the cap ring hole from the "Close" position to the "Open" position.  The fill hole must be OPEN when pH measurements are being made.

(4)  Check to see that the "temperature" control knob is at 25oC; if not, set it to that temperature.

(5)  Obtain a sample of pH 7.0 Standard Buffer solution from

a plastic beaker labeled "Standard Buffers/Unknown", near the pH meter.  Use these screw-cap tubes "as is",  i.e., do not pour samples into other tubes.  They can be re-used many times.

These pH standards come from the bottles of Standard Buffer Stock Solutions located in the "CHEMICALS" cabinet.  Look for them in the cabinet, and examine their labels.  Note that the standard buffer solutions used in the Tech Facility are color coded with dyes (pink, blue, yellow) that make it easier to tell what their pH is.  In other labs, they may not be color coded.

(6)  Carefully lift the support arm holding the electrode so that the electrode is raised out of the bottle of electrode storage solution.

NOTE:  When handling the electrode, do so very carefully.  Even a slight scratch on the glass bulb tip can render the electrode useless for pH determinations.

(7)  Using the "squirt" bottle labeled "pH meter only", rinse electrode off with deionized water by squirting the upper region and letting the water run down into the "waste beaker" held underneath it.  (Note: The waste beaker is labeled, and is near the pH meter).

(8)  Dry the electrode gently in a downward direction with a Kimwipe.  Do not squeeze the electrode, especially the bulb tip, when wiping.

(9)  Carefully lower the electrode into the pH 7.0 buffer solution in plastic tube.  DO NOT TOUCH THE ELECTRODE TIP TO THE BOTTOM OR EDGE OF THE TUBE.  IT CAN CRACK VERY EASILY!

NOTE:  The friction knob on the armature supporting the electrode has already been adjusted to maintain a set position but enough to allow the arm to easily be raised and lowered without continually loosening or tightening the screw.  If the support arm does not hold its position, then the screw should be tightened accordingly.  Adjustments of the friction knob should be made only when necessary.

(10)  Rotate the "function" selector from "Standby" mode to "pH"

mode.

(11)  Set the "slope" control knob to 100%.  Wait several seconds to allow the pH indicator to stabilize.

(12)  Gradually adjust the "standardize" control knob until the analog display indicates pH 7.0 for the Standard Buffer.

NOTE:  The most accurate way to read the pH is to line up the pH indicator (pointer) with its reflected image on the mirrored strip behind it, doing so with one eye closed.

The pH meter is now standardized at pH 7.0.

(13)  Once Standardization is completed, switch the "function"

selector back to "Standby" mode.

(14)  Raise electrode, then remove and re-cap the pH 7.0 Standard Buffer and return it to its original place.

(15)  Rinse off the electrode with deionized water (squirt bottle) and dry with Kimwipe as before (Steps 7, 8).

(This is done in order to avoid contamination of one solution with another.  Thus, it is very important that proper rinsing of the electrode always be done between sample measurements.)

(16)  This completes the steps for performing a 1-point pH Standardization.  The pH meter is now ready for making pH measurements; go on to the next exercise.

EXERCISE #5:  MAKING pH DETERMINATIONS WITH THE "ACCUMET MODEL 5" pH METER

(17)  Obtain a sample of Standard Buffer solution known to be

pH 4.0 (in labeled tube near pH meter).

(18)  Immerse electrode in the solution, again being careful that its tip does not come into contact with bottom or edge of tube.

(19)  Rotate "function" selector to "pH" mode.

(20)  Allow a few seconds for the pH indicator to stabilize its reading.  (Do not make any adjustments to the "standardize" control knob.)  Once stabilized, record the pH measurement of the solution in your Tech. Facility notebook.  Does the meter read pH 4.0?

(21)  Switch meter back to "Standby", before removing the electrode from the solution.

(22)  Rinse off electrode with deionized water and dry with Kimwipe, as before (Step 7,8).  The electrode is again ready for next pH measurement.

(23)  Repeat Steps 17 - 22, except using tube containing Standard Buffer solution of pH 10.0.  Record reading obtained in Tech. Facility notebook.  Does the meter read pH 10.0?

(24)  Repeat  Steps 17 -20, except using tube containing a buffer solution "Unknown".  Record reading obtained in Tech. Facility notebook.

