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LABORATORY ONE

INTRODUCTION TO LABORATORY EQUIPMENT AND TECHNIQUES

The purpose of this laboratory is to introduce you to the equipment and techniques

that you will use over the course of the term in Kinesiology 307. Good scientific techniques are necessary to ensure both reliable and valid results. Care should be taken in each experiment that follows to ensure that you know the proper handling and accurate use of each piece of equipment. This will only be attained with repeated use of the equipment and strict attention to proper procedures.

The preparation and handling of chemical solutions are an essential part of experimental biochemistry. Both accuracy and precision are important in the biochemical sciences. Accuracy - is concerned with the extent to which the measured value represents the true score. More often than not, the true score is not known. The average of an infinite number of trials is said to be the best estimate of the true score. The average can be used to determine the constant error by subtracting the true score. The constant error can be expressed as percentage of the true score.

Precision - is defined as the agreement of a group or series of measurements of the same thing. When measurements of the same thing agree or when the measured values are clustered closely together, then the measure is said to be precise. Standard deviation(s) or coefficient of variation (c = s/x_) are often used to estimate the precision of a measurement. The standard deviation can also be expressed as a percentage of the true or expected value. A. TRANSFERRING AND DISPENSING VOLUMES - PIPETTING

One specific purpose of the laboratory is to transfer accurately a known amount of material from one container to another. A lack of accuracy and reliability in pipetting skills is the most frequent cause of frustration in perfecting and refining laboratory techniques. There are many pieces of glassware which can be used to transfer solutions. You will be introduced to several of these. Each piece of glassware has a specific use and a certain amount of error associated with it. Beakers - come in various sizes. They are very inaccurate as a piece of measuring glassware.

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Graduated Cylinder - a piece of calibrated glassware used to measure liquid volumes directly. Depending on the size and volume graduated cylinders are normally used to measure volumes within 0.5 mL. The volume of liquid is always read at the bottom of the meniscus. Pipets - a piece of calibrated glassware used to accurately measure and dispense volumes of liquids. Precision of experimental results usually relies on the precision of the pipets used. Therefore in selecting a pipet, you should select one that will provide the greatest precision for the required conditions of the experiment. 1. Volumetric Pipets - are calibrated to deliver only one volume. They are usually

calibrated in smaller units. The blow-out pipets are usually graduated to the tip. Others with a band around the top indicate that they too are graduated to the tip and that they should be left to drain in order to dispense the required volume.

2. Measuring Pipets or Transfer Pipet (Mohr) - are graduated to deliver varying

volumes of liquid. Mohr varieties are graduated to deliver volumes of liquid between the calibration marks on the outside of the pipet.

3. Air Displacement Pipets (Automatic (Eppendorf) pipets) - used to transfer

either set or adjusted volumes from one container to another. They are usually calibrated in smaller units. Depressing the plunger at the top of these pipets causes air to be displaced in the tip. The plastic tip is then placed in the solution and the plunger is allowed to return to its original position. This draws the volume of solution into the pipet tip. They have automated pipetting and when used properly, are efficient. They are particularly useful in pipetting toxic or poisonous solutions.

4. Syringe Pipets (Positive Displacement Pipets) - operate like a syringe leaving

no air bubbles between the liquid and the plunger. They are particularly useful when pipetting liquids which are difficult to pipet by traditional techniques or using other automatic pipets, such as very viscous liquids.

5. Eppendorf Repeater (Repeater) - is a hand-held device for repetitive pipetting. It

eliminates filling the pipette with the reagent for every single pipetting operation. It consists of the Repeater, and the Combitip, and a reusable plastic reservoir which locks into the bottom of the Repeater. Five different volumes may be dispensed per Combitip, and this is set by the volume selection dial on the Repeater. The volume selection chart should be consulted when determining the dial selection and Combitip to use for the volume that you wish to pipet.

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B. PREPARATION OF SOLUTIONS: EXPRESSION OF CONCENTRATION AND

DILUTION

The three most useful expressions of concentration are molarity, molality and normality. 1. Molarity (M) is the number of moles of solute dissolved in a litre of solution. This

is the most common expression of concentration for solids of known molecular weight. The solute is weighed on an analytical balance and the solution volume is measured using a volumetric flask.

a. Molarity

M = moles of solute

Litres of solution

Eg. A 1M (molar) solution of NaCl contains 58.5 g of NaCl per litre of solution.

How would you make up 250 mL of a 0.3 M solution of NaCl?

First you must determine the number of grams of NaCl to weigh out? One method of determining this is using the factor label method.

Start with what you what to determine: I want to know the #? g NaCl to weigh.

And also with what you know: I know that there are 58.5 g of NaCl in a mole of NaCl and I also know that 1M equals 1 mole per Litre of solution.

Put what you want to find on the left of the equal sign, start with what you know and find the appropriate conversion factors so that you are left with the correct units that you wish to determine.

#?g NaCl = 58.5 g NaCl * 0.3 mole * 1L * 250 mL

Mole L 1000mL = 4.3875 g NaCl

Notice that in the example above that the units cancel and that we are left with g.

We would therefore need to weigh out 4.388 g NaCl. Place this in a graduated cylinder and add the required volume of dH2O to make a final volume of 250 mL.

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2. Molality is the number of moles of solute per 1000 grams of solvent and is denoted by m. Molality can be more precise than molarity since both the solute and the solvent are weighed, however it is not widely used in biochemistry.

3. Normality (N)

One litre of 1 M H2SO4 contains twice as many hydrogen ions as does 1 litre of 1 M HCl. It is convenient to deal with solutions that contain the same number of hydrogen ions or hydroxide ions per litre. Normality of an acid is the number of moles of hydrogen ion present per litre of solution and the normality of a base is the number of moles of hydroxide ion per litre of solution. For monoprotic acids such as acetic acid, which yield one mole of hydrogen ions upon complete dissociation: 1 mole of CH3COOH is a 1 normal (1 N) solution as well as 1 M solution.