(25)  Switch pH meter to "Standby", and remove the electrode from the sample solution.

(26)  Rinse electrode and dry as before (Step 7,8).

EXERCISE #6:  TURNING THE "ACCUMET MODEL 5" pH METER OFF

(27)  When you are done using the pH meter, re-immerse the electrode in its storage solution (bottle labeled "for pH electrode").

(28)  Turn the cap ring hole from the "OPEN" position to the "CLOSE" position.

(29)  If the meter is still to be used during the rest of the day, leave the power strip "ON" and the pH meter on "Standby" mode.  If the meter will no longer be used, shut it "OFF" completely using the off switch on the power strip into which it is plugged.

NOTE ON ELECTRODE STORAGE:  According to the pH meter manual, between uses the combination electrode is stored immersed in a solution consisting of equal volumes (50/50) of saturated KCl and pH 7.0 buffer.  The manufacturer recommends that the fill-hole always be kept open to keep the junction free-flowing.  However, we have found that keeping the fill-hole open results in considerable leakage of the electrode electrolyte fill solution into the storage solution.  This then requires constant re-filling of the electrode, which is not very practical. Therefore, for our purposes in the Techniques Facility, when the pH meter is not in use, the fill-hole should be closed.  Note that this may differ from the electrode storage system used in a research lab that you may work in; be sure to check at the time!

ONE-POINT VS. TWO-POINT STANDARDIZATION

Standardization of the pH meter must be performed prior to any pH measurements being made on experimental solution.  In the next exercise you will be performing what is known as a 2-point Standardization.  This procedure differs from the previous one in that 2 different Standard Buffer solutions are used to Standardize the pH meter before pH determinations can be made on experimental solutions.

In the next exercise the 2-point calibration will be set between pH 4.0-pH 7.0.  The two points (pH values) are chosen to bracket the desired measuring range of the experimental solution. This narrows the pH measurement range closer to that of the experimental sample, and improves accuracy of the pH readings.  To do a 2-point Standardization on the Model 5 pH meter, the "standardize" control knob will be used to set the first pH point and the "slope" control knob to set the second.  NOTE:  When setting a 2-point Standardization, it is recommended that pH 7.0  be set as the first standard since, according to equipment specifications, at pH 7.0 the "slope" control will have no effect upon the reading.

EXERCISE #7:  TWO-POINT STANDARDIZATION OF THE ACCUMET

MODEL 5 pH METER:

(1)  Turn on pH meter by using the "ON" switch of power strip to into which meter is plugged.  Allow about 10 minutes for pH meter to warm up.

(2)  Check to see that pH meter "function" selector is on "Standby" mode before starting.

(3)  Lift the electrode carefully from its arm, just enough to be able to turn the cap ring hole from the "Close" position to the "Open" position.  The fill hole must be OPEN when pH measurements are being taken.

(4)  Check to see that the "temperature" control knob is at 25oC.

(5)  Obtain samples of two Standard Buffer solution at  pH 4.0 and pH 7.0 from beaker labeled "Standard Buffers/Unknown", near the pH meter.  Use these tubes "as is",  i.e., do not pour samples into other tubes.  They can be re-used many times.

(6)  Carefully lift the support arm holding the electrode so that the electrode is raised out of the bottle of electrode storage solution.

REMEMBER! - When handling the electrode, do so very carefully.  Even slight damage to the tip can make the electrode useless for pH determinations.

(7)  Rinse electrode off with deionized water, and dry carefully with a Kimwipe as previously.

(8)  Carefully lower the electrode first into the pH 7.0 buffer solution in plastic tube.  [REMEMBER! -Do not touch electrode tip to anything.]

(9)  Rotate "function" selector from "Standby" to "pH".

(10)  Set "slope" control knob to 100%.  Wait several seconds to allow the pH indicator to stabilize.

(11)  Gradually adjust the pH reading to 7.0 using the "standardize" knob.  The meter is now standardized at pH 7.0, the first of the two pH points.

(12)  Switch the "function" selector back to "Standby".

(13)  Raise electrode, then remove and re-cap the pH 7.0 Standard Buffer and return it to its original place.

(14)  Rinse off the electrode with deionized water and dry as before.