CH3-COOH <------------>CH3-COO- + H+

For diprotic acids, such as furmaric acid, that yield two moles of protons when reacted completely with a base, it follows that: 1 M of furmaric acid is a 2 N solution, since two moles of H+ ions are supplied by each mole of furmaric acid.

HOOC-CH2 =C-COOH <------------> -OOC-CH2 =C-COO- + H+ + H +

A solution containing 1 mole of KOH per litre is a 1 N solution as well as a 1 M solution. A solution that contains 1 mole of H2SO4 per litre is a 1 M solution. However, it is a 2 N solution, since two moles of H+ ions are supplied by each mole of sulphuric acid. One mole of any 1 N acid solution will exactly neutralize the same volume of a 1 N base. The equivalent weight of an acid is that weight of the substance which furnishes 1 mole of H+ and for a base that furnishes 1 mole of OH-. Thus the equivalent weight of NaOH is 40/1 = 40 g H2SO4 is 98/2 = 49 g. equivalent weight = molecular weight

number of H+ or OH- per molecule

Normality is therefore defined as the number of equivalent weight in a litre of that solution (eq. weights/litre). A convenient way to calculate the amount of one solution required to neutralize another is:

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Nacid x volume of acid = Nbase x volume of base

It follows that 1 L of 5 N HCl will exactly neutralize 2 Litres of 2.5 N NaOH.

The gram-equivalent weight of an oxidizing or a reducing agent in an oxidation - reduction (redox) reaction is the amount of compound that will transfer one mole of electrons. In the reduction of nicotinamide adenine dinucleotide (NAD+), two moles of electrons are transferred per mole of NAD+. Therefore, the gram-equivalent weight of NAD+ is half its molecular weight. NAD+ + 2e- + H+ <---------> NADH

4. Percent Solution

When it is not necessary to know the number of moles of solute in a solution, percent solution may be used (particularly, if their molecular weight is unknown). Percent solution refers to the relative amounts of solute and solvent. There are three types of percent-of-solution used: volume per volume (v/v), weight per volume (w/v) and weight per weight (w/w). A solution that contains 5.5 mL of liquid solute in 100 mL of solution is a 5.5% (v/v) solution and 5.5 g of solid solute in 100 mL of solution is a 5.5% (w/v) solution and 5.5 g of solid solute in 100 g of solution is a 5.5% (w/w) solution. The most commonly used expression is the weight per volume, i.e., 5.5 g of solid solute diluted to 100 mL with solvent (w/v). a. Weight-to-Volume

%(w/v) = g of solute X 100 mL of solution

E.g. 10 g of NaCl in sufficient volume of dH2O to make a final volume of 100 mL is a 10% (w/v) solution. b. Volume-to-volume

%(v/v) = mL of solute X 100 mL of solution

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c. Weight-to-weight

%(w/w) = g of solute X 100 g of solution

5. Making Dilutions

Frequently, standards or extracts are obtained in a concentrated form and must be diluted to a less concentrated form in an assay. Dilute solutions are not as stable as concentrated ones, and therefore it is common place to prepare and store concentrated stock solution and make working dilute solutions as they are needed. There are several

ways for indicating the method of dilution. For example, A1-to-5 dilution@ can indicate

two different cases. First, one part of the original solution could be diluted with four parts of solvent to give a final diluted volume of five parts. This dilution is 1/5, since the concentration of the diluted solution is 1/5 the concentration of the original solution.

However, A1-to-5" also could indicate that one part of the original solution is diluted with

five parts of solvent to give a solution with 1/6 of the original concentration. The former convention appears to be more suitable and is more frequently used because the dilution factor is immediately apparent. In this case, it is best to read Adilute A1:5" as Adilute

one part solution with solvent to give five parts of total volume.@ We will use this

method for future laboratories to avoid confusion.

Serial dilutions involve the systematic dilution of an original Astock solution@ in a

series of fixed steps, such as: 1:10, 1:100, 1:1000. In this example 1 mL of the original stock solution (1M) might be diluted to 10 mL with 9 mL of distilled water to yield a solution that is one 10th as concentrated as the stock solution, ie., 0.1M solution. 1 mL of this 0.1 M solution would then be diluted in 10 mL of distilled water to yield a solution with a concentration of 0.01 M. Extra care should be taken when performing serial dilutions because error made in an early part of the process will be carried to all subsequent dilutions.

It is common practice in laboratories to purchase concentrated stock solutions of reagents and to dilute these with water to achieve the desired concentrations. This is accomplished by adding solvent to the concentrated stock solution to lower the original concentration. Eg. Given that you purchase 100 mL of 5.0 M stock glucose solution, how could you lower its concentration to 0.5 M? The question is, how much solvent should you add:

Remember;

final concentration = initial concentration X dilution factor

OR

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dilution factor = final concentration

initial concentration

In the above example: dilution factor = 0.5 M 5.0 M = 0.1 How do we use the above dilution factor? A dilution factor has two meanings. It is the factor by which a concentration is reduced and it is also the ratio of initial volume to final volume. If the entire original volume of 100 mL is used in the dilution, then this volume must be increased ten times by adding solvent.

dilution factor = initial volume

final volume

Rearrange:

Final volume = initial volume

dilution factor = 100 mL

0.1 = 1000 mL

Therefore, in this example you must add enough water to the 100 mL stock solution to bring its final volume to 1000 mL. We can summarize the above steps by the following equation:

final concentration = initial concentration X initial volume

final volume OR initial concentration = final volume final concentration initial volume

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C. pH METRE, pH and BUFFERS Hydrogen ions in a solution arise from the dissociation of an acid. The generalized equation to describe this is:

HA ----------------------> H+ + A-

Bronsted and Lowry have defined an acid as a proton donor, and a base as a proton acceptor. Consider the reaction of a strong acid such as HCl in water:

HCl + H2O --------------> H3O+ + Cl-

The HCl acts as proton donor and is therefore an acid. H2O accepts a proton and is therefore a base. Therefore the concentration of H+ ions or hydronium (H3O+) in the solution increases. Since HCl is a very strong acid, the above reaction goes essentially to completion in a dilute solution. A weak acid such as acetic acid is a poor donor of protons. Such a weak acid can donate protons to the water only to a limited extent, and in the resulting equilibrium, the concentrations of the undissociated reactants are much greater than the concentration of the products. The reaction of acetic acid with water can be expressed as:

CH3COOH + H2O ========= H3O+ + CH3COO-

Because this is an equilibrium reaction, we can write:

[H3O+] [CH3COO-] Ka = ----------------------- [CH3COOH]

Where Ka is the acid dissociation constant and is defined as:

[H+] [A-] Ka

== -------------

[HA]

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From this we can derive the Henderson-Hasselbalch equation, or the buffer equation, which relates the pH of a solution to the pKa of the acid and the relative concentrations of undissociated acid and the conjugate base forms:

[HA] [H+] = Ka

= x -------

[A-]

The above equation can be converted to the logarithmic form, to obtain:

[HA] log [H+] = log Ka

= + log ---------

[A-]

Multiply each term by -1, gives:

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

= - log --------

[A-]

Since pH = the -log [H+] and pKa >= - log Ka=, we can substitute:

[HA] pH = pKa

= + - log -------

[A-]

Inverting the log [HA]/[A-] term, the equation becomes:

[A-] pH = pKa

= + log --------

[HA-]

OR

[Conjugate Base] pH = pKa

= + log ---------------------------

[Undissociated Acid]

This is the Henderson-Hasselbalch equation. It is useful in the preparation of buffers and also in understanding how the concentrations of undissociated acid and the conjugate base of a weak acid affect the pH.

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Water Water is a very unique material because it can act as both a Bronsted acid and a Bronsted base:

H2O + H2O ====== H3O+ + OH-

Here, one molecule of water donates a proton to another molecule. This reaction occurs

to a very small extent. At 25C, pure water has equal concentrations of H3O+ ions and OH- ions.

[H3O+] = [OH-] = 1 X 10-7 at 25C

[H3O+] x [OH-] = 1 X 10-14 at 25C

This expression holds for any aqueous solution at 25C, whether other ions are present or not. The Concept of pH The concept of pH was developed as a convenient means of expressing the concentration of hydronium ions in dilute solutions of acids or bases. The pH of a solution is defined as the negative logarithm of the hydronium ion concentration of that solution. The pH scale runs from 0 to 14; the lower the number the more acid, the higher the more basic. A pH of seven is considered neutral.

pH = -log [H3O+] or -log [H+]

[H+] (M) 1 10-2 10-4 10-6 10-8 10-10 10-12 10-14

---acidic-------------B -neutral-B -------basic-------

pH 1 2 4 6 7 8 10 12 14

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As an example, the pH of a 0.1 molar solution of HCl is:

-log 0.1 = -log 1 X 10-1 = 1

The pH of a solution of 0.01 molar NaOH (a strong base) is worked out as follows:

[OH-] in 0.01 molar NaOH is 0.01 moleslitre-1 (1 X 10-2) Kw = [H3O+][OH-] = 1 X 10-14 [H3O+] = 10-14 = 1 X 10-12 10-2 pH = -log 10-12 = 12

The pH of a 0.033 molar solution of HCl is:

pH = -log (0.033) = -log (3.3 x 10-2)

= -(log 3.3 + log 10-2) = -(0.52 - 2.0) = 1.48

Neutralization Reactions The reaction of an acid with a base is called a neutralization reaction. For example, if a solution of NaOH is added to a solution of HCl, the acid and base neutralize the effects of each other. NaOH in water contains Na+ ions and OH- ions. When this is added to the HCl solution that contains H3O+ ions and Cl- ions the following equation results:

H3O+ + Cl- + Na+ + OH- ---------- 2H2O + Na+ + Cl-

If we cancel out the Na+ and Cl- that appear on both sides of the equation, we are left with:

H3O+ + OH- ----------- 2H2O

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Buffers Buffers are solutions that resist changes in pH when a small amount of acid or base is

added. A buffer=s chemical composition therefore, must be such that it can neutralize

both added acid and base. A buffer usually consists of a mixture of a weak acid and its salt or a weak base and its salt.

CH3COONa / CH3COOH

NH4Cl / NH3

A solution of CH3COOH and CH3COONa is a typical buffer. This solution can be represented:

CH3COOH + H2O ====== CH3COO- + H3O+

CH3COONa ----------- Na+ + CH3COO-

In this mixture, the molecular acetic acid is in equilibrium with hydronium and acetate ions, although the equilibrium is shifted more toward molecular (undissociated) acetic acid by the presence of excess acetate ions from the sodium acetate. When a base such as NaOH is added to this buffer solution, the added OH- ions react with H3O+ ions present to yield water.

OH- + H3O+ ----------- 2H2O

Some of the molecular acetic acid then dissociates to replace the hydronium ions that were used up and the pH remains the same. If a small amount of an acid such as HCl is added, the added hydronium ions react with some of the excess acetate ions present to give molecular acetic acid.

H3O+ + CH3COO- ----------- CH3COOH + H2O

Again, the pH remains essentially unchanged. Amino acids and proteins are excellent buffers because they have both acid and basic groups that can eliminate added base or acid:

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-COO- + H3O+ ----------- -COOH + H2O

-NH+3 + OH- ----------- + H2O

The monoamino, monocarboxylic amino acid in a neutral solution exists as a Zwitterion. It has a positive and negative charge.

COO- carboxylate group

ammonium group H3N+B C --- H

CH3 ALANINE (ALA)

Now if we have alanine in a solution at pH 7 and we added H3O+, then the reaction would be:

COO- COOH

∣ ∣ NH3

+ – C -- H + H3O+ -----------⟩ NH3+– C -- H + H2O

∣ ∣ CH3 CH3

The carboxylate group (COO-) is acting as a base because it is accepting a proton (H3O+) and becomes a carboxyl group (COOH). If the alanine, dissolved in water at pH 7, had some NaOH added, the reaction would be:

COO- COO-

NH3+B C B H + OH- ----------- NH2 B C B H + H2O

CH3 CH3

The ammonium group (NH3+) acts as an acid because it donates a proton (H+) to OH-, producing an amino group (NH2).