(15)  Obtain sample of second pH Standard Buffer solution, in this case pH 4.0.

(16)  Carefully lower the electrode into the pH 4.0 buffer solution in plastic tube, being careful not to touch the tip of the electrode to bottom or edge of tube.

(17)  Rotate "function" selector from "Standby" to "pH".

(18)  To set the second pH point, the "slope" control knob is rotated until the meter reads pH 4.0.  (Do not adjust the "standardize" control setting.)  The meter is now standardized at pH 4.0, the second of the two pH points.

(19)  Switch "function" selector back to "Standby".

(20)  Raise electrode, then remove and re-cap the pH 4.0 Standard Buffer and return it to its original place.

(21)  Rinse electrode as previously.

(22)  This completes the 2-point pH Standardization.  The pH meter is now ready for making pH measurements.

NOTE:  From this point on, when making pH determinations of sample solutions, take care not to change the position settings of either the "standardize" and "slope" controls.

EXERCISE #8:  MAKING pH DETERMINATIONS WITH THE "ACCUMET

MODEL 5" pH METER

(23)  Obtain a tube containing "Unknown" pH buffer solution.  (in labeled tube near pH meter).

(24)  Immerse electrode into the solution, being careful about the tip as before.

(25)  Rotate "function" selector to "pH".

(26)  Allow a few seconds for pH reading to stabilize.  (Do not make any adjustments to the "standardize" control or "slope" control knobs.)  Once stabilized, record the pH measurement of the solution in Tech. Facility notebook.

(27)  Switch meter back to "Standby", and remove electrode from the solution.

(28)  Rinse off electrode and dry with Kimwipe, as before.

TURNING THE "ACCUMET MODEL 5" pH METER OFF

(29)  When you are done using the pH meter, re-immerse electrode into its bottle of storage solution, (labeled "for pH electrode").  Turn the cap ring hole from the "OPEN" position to the "CLOSE" position.

(30)  If meter is still to be used this day, leave power strip "ON" and meter on "Standby" mode.

(31)  If meter will no longer be used, shut it "OFF" completely by turning off switch on its power strip.

(32)  Results:  Compare the pH reading obtained after the 2-point standardization with that obtained for the same sample when a

1-point Standardization was performed.  How close are the two results?

Ask your Tech. Facility supervisor for the correct pH of the unknown solution.  Which of the two Standardization procedures gives a more accurate reading?  Record observations in Tech. Facility notebook.

USING A DIGITAL pH METER

In the previous exercises (Exercises #3 - #8), you used an analog pH meter.  In the following exercise you will acquaint yourself with a digital pH meter, the Corning Model 340, which displays pH measurements as a liquid crystal display (or LCD) digital readout.  As previously, it is important that you familiarize yourself with the locations and functions of various panel buttons.

In the next exercise you will use the "Corning 340" digital pH meter to do a 2-point standardization on the same unknown used previously.  You already have an idea in what range the pH of the "Unknown" falls, based on your measurement using the "Accumet Model 5" Analog pH meter.  Therefore, when doing a 2-point standardization on the digital pH meter, the two points set will be pH 4.0 and pH 7.0.

EXERCISE #9:  USING THE CORNING 340 DIGITAL pH METER TO DO A 2-POINT STANDARDIZATION

(1)  Before starting, examine the pH meter electrode.  At the top of the electrode is a thin rubber strip covering the fill hole, which much be unplugged during meter use.

(2)  Unplug the fill hole stopper as follows:

(a) Hold the electrode firmly at the top (orange portion) with one hand.

(b) With other hand, grab end of grey rubber strip and pull outward.

Once the fill hole is unplugged, pH measurements can be made.

(3)  Obtain the sample of Standard Buffer solution at pH 4.0.

(4)  Carefully lift the arm holding the electrode, raising it out of the bottle of electrode storage solution.

(5)  Using the "squirt" bottle labeled "pH meter only", rinse the electrode with deionized water and dry gently with a Kimwipe. Do not squeeze electrode, especially electrode bulb tip, when wiping.

(6)  Carefully lower the electrode into pH 4.0 buffer solution in plastic test tube (NOTHING TOUCHING TIP!!).

(7)  Turn the pH meter "ON" by pressing "on/off" button; LCD display will light up.

(8)  Check to see that meter is in pH mode by pressing "mode" button.  The "pH" indicator will appear at lower right, under the main display.

(9)  To calibrate the pH electrode, press the "cal" button.

(10)  Immediately, press the "auto read" button, which will automatically freeze the display once a stable endpoint is reached.

(When the auto read mode is ON, the auto endpoint indicator [ ] appears in the upper right of the display.)

(11)  Wait until a stable endpoint measurement is reached.  At that time the flashing decimal point stops, the display freezes and a beep sounds.

The pH meter is now standardized at pH 4.0.  This is the first calibration point.

(12)  Do not turn pH meter "OFF" at this point; leave the display on "frozen".

(13)  Carefully raise electrode from tube; remove and re-cap the pH 4.0 Standard Buffer and put it back in its original place. Then rinse off the electrode with deionized water and dry as before.

(14)  To set the second calibration point obtain a obtain a sample of Standard Buffer solution at pH 7.0.

(15)  Carefully lower the electrode into the pH 7.0 buffer solution in plastic test tube (DO NOT LET TIP TOUCH ANYTHING!!).

(16)  REPEAT STEPS 9-11 (above).  The pH meter is now standardized at pH 7.0.  This is the second calibration point.

(17)  Do not turn pH meter "OFF" at this point, leave the display on the "freeze" mode.

(18)  Carefully raise electrode from tube; remove and re-cap the pH 7.0 Standard Buffer and put it back in its original place. Then rinse off the electrode with deionized water and dry as before.

(19)  The pH meter is now ready to make pH measurements.

MAKING pH MEASUREMENTS WITH THE

CORNING 340 DIGITAL pH METER

After performing "Standardization" (Steps 1-15), the meter is now ready to be used for making sample pH measurements. Since the pH meter is still in the "auto-read" mode, to start each new pH measurement, just press the "read" button.

(20)  Obtain a tube containing a sample of Buffer solution "Unknown" (in labeled tube near pH meter).

(21)  Immerse the electrode into solution (CAREFULLY!! EACH ELECTRODE COSTS >\$100!!).

(22)  Press "read" button to take pH measurement; allow a few seconds for the meter display to freeze automatically once the stable endpoint is reached, when a beep will sound.

(23)  Once stabilized, record the pH measurement of the solution in Tech. Facility notebook.

(24)  Do not turn pH meter "OFF" at this point, leave the display on the "freeze" mode.

(25)  Raise electrode, then remove and re-cap the "Unknown" Standard Buffer and put it back in its original place.  Then rinse off the electrode with deionized water and dry with kimwipe as before.  (Step 4)

TURNING THE CORNING 340 DIGITAL pH METER OFF

(26)  When you are done using the pH meter, immerse the electrode into the bottle of electrode storage solution, labeled "for pH electrode".

(27)  Cover up fill hole.

(28)  Check to see that pH meter is "OFF".

(29)  Clean up area around pH meter.

(30)  Compare the readings obtained for the same "Unknown" sample when a 2-point Calibration was performed using the digital vs. analog pH meters.  How close are the two results?

(31)  Check your Tech. Facility notebook (Exercise #8) for the correct pH of the unknown solution.  Of the two 2-point Standardization procedures, which gives a more accurate reading?  Record observations in Tech. Facility notebook.

pH BUFFERS

In general, according to the dictionary, a "buffer" is a kind of shield, a device or method that can protect against something.  With respect to pH, a buffer is a solution containing chemical compounds that prevent the pH (i.e., [H+]) from changing significantly when either acid or base is added to the solution.

EXAMPLE: NO BUFFER:

Several drops of HCl are added to a beaker containing pure water, with stirring.  The pH is measured and found to have fallen to 2.0.  Now several drops of NaOH solution are added with stirring, and the pH is found to be 10.0.

EXAMPLE: + BUFFER:

Several drops of HCl are added to a beaker containing "phosphate buffer", in this case a solution of NaH2PO4 and Na2HPO4.  The pH stays very close to 7.0.  Now several drops of NaOH solution are added with stirring, and it is found that the pH is still nearly 7.0.

What is going on here?