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From the Henderson-Hasselbach equation;

[A-] pH = pKa

’ + log ---------- [HA-]

it can be seen that when the [conjugate base] = [dissociated acid], the pH of the solution is equal to the pKa of the buffer.

pH = pKa

= + log 1.0 since log 1.0 = 0, then the pH = pKa=

When the [conjugated base] = 10 x [undissociated acid], then:

log [conjugated base] = log 10 = 1.0 [undissociated acid] 1 therefore; pH = pKa

= + 1.0

When the [conjugated base] = 1/10 x [undissociated acid], then:

log [conjugated base] = log 1 = -1.0 [undissociated acid] 10 therefore; pH = pKa

= - 1.0

Therefore, buffers are most effective in the range pH = pKa= 1. Outside that range,

either the concentration of undissociated acid is too small to effectively buffer added base, or the concentration of conjugate base is too small to effectively buffer added H+ ions. Once the desired pH range for an assay is determined then a buffer is selected on the basis of the pK value. There are many buffers that can be used to control the pH in the

range of pH = 7.0 eg., For example, Tris [tris(hydroxy-methyl)aminomethane] has a pKa=

= 8.3. Temperature affects the dissociation of some weak acids and bases the pKa= will

change with temperature.

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Properties of Some Common Biological Buffers

Buffer Name

pKa= at 20C

pKa= /C

Phosphoric acid Citric acid Formic acid Succinic acid Sodium dihydrogen citrate Acetic acid Disodium hydrogen citrate Sodium hydrogen succinate MES [2-(N-morpholino)ethanesulfonic acid] ADA [N-2-acetamindo)iminodiacetic acid]

PIPES [piperazine-N,N=-bis92ethanesulfonic acid)]

Imidazole hydrochloride

BES [n,n=-bis(2hydroxyethyl0-2-aminoethanesulfonic acid] MOPS [3-(N-morpholino)propanesulfonic acid] Sodium dihydrogen phosphate

HEPES [N-2-hydroxyethylpiperazine-N=-2-ethanesulfonic acid] Tricine [N-tris(hydroxymethyl)methylglycine] Glycine amide hydrochloride Tris [tris(hydroxymethyl)aminomethane] hydrochlorine

Bicine [N,N=-bis(hydroxyethyl)glycine]

Glycylglycine Boric acid Disodium hydrogen phosphate

2.12 3.06 3.75 4.19 4.74 4.75 5.40 5.48 6.15 6.6 6.8 7.0 7.1 4.20 7.21 7.55 8.15 8.2 8.3 8.35 8.4 9.24 12.32

B B B B B B B B

-0.011 -0.011 -0.0085 B

-0.027 B B

-0.014 -0.021 -0.029-0.031 -0.018 -0.028 B B

eg. Tris at 10C = pKa= = (10C - 20C) (-0.031) = 0.31; pKa= = 8.3 + 0.31 = 8.6

Tris at 30C = pKa= = (30C - 20C) (-0.031) = -0.31; pKa= = 8.3 - 0.31 = 8.0

Buffer Capacity and Ionic Strength The molar strength of the two components of a buffer should be chosen to provide adequate buffering capacity. The total ionic strength of the buffer may be important also for some assays. Buffer strength is expressed in terms of the total concentration of conjugate base or conjugate acid.

CH2OH CH2OH HOCH2 – C – NH3

+ ⟨----------------⟩ HOCH2 – C – NH2 +H+ CH2OH CH2OH (I) (II) Tris pKa

’ = 8.3 at 20̊C

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Since the pKa= of Tris ([tris(hydroxymethyl)aminomethane] hydrochlorine) is 8.3, it acts

as an effective buffer between pH 7.3 and 9.3. An assay that requires 50 mM Tris buffer pH = 7.8 can be made from 6055 mg of Tris in its free base form (II) (FW = 121.1 g) or the hydrochloride salt.

50 millimole x 121.1 mg = 6055 mg Tris 1 millimole

The pH of solution would have to be adjusted to reach the required 7.8 by adding acid. Hydrochloric acid is the best choice because Tris-HCl is soluble and chloride ions are compatible with most biological systems. Henderson-Hasselbach equation can be used to determine the amount of HCl required to adjust the pH to 7.8.

eg. log [conjugate base] = pH - pKa’

[acid] log [conjugate base] = 7.8 - 8.3 [acid] log [conjugate base] = -0.5 [acid] OR log [acid] = 0.5 [conjugate base]

[acid] = 3.16 [conjugate base] Since we started out with 50 millimoles of the conjugate base form of Tris, the final ratio would be:

[acid] = 3.16 50 mM - [ acid]

[acid] = 158 mM - 3.16 [acid] 4.16 [acid] = 158 mM [acid] = 37.89 mM or 38 mM

Therefore, 50 mmol of Tris plus 38 mmol of HCl diluted to 1 L would provide a 50 mM Tris

buffer, pH 7.8 at 20C. However, the final pH should be verified using a pH metre.

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The solubility, stability and interactions with other ions in the system as well as the light absorption, potential for inhibition or stimulation are important considerations when choosing a buffer system. Tris (tris-[hydroxymethyl]-aminomethane) is a commonly used buffer in biochemistry because it has a high buffering capacity, low toxicity and it does not interfere with most biochemical reactions or interact with metal ions. However, Tris can form a thin film on the glass of certain types of electrodes causing errors in the pH reading. The film has to be removed by treating the electrode with detergent or acid. Manufactures will usually advertise electrodes that can be used with Tris. Tris buffers usually vary more with temperature than most buffers do. pH decreases approximately 0.3 pH units per

degree Celsius from 25C to 5C. pH Measurement The pH metre is a device, used to measure the pH of a solution. The pH metre is simply a potentiometer used to measure the [H+] in a solution. It actually measures the electrical potential between two electrodes placed in the solution. The pH metre records the electrical potential difference between a reference electrode and a measuring electrode that is sensitive to [H+]. The measuring electrode has a selective permeable membrane that is permeable to H+ but not to other cations and anions. This results in the creation of a potential difference across the membrane that varies linearly with pH. The most commonly used pH-dependent unit is the glass electrode. A glass electrode usually consists of a glass electrode containing 0.1 N HCl in contact with the H+ permeable membrane. This is connected to the voltmeter by a silver wire coated with silver chloride which is immersed in the HCl. This in turn is connected to a reference electrode which commonly contains a Hg-HgCl2 paste or Ag-AgCl2 in saturated KCl. The reference electrode is encased in a glass tube which is impermeable to H+ ions. The voltage measured is the result of the difference between the glass and the reference

electrode. Many electrodes are Acombination electrodes@.