ACIDS AND BASES - A BROADER DEFINITION:

An understanding of buffers requires use of a broader definition of acids and bases (the "Bronsted-Lowry" definition).  An acid  = a proton donor (H+ = hydrogen ion = a proton), and a base = a proton acceptor.  The equivalent weight of an acid, by this definition, is that weight of the acid which yields 1 mole of H+, and the equivalent weight of a base is that weight of the base which accepts (or binds) one mole of H+.

When an acid is added to water, what actually happens is as follows:

HCl (hydrochloric acid) + H2O <--> H3O+ (a "hydronium ion") + Cl-

Here, according to the broader definition, the hydronium ion (H3O+) is an acid because it could, potentially, yield the proton it has accepted.  The Cl- (chloride ion) is, by definition, a base because it can, potentially, accept a proton after having yielded one.  The water (H2O) itself, by definition, is acting as a base because it accepts a proton!

In the above reaction, the equilibrium lies far to the right.  In other words, essentially all of the HCl will react with H2O to produce H3O+ and Cl-.  This is true for all strong acids such as HCl.  Correspondingly, the acid produced in the reaction, the H3O+ ("hydronium ion") is a weak acid.  If we wrote the reaction backwards (H3O+ + Cl- <--> HCl + H2O), the equilibrium would lie far to the left.

The base that is produced when an acid (proton donor) yields its proton is called the conjugate base of the acid.  Conversely, the acid that is produced when a base (proton acceptor) accepts a proton is called the conjugate acid of the base.  When an acid loses a proton, that acid plus the conjugate base produced are called a conjugate acid-base pair.  In the reaction written above, HCl and Cl- are a conjugate acid-base pair, and H2O and H3O+ are also a conjugate acid-base pair.

ANOTHER EXAMPLE:

H3PO4  (phosphoric acid) + H2O <--> H3O+ + H2PO4-

Here, again, the phosphoric acid is a relatively strong acid compared with the hydronium ion.  The H2PO4- (dihydrogen phosphate ion; "di" = 2) is the conjugate base of the phosphoric acid, produced by loss of a proton from phosphoric acid.

The H2PO4- (dihydrogen phosphate ion) can itself act as an acid, although a weaker one than phosphoric acid, as follows:

H2PO4- + H2O <--> H3O+ + HPO4-- (hydrogen phosphate ion)

Here, the HPO4-- (hydrogen phosphate ion) is the conjugate base of the H2PO4- (dihydrogen phosphate ion).  The hydrogen phosphate ion can in turn act as a very weak acid, giving up a proton to water to produce a hydronium ion plus a phosphate ion:

HPO4-- + H2O <--> H3O+ + PO4--- (phosphate ion)

How does a pH buffer work?  Such buffers are usually mixtures of weak acids and their conjugate bases (sometimes referred to as the salts of the weak acids).  Phosphate buffer, for example, is typically a mixture of dihydrogen phosphate ion and hydrogen phosphate ion, obtained by mixing specified amounts of NaH2PO4 (sodium dihydrogen phosphate, the weak acid) with Na2HPO4 (disodium hydrogen phosphate, its conjugate base or salt).  To be precise, such a buffer system would be called "sodium phosphate buffer.  Sometimes the potassium (K) salts are used, i.e., KH2PO4 plus K2HPO4, and this would be called potassium phosphate buffer.  If HCl is added, the H+ reacts with Na2HPO4 (the proton acceptor), producing more NaH2PO4 (the conjugate weak acid), which is already a component of the buffer.  In effect, free H+ ions are removed from the solution.  When NaOH is added, it reacts with NaH2PO4 (the weak acid) to produce more Na2HPO4 (the conjugate base, or salt), which is the other component of the buffer.  The effect is to remove free OH- from solution, again preventing any dramatic change of pH.

Can the other "H" left in Na2HPO4 similarly be yielded?  Yes!  If enough additional NaOH is added, Na3PO4 will be produced.  Similarly, going in the direction of lower pH, if enough HCl is added to NaH2PO4, H3PO4 will be produced.

EXERCISE #10:

Pick up a bottle of Na2HPO4 from the chemicals shelf in cabinet.  How is it described on the label?  Similarly, examine the label of a bottle containing NaH2PO4; how is it described?  What might the terms "monobasic" and "dibasic" refer to?  Note that you can remember which is which by starting with phosphoric acid in mind (H3PO4), which can give up as many as three protons and might be described as "nonbasic".