They contain both the reference and the measuring electrode incorporated into one slim glass or gel tube. If a thin layer of this glass separates two solutions with different H+ concentrations, H+

ions will move across the membrane from the solution of high concentration to that of low H+ concentration. The movement of the H+ ions across the membrane adds a positive ion to the solution of lower H+ concentration and leaves behind a negative ion, thereby developing an electric potential across the membrane. The magnitude of this potential difference is measured by the voltmeter (pH metre).

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The magnitude of this potential is given by the Nernst equation:

E = constant + 2.303 RT x pH

F E = magnitude of the potential; R = gas constant; T = absolute temperature, F = Faraday constant

The constant in the equation above depends on a number of factors and varies from electrode to electrode. In order to avoid determining the constants in the equation above, and because the concentration of HCl in the glass electrode can change with repeated use, the pH metre is standardized using solutions of known pH. The pH reading of the metre is adjusted to read the pH of the buffer standard solution(s), thereby eliminating the constant from the pH reading. The measured potential is a function of the temperature. The measured potential is therefore, affected by the temperature of the solution. Making pH Measurements The pH electrode is rinsed with dH2O and blotted dry. The metre is standardized with at least two buffers, one with a pH below and one with a pH above the pH of the solution to be measured. Purchased buffers (4.0, 7.0 10.0) are used to calibrate the pH metre. A standard calibration requires the measurement of the response of the electrode in the two pH buffer solutions and creating a linear map of the electrode response to these two

points. The pH buffer pH 7 should provide a reading around 0 - 30 mv, pH buffer 4 approximately 160 - 180 mv greater than the pH 7 buffer and pH buffer 10 approximately 160 - 180 mv less than the pH 7 buffer. This results in offset and slope correction factors, where the offset is the mV reading at pH 7 and the slope is the change in mV response per pH unit, usually expressed in mV/pH as a percentage of the ideal slope of the

electrode (59.16 mV/pH at 25 C). Usually the day to day changes in the offset and slope are monitored in a laboratory. In this way, errors in calibration can be identified, since the electrodes are not expected to age suddenly. For best results, the sample and the buffers should be allowed to equilibrate to the same temperature. The electrode should be cleaned prior to use.

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Figure 1: Relationship between electrical potential (mv) and pH

Digital Analysis Corporation http://www.phadjustment.com/pH_Systems.html

Types of Electrodes Although purchasing an electrode can be confusing because of the many advertised in a

manufacturer=s catalogue, there are really only a few different types and combinations

of these: Glass (H+ permeable) electrodes - either general purpose, high temperature, low Na+ electrodes (high pH electrode) and Na+ ultra-sensitive (used to measure Na+ concentration). Note: usually general-purpose electrodes will be somewhat permeable to Na+. If H+ and Na+ both diffuse across the membrane, then the H+ concentration will appear to be greater than it is, and the pH will be lower than the actual pH. This can be more noticeable at high pH when NaOH is present. Certain types of borosilicate glass are permeable to H+ but not to other cations or anions. Reference (H+ permeable) electrodes - either general purpose, high temperature or chloride free.

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Combination electrodes - the glass and reference electrode are contained in a single unit. Smaller volumes can more easily be measured. However, these can be more expensive. Other electrodes - used for measurements other than pH. E.g., metallic electrodes measure redox potential, radiometer measures pO2, pCO2, lactate and glucose analysers. pH paper - various colour indicator papers for pH determinations. Some are sensitive to

0.2 pH units. However, these are usually used to provide a rough Aball park@ of the

solution pH. They can be affected by other salts in the solution. Alternatively, the dye may react with organic substances in the solution. D. WEIGHING The Analytical Balance Analytical balances vary in capacity and sensitivity, as well as the method of weighing. Sensitivity is defined as the minimum load that will produce a noticeable change in the scale divisions of the readout. Care should be taken to read all instructions careful to ensure accurate weighing. However, in general, the balance should be levelled and free of dirt. The balance should be properly zeroed. The sample to be weighed is placed in the centre of the pan (avoid drafts and air movement). Usually a weighing boat or container is used for chemicals. Most balances have automatic tarring, which can be used to compensate for the weight of the container. When the balance is stable, record the weight of the sample immediately. Electronic single-pan, top-loading balances and automatic tarring are convenient and fast to operate. They usually have a large range and

a precision of 0.5%. E. Water Purification of water to remove impurities is essential in most analytical techniques. Distillation or deionization of water are common methods of purification. However, water can also be purified using carbon filtration, microporous filtration, ultra-filtration, reverse osmosis, ultraviolet oxidation or electrodialysis. Purification will remove unwanted material such as:

Suspended particles - materials suspended in water, making it appear Aunclear@ or

turbid. Colloid particles - either organic or inorganic.

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Dissolved inorganic solids include silicates, chlorides, fluorides, bicarbonates, sulphates, phosphates, nitrates and ferrous compounds which are present in water as either positively charged ions (cations) or negatively charged ions (anions). (Note: This seems to be confusing for some students. Cations are positively charged ion which migrate toward the cathode, the negatively charged electrode. Anions are negatively charged ions which will migrate toward the anode, the positively charged electrode. Dissolved inorganic gases, especially CO2 which can form carbonic acid when combined in water are also removed using exchange resins. Dissolved organics - include proteins, alcohol and detergents. Microorganisms - bacterium, amoebae and algae. Pyrogens and bacterial endotoxins - fragments of cell walls, while viruses are often considered to be non-living nucleic acids. Nucleases (DNases and RNases) - naturally occurring enzymes. Where high purity is required, a combination of the above processes may be used to remove trace contaminants. Efficacy of Purification Processes

Deionization

Distillation

Reverse Osmosis

Dissolved Inorganic Solids

Dissolved inorganic Gases

Dissolved Organics

Particulate

Bacteria

Pyrogens

- Excellent at removing, - Good, - Poor

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Distilled Water Distillation is the most common water purification technology. Water undergoes a phase change during the process that involves boiling the water and then condensing the steam into a clean container. Vaporization results in the separation of water from its contaminants. Pure water (distilled water, dH2O) is produced and the solid contaminants are left behind. However, white or yellowish mineral scale is also left behind which can mean that the apparatus requires frequent cleaning to ensure a high quality of pure water. Where highly purified water is required for biochemical and trace analysis, water may be double distilled (ddH2O).This is used to be the standard prior to the development of commercial systems which combine a combination of methods to purify water. Deionized Water Deionized water, also referred to as demineralized water has had mineral cations such as sodium, calcium, iron, copper and anions such as chloride and bromide removed using a process in which ions in the water are exchanged with a commercially manufactured resin. Ion exchange refers to the exchange of one ion (A+) and replacement with another ion (B+). For example, a solid column or matrix may have associated with it the molecule, X-A+. As a solution containing B+ passes over the column, A+ is replaced by ion B+ which is of similar charge. This exchange is in equilibrium and the degree to which the ion exchange occurs is determined by the relative concentrations of the two ions and their relative affinities for the matrix. Affinity for the matrix is determined by valence. Ca+2 have a greater affinity and will bind tighter than H+. Where the valences are the same, the relative affinity is determined by atomic mass (Li+ typically binds more tightly than H+). X-A+ + B+ <-----> X-B+ + A+ Ions in water can be removed using a two-stage process of ion exchange. First, water is passed through a cation exchange resin. Cations in the water bind to the resin in exchange for H+. The effluent is then passed through an anion exchange resin and the anions in the water bind to the resin and are exchanged for OH-. The exchanged H+ and OH- ions interact to form water. If the water is passed through a series of such columns, the water will become progressively cleaner. Because most of the impurities in water are dissolved salts, deionization can quickly remove these ions, producing high purity water without scale building up in the apparatus. Deionization does not remove uncharged organic molecules, viruses or bacteria. There are specially made strong base anion resins which can remove gram-negative bacteria.

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Reverse Osmosis Reverse osmosis involves pushing water through a semipermeable membrane with pores sizes that are very small (300 Dalton range) using external pressure. Most impurities will not pass through the membrane and collect on the surface of the membrane. Reverse osmosis may be a pre-treatment prior to either distillation or deionizing water. Filtration Either depth or membrane filters may be used prior to other treatments to remove particles and bacteria or as a final step in deionization systems to prevent submicron particles from bacteria from entering the product water. Ultrafiltration is used to remove pyrogens (bacterial endotoxins) from water. Adsorption High surface area activated carbon is used to remove organics and chlorine from feed water. It is therefore a pre-treatment prior to other purification steps. Ultraviolet (UV) Oxidation Photochemical oxidation with ultraviolet light at dual (185 and 254 nm) wavelengths will eliminate trace organics and kill microbes in pure water. It can also be used to prevent bacterial growth. UV light kills microbes by disrupting cell metabolism and reacts with dissolved oxygen to create ozone, promoting hydroxyl radical formation, which oxidize organics. The combination of UV oxidation and ultrafiltration is effective in removing nucleases (DNA, RNase, DNase).

Water quality Anorms@ have been established by a number of professional organizations.

It is important that you know where to obtain the deionized, distilled and ultrapure water. Our ultrapure water purification system uses an ion exchange and filtration technologies as well as a dual wavelength UV lamp and ultrafilters. The system claims to remove nucleases such as RNase and DNase and DNA. This is why you will see this water requested in demanding molecular biology applications such as PCR, electrophoresis and cell culture. However, water for molecular biology experiments often requires further purification to be DNase or RNase-free. Sterile water is obtained by autoclaving.

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Autoclave An autoclave is a pressurized device designed to heat aqueous solutions above their boiling point at normal atmospheric pressure to achieve sterilization. Under standard

pressure, liquid water cannot be heated above 100C because it boils. However, under pressure it is possible to sterilize the liquid by heating the liquid to a much higher temperature because the amount of energy needed to form steam is increased. Most

liquids (<2L) can be sterilized by heating to 121 C for 15 minutes. Sterilization of surgical

supplies may require higher temperatures (121- 135 C) for either shorter or longer periods of time depending on their nature and whether they are wrapped or packed. Plastic containers cannot be used and it is important to remove the trapped air from the liquid as hot air is very poor at achieving sterility. Materials Data Safety Sheets (MSDS) MSDS Data sheets are available through: Safety Office. University of Waterloo. http://www.safetyoffice.uwaterloo.ca/index.html http://www.safetyoffice.uwaterloo.ca/hse/whmis/msda_databases.html Concentration of Common Acids and Bases (Common Commercial Strengths)

Molecular weight

Moles per Litre

Grams per Litre

% by weight

Specific Gravity

mL/L for 1N Solution

Acetic acid, glacial Acetic acid Butyric acid Formic acid Hydrobromic acid Hydrochloric acid Hydrofluoric acid Lactic acid Nitric acid Perchloric acid Phosphoric acid Sulfuric acid Sulfurous acid Ammonium hydroxide Potassium hydroxide Sodium carbonate Sodium hydroxide

60.05 60.05 88.1 46.02 80.92 36.5 20.01 90.1 63.02 100.5 98 98.1 82.1 35.0 56.1 106.0 40.0

17.4 6.27 88.1 23.4 8.89 12.1 32.1 11.3 15.99 100.5 18.1 18.0 0.74 14.8 13.5 1.04 19.1