THE pH OF BUFFER SOLUTIONS:

The degree to which a weak acid will ionize in solution is measured by Ka, an equilibrium constant for the ionization reaction.

Ka = [H+][A-]/[HA],

where [HA] = the concentration of weak acid and [A-] is the concentration of ion (or, of salt) derived from the acid.  The higher the Ka is, the more complete the dissociation of the acid into ions.

This equation can be rearranged to describe the hydrogen ion concentration in terms of the other variables (for an illustration of the math involved, see note "a" at end of module).  First,

[H+][A-] = Ka[HA], and thus [H+] = Ka[HA]/[A-]

The equation can be further manipulated into a more useful form:

Taking the log of both sides, we get:

log [H+] = log Ka + log [HA]/[A-]

Converting to the negative log of both sides gives:

- log [H+] = - log Ka + log [A-]/[HA]

Shorthand notation for - log [H+] is "pH" (by definition and mutual agreement), and shorthand notation for - log Ka is "pKa".  Thus, re-writing the equation above,

pH = pKa + log [A-]/[HA]

This is known as the "Henderson-Hasselbalch equation".  It says that the pH is a function of the dissociation constant of the acid and the ratio of salt concentration [A-] to acid concentration [HA].  Note that when the salt concentration = the acid concentration, [A-]/[HA] = 1 and log [A-]/[HA] = 0 (i.e., log 1 = 0), so that the pH is then equal to the pKa.  The ability of a buffer to resist changes of pH is greatest when this condition exists.  Said another way, when the pH = pKa, a buffer will be most effective.  This means that one of the criteria for selecting a buffer for experimental work is that it should maintain the pH well (i.e., have its pKa) near or at the desired pH.

Looking at the Henderson-Hasselbalch equation above, since the pKa is constant, the pH will be a function of the ratio of [A-]/[HA].  You can see that as more base (such as NaOH) is added, more of the acid [HA] is converted to salt [A-], causing the pH to rise.  Conversely, as more acid is added, the pH will fall.

EXAMPLES:  Some weak acids commonly used to make buffers in research labs, and their approximate pKa values (published values for pKa very somewhat due to differences in conditions of measurement, ionic species and contaminants present, etc.):

Weak acid                    pKa     Typical buffer system

acetic acid (CH3COOH)        4.8  acetic acid + Na-acetate

carbonic acid (H2CO3)         6.4  carbonic acid + Na-bicarbonate

dihydrogen phosphate (H2PO4-) 7.2  sodium phosphate (monobasic) +

sodium phosphate (dibasic)

Some other commonly used buffer compounds and their pKa are:

"TRIS" (= tris hydroxymethylaminoethane), 8.0; "PIPES" = (= piperazine-N,N'-bis[2-ethane sulfonic acid]), 6.8; and "HEPES" (N-2-hydroxyethylpiperazine-N'-2-ethanesulfonic acid, 7.5.  PIPES and HEPES are two members of a series of related buffering compounds developed relatively recently.

Compounds that can yield (or accept) more than one H+ have multiple pKa values.  Phosphate buffers provide a good example.  H3PO4 (phosphoric acid) can donate its protons in a series of steps, as outlined above.  For H3PO4 and H2PO4- (conjugate base), pKa = 2.1; for H2PO4- and HPO4--, pKa = 7.2; for HPO4-- and HPO4--- (phosphate ion), pKa = 12.7.

EXERCISE #11:

Look at the tables and chart posted on the wall above the pH meter.  Note the useful pH ranges for the buffers listed.  These center about the pKa for each compound.

BUFFER CAPACITY

In addition to pKa, the ability of a buffer to resist changes of pH  will depend upon the concentration of the buffering chemicals.  Obviously, the pH of a very dilute buffer solution cannot resist addition of a given amount of acid or base as well as a concentrated buffer solution.  "Buffer molarity" = the concentration of the weak acid plus the concentration of the conjugate base.  If we dissolved 1 mole of NaH2PO4 (the weak acid) plus 1 mole of Na2HPO4 (the conjugate base) in 1 liter total volume, the buffer molarity would be 2M.  If, instead, we dissolved 0.1 mole of each in 1 liter, the buffer molarity would be 0.2M.  Adding, say, 0.4 moles of HCl to each buffer would overwhelm the capacity of the 0.2M buffer, but would be handled well by the 2M buffer.