1,045 376 912 1,080 720 424 642 1,020 1,008 1,172 1,445 1,766 61.2 251 757 1.04 19.1

99.5 36 95 90 48 36 55 85 71 70 85 96 6 28 50 10 50

1.05 1.045 0.96 1.20 1.50 1.18 1.167 1.2 1.42 1.67 1.70 1.84 1.02 0.898 1.52 1.10 1.53

56.9 -- -- -- -- 85.5 -- -- 64.0 -- 22.7 28.4

-- -- B

--

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Format

A: Pipetting and Dispensing Purpose: To familiarize yourself with the various types of pipets available and to learn how each operates. Pipet each of the following volumes using the specified pipet. 1. Pipet the following volumes of distilled H2O: (Repeat until each member of your group is competent in the operation of each

pipet. You should also become familiar with what each of these volumes looks like when it fills the pipet tip.)

a. 9.5 mL using 10 mL glass pipet with propipet

b. 2.0 mL using the pump bottle

c. 1.0 mL using the Eppendorf Repeater

d. 5000 µL using the Eppendorf digital adjustable pipet (air displacement

pipet)

e. 1000 µL using the Eppendorf digital adjustable pipet

f. 100 µL using the Eppendorf digital adjustable pipet

g. 200 µL using the Positive displacement pipet

h. 25 µL using the Positive Displacement pipet

i. 5.0 µL using the Eppendorf digital adjustable pipet 2. Once you have practised with each of the pipets above, pipet dH2O five times,

weighing the volume after each addition using the following pipets:

e. 1000 µL using the Eppendorf digital adjustable pipet f. 100 µL using the Eppendorf digital adjustable pipet

h. 25 µL using the Positive displacement pipet

Record all five trials for each pipet in your lab book. The mass of the solution dispensed should be used to judge your accuracy and precision. You might use a chart such as the one below to assist you in evaluating your accuracy and precision with each pipet.

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Pipet

T1

T2

T3

T4

T5

Mean

S.D.

e. 1000 L Eppendorf

f. 100 L Eppendorf

h. 25 L PDP

B. Preparation of Solutions: Expression of Concentration and Dilution

Purpose: To familiarize yourself with the various ways in which solutions are prepared and the useful means of expressing concentrations of solutions.

1. Prepare the following solutions. Retain these solutions for the pH section of

this laboratory. Be sure to label your beakers.

a. Prepare 100 mL of a 0.89% (w/v) solution of NaCl. Weigh the appropriate amount of NaCl using on a top loading balance. Add this to a graduated cylinder and add the appropriate volume of dH2O to bring the solution to volume (100 mL).

b. Use a graduated cylinder prepare 100 mL of 0.1 M NaOH.

c. Use a graduated cylinder to prepare 100 mL of 0.01 M NaOH using the

0.1M NaOH solution above (#2).

d. Use a graduated cylinder to prepare 100 mL of 0.2 M HCl. Prepare this in the fume hood. Always add the acid to the water.

e. Use a graduated cylinder to prepare 100 mL of 0.02 M HCl using the 0.2 M

HCl solution above (#4).

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C. pH Metre, pH and Buffers

Purpose: To familiarize yourself with the operation of the pH metres. You should know how to calibrate any pH metre. You should also be capable of determining the pH of a solution and know how to adjust it to a required pH.

1. Calibrate the pH metre according to the directions. If your pH metre is only capable

of a two point calibration then calibrate using the pH 7 calibration buffer, followed by the 10. If your pH metre is capable of three point calibration, then calibrate using the pH 4 calibration buffer, followed by the 7 and then 10.

2. While stirring your prepared solutions in a 100 - 150 mL beaker, read the pH of each

solution: b. 0.1 M NaOH. c. 0.01M NaOH.

d. 0.2 M HCl. e. 0.02 M HCl.

3. a. While stirring your 100 mL 0.89% (w/v) NaCl solution in a 100 - 150 mL

read the pH. b. Add 2 drops of your 0.2 M HCl solution using a Pasteur pipet. Observe the

changes in pH as you add these drops and record the final pH value once stable.

. 4. Obtain 100 mL of 1.0 M Glycine.

a. While stirring your 1.0 M Glycine, read the pH of this solution. b. Add 2 drops of your 0.2 M HCl solution using a Pasteur pipet. Observe the

changes in pH as you add these drops and record the final pH value once stable.

5. Adjust the pH of the 1.0 M Glycine solution to 9.95. D. Weighing using a Balance

Purpose: To familiarize yourself with the various balances available for weighing.

You will be required to familiarize yourself with the balance from procedures you have already performed (i.e., Weighing the volumes of water suggested in section A and the required amounts of NaOH and NaCl (section C) using one of the top loading balances.)

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Questions

A: Pipetting and Dispensing 1. Use your recorded observations from pipetting water with the two Eppendorf pipets

and the positive displacement pipets to determine your accuracy and precision. Calculate your accuracy and precision and demonstrate your understanding of the values by commenting on your accuracy and precision. ie, How reliable and valid were you? Include a comment regarding the most likely sources of error.

Pipet T1 T2 T3 T4 T5 Mean S.D.

e. 1000 µL Eppendorf

f. 100 µL Eppendorf

h. 25 µL PDP

2. Complete the table by adding your comments regarding when you would use each

of the pipets in the student biochemistry lab.

Pipet Preferred Use

Transfer Pipet

Brinkman Dispensette

Automatic Pipette

Positive Displacement Pipet

Repeater

3. How you would calibrate one of the air displacement pipets. Also how often should

a pipet be calibrated and when would you calibrate them? 4. Your assay requires that you pipet 100 µL of a standard solution into tubes. You

decide to use an air displacement pipet but which one, the blue Eppendorf or the yellow? Why?