HOW DOES ONE CHOOSE A BUFFER FOR EXPERIMENTAL USE?

Some buffer compounds may interfere with the biological processes under study, or they may introduce artifacts at some point of the procedure.  For example, phosphate buffer is often used to control the pH of fixative solutions in which cell structure is preserved for electron microscopy.  However, the phosphate ions can form a precipitate with intracellular calcium or other ions, causing insoluble granules to appear within the observed structure.  Special washes must be used to avoid this.

In general, a buffer is chosen because it has been used before for the same purpose.  Or, if no such experiment has been performed previously, a number of different buffers will be tested which have their pKa close to the pH desired.

PREPARING BUFFERS:

Having read through all of the above background material, you will be pleasantly surprised to learn that making up buffers for experimental use does not usually require calculations.  In most cases you will simply follow a recipe that tells you what stock solutions to prepare, and how to mix them.  You will then use the pH meter to check whether the pH has come out as expected, or whether it requires some adjusting.

In many cases, moreover, you do not need to prepare solutions of both the weak acid and its conjugate base!  You need only prepare a stock solution of the weak acid, and add a specified amount of strong base (usually NaOH), or, in some cases, of strong acid (usually HCl).  The reason for this is that addition of the strong base (NaOH) to the weak acid will produce the conjugate base of the acid.

Recipe for ACETATE BUFFER, useful pH range 3.6 - 5.6.

Volume:  1 liter; Concentration: 0.1M

Stock solutions: (a) 1N acetic acid, (b) 1N sodium hydroxide

In each case, a 1 beaker (with magnetic stirring bar) is partially filled with reagent-grade water (say, to about half-full), and then 100ml of the 1N acetic acid is added.  Next, a specified volume of the 1N NaOH is added, using the following table to achieve a particular final pH:

Table I. Acetate Buffer

pH--> 3.6  3.8  4.0  4.2  4.4  4.6  4.8  5.0  5.2  5.4  5.6

vol. (ml)

1N NaOH)--> 7.5 12.0 18.0 26.5 37.0 49.0 60.0 70.5 79.0 85.5 90.5

Water is added to about 900 ml with gentle stirring.  The pH meter can then be used to check that the pH is as expected.  The solution is then transferred to a 1 liter volumetric flask, and the flask is filled to the 1 liter mark with water, including rinses of the original beaker and electrode.  The solution is either mixed by inversion of the flask many times (end covered by Parafilm), or is mixed after pouring it into a reagent storage bottle.  To check that the final addition of water bringing the volume to 1 liter had no effect on the measured pH, the pH of a small sample can be measured.

Note that, regardless of the final pH, the dilution of stock 1N acetic acid is the same: 100 ml --> 1000 ml, i.e., a dilution factor of 1/10.

Recipe for POTASSIUM PHOSPHATE BUFFER, useful pH range 5.8 - 8.0.

Volume:  1 liter; Concentration: 0.1M

Stock solutions: (a) 1M KH2PO4 (potassium phosphate - monobasic),

(b) 1N KOH (potassium hydroxide)

This buffer uses potassium hydroxide rather than sodium hydroxide.

A 1 liter beaker (with magnetic stirring bar) is partially filled with reagent-grade water (about half-full), and then 100ml of the 1M KH2PO4 stock is added.  Next, a specified volume of the 1N KOH is added, with stirring, using the following table to achieve a particular final pH:

Table II. Potassium Phosphate Buffer (0.1M)

pH --> 5.8     6.2     6.6    7.0     7.4     7.8     8.0

vol. (ml)

1N NaOH)-->  7.32   17.10   35.48  59.08   78.68   90.34   93.70

Water is added to about 900 ml with gentle stirring, and the buffer solution completed as above.

Note again that, regardless of the final pH, the dilution of stock 1M potassium phosphate is the same: 100 ml --> 1000 ml, i.e., a dilution factor of 1/10.

EXERCISE #12: Using the information in Table II, prepare 100 ml of 0.1M potassium phosphate (KH2PO4) buffer pH 6.2 and 100 ml of the same buffer at pH 7.4.  You will use these buffers in the spectrophotometer exercises.  Store each in a properly labeled bottle in the refrigerator.