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B. Preparation of Solutions: Expression of Concentration and Dilution 5. a. How many grams of glucose (C6H12O6) would be needed to make 750 mL

of 5 mM solution? b. How many grams of sodium nitrate (NaNO3), would 2 litre of a 0.7 M solution

contain? c. How many grams of Magnesium chloride, MgCl2 are in 50 mL of a 0.25 M

solution of MgCl2?

d. How would you make up 150 mL of 0.75 M CH3COONa⋅3H2O? 6. a. Powerade has 19.7 g of glucose dissolved in 237 mLs. What is the percent

(w/v) concentration of glucose in the solution? b. Blood variables may be expressed in mg or g%. If a client has a

Haemoglobin of 18 g% and a cholesterol of 240 mg%. What does each of these represent?

c. Your American friend goes to the lab and they determine his blood glucose to be 150 mg%. Convert this blood glucose (C6H12O6) value of to a value in mM. Should he be concerned?

d. What is the concentration in g % if 50 g of NaCl dissolved in 2000 mL of solution?

e. Your wine states that the alcohol content is 14 %. How many mLs of ethanol (alcohol) (v/v) are in the 1.5 L bottle?

f. What is the final concentration of a 2.5 M solution, if 100 mL of it is diluted by adding 400 mLs of dH2O?

g. To what final volume would 50.0 mL of a 2.5 M solution be diluted to produce 1.0 M solution?

C. pH Metre, pH and Buffers 7. You are preparing a reagent solutions for your glucose assay, which must be

adjusted to pH = 8.1. How would you calibrate the pH metre, which only accommodates a two point calibration?

8. How often should a pH metre be calibrated? 9. Record the pH readings of the solutions in the table below. Why might the pH of

the solutions differ from the expected values?

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Solution Measured pH Calculated/ Expected pH

0.1 M NaOH

0.01M NaOH

0.2 M HCl

0.02 M HCl

10. Contrast the readings of your NaCl solution and the 1.0 M Glycine solution

provided, before and after you added 2 drops of 0.1 M NaOH. Explain any differences observed.

Solution

Initial pH

2 drops of HCl

0.89% (w/v) NaCl solution

Glycine

11. What is 0.89% (w/v) NaCl and where would you find it naturally?

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References

Freifelder, D.M. Physical biochemistry applications to biochemistry and molecular biochemistry. W.H. Freeman and Company. San Francisco, California. 1976. Robyt, John F. and Bernard J. White. Biochemical Techniques: Theory and Practice. Waveland Press, Inc., Prospect Heights, Illinois. 1987. Stryer, Lubert. Biochemistry. Third Edition. W.H., Freeman and Co., New York, 1988. Voet, Donald and Judith G. Voet. Biochemistry (2nd Edition). John Wiley & Sons, Inc. Toronto, ON, Canada. 1995. Werner, Rudolf. Essentials of Modern Biochemistry: A Comprehensive Review for Medical Students. Jones and Bartlett Pub. Inc., Boston, 1983, Chapter 11. Websites: Safety Office. University of Waterloo. http://www.safetyoffice.uwaterloo.ca/index.html MSDS Link from U of Waterloo Safety Office. http://www.safetyoffice.uwaterloo.ca/hse/whmis/msda_databases.html The MEDLINEplus Health Information. Medical Dictionary. February 2003. http://www.nlm.nih.gov/medlineplus/mplusdictionary.html Access Excellence. Access Excellence at the National Health Museum. The Site for Health & Bioscience Teachers and Learners. http://www.accessexcellence.org/ Chemistry Review Websites: The University of Arizona. The Biology Project: An online, interactive resource for learning biology. http://www.biology.arizona.edu/ MIT Biology Hypertextbook. http://web.mit.edu/esgbio/www/ Chemistry Review Chapter. http://web.mit.edu/esgbio/www/chapters.html Roche Diagnostics Corporation. (2000 http://biochem.roche.com/cfm/country_id_a.cfm/

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Additional Procedures

A: Transferring and Dispensing Volumes - Pipetting Procedures 1. Operating a Transfer, Volumetric or Measuring Pipet

1. Keep the tip of the pipet well below the surface of the liquid and apply suction

gently. Suction can be supplied by using a rubber bulb or propipet. For safety reasons we will pipet all liquids using a rubber bulb or propipet. DO NOT USE YOUR MOUTH TO SUPPLY THE SUCTION.

2. Draw the liquid up the pipet about an inch above the desired volume. 3. Quickly place your index finger over the top of the pipet. 4. Withdraw the pipet from the liquid and allow the liquid to drain out of the pipet until

the bottom of the meniscus coincides with the desired volume. 5. Wipe the outside of the pipet with a Kimwipe to remove liquid which adheres to the

outside of the pipet. 6. Dispense the liquid from the pipet down the sides of the glassware until the bottom

of the meniscus falls on the desired volume. 7. Touch the tip of the pipet to the glassware to remove any residual drop adhering

to the pipet.

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2. Operation of Brinkmann Dispensette (Pump bottle)

1. The dispensette is placed into a bottle containing the liquid to be dispensed and

tightened onto the bottle. 2. Remove the stopper cap if it is in place. 3. Adjust the dispensette to the volume to be dispensed. With the plunger down,

loosen the red button, slide the pointer to the desired volume and tighten the red button.

4. The air must be removed from the discharge tube. Hold a vessel under the

discharge tube and pull the plunger up gently and then press down. Repeat until no more bubbles appear in the cylinder and discharge tubes.

5. To dispense: With the discharge tube pointing away from you, hold your test tube

under the discharge tube. Pull the plunger up gently to the upper stop and then push the plunger down gently and evenly.

accuracy: 0.7%

precision: 0.2%

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3. Operating Air Displacement (Eppendorf) Automatic Micro Pipets

A. Standard Forward Pipetting Standard forward pipetting will generally yields better accuracy and precision than reverse mode for all but viscous or volatile liquids. 1. If the pipet is adjustable then the volume may be selected in one of three ways:

a. For our oldest air displacement pipets, set the volume by turning the plunger knob past the desired volume and then back to the volume to be used. Lock the setting in place by pushing the control button (cap) down until you hear a click.

b. If the pipet has a grey locking button on the left side of the volume window,

then this should be held down while setting the volume by turning the plunger