OTHER APPROACHES:

In some buffer recipes, stock solutions are not made up, but rather the amounts of each ingredient are specified by weight.

EXAMPLE:

Recipe for SODIUM PHOSPHATE BUFFER, useful pH range 5.8 - 8.0.

Volume:  1 liter; Concentration: 1/15 M (= 0.067M).

Stock solutions: none

In each case, a 1 liter beaker is partially filled with reagent-grade water (say, to about half-full), and then the amounts of NaH2PO4 (sodium phosphate - monobasic) and Na2HPO4 (sodium phosphate - dibasic) specified in the following table (Table III) are dissolved in it (magnetic stirring).  The solution is then completed, with pH measurement, as above.

WATER OF HYDRATION: In Table III, the weight of NaH2PO4 given is for the "monohydrate", i.e., NaH2PO4.H2O.  This means that there is one mole of water of hydration per mole of NaH2PO4, so that, when the weights were calculated, the formula weight of water was added to the formula weight for the NaH2PO4.  This is water that crystallizes out with the compound, even though not part of its molecular structure.  If you are using this table to make up the buffer, you would have to be sure that you were weighing out the same chemical, i.e., NaH2PO4.H2O.  The waters of hydration are given on the bottle of chemicals in the formula weight.  The number of waters of hydration can vary for a given salt, so that you should always check for a match before using what you think is the same chemical specified in a recipe.

Table III. Sodium Phosphate Buffer (1/15 M)

pH --> 5.8     6.2     6.6    7.0     7.4     7.8     8.0

monobasic -> 8.442   7.545   5.773  3.588   1.794   0.805   0.506

NaH2PO4.H2O

(grams)

dibasic  --> 0.781   1.704   3.526  5.774   7.620   8.637   8.945

Na2HPO4

(grams)

EXERCISE #13:

Using Table III, above, prepare 100 ml of 1/15 M sodium phosphate buffer at pH 7.4.  Before starting, write down the procedure to be followed (in your Tech Facility notebook).  What volume of buffer is made up from the amounts listed in the above table?  What volume are you making up?  What amounts of monobasic and dibasic salts will you be weighing out?  To what decimal place can you weigh these amounts using the electronic balance?  Note that only the monosodium phosphate has water of hydration.  Obtain the bottles of monobasic and dibasic sodium phosphate from the chemicals cabinet, and compare their chemical formulas.  Using the following atomic weights (=rounded-off values from the periodic chart of elements on the wall), calculate the formula weight for NaH2PO4, NaH2PO4.H2O and Na2HPO4: Na=23, H=1, P=31, O=16.  Do your calculated formula weights agree with those given on the bottles?

Proceed to make the solution, following the steps as outlined above and modified in your notebook protocol.  Did the pH come out as expected?

EXERCISE #14:

You can prepare a buffer even if you do not have a Table to follow.  Repeat the exercise just above, except this time do not check Table II to see how much 1N KOH to add.  Simply add the water and potassium phosphate stock to the beaker, and then add 1N KOH in measured volumes (using pipette) while continuously stirring and continuously reading the pH.  Make sure that the electrode is off to one side of the beaker, so that you do not accidentally hit it with the pipette!

Stop adding the 1N KOH when you reach pH 6.6.  What volume have you added?  Does it match that in Table II?  Now, continue adding the 1N KOH until you reach pH 7.0.  What is the total volume added now?  Does it match that in Table II?

In practice, buffers can be prepared routinely in the laboratory in this way.  The pH of the measured amount of weak acid or of weak acid plus other ingredients is adjusted to the final pH desired by continuously adding NaOH or KOH solution in some cases) and monitoring pH.

Note "a":

The math here can be illustrated as follows:

Original form of the equation Ka = [H+][A-]/[HA] is: a = b x c/d; then log a = log b + log c/d, and -log a = -log b + log d/c.

Example: let a = 1000, b = 10, c = 200, and d = 2

Original equation: 1000 = 10 x 200/2;

then log 1000 = log 10 + log 100 [i.e., 3 = 1 + 2], and

-log 1000 = -log 10 + log 2/200  [i.e., -log 10 + log 1/100

= -1 + log 10-2 = -1 + (-2)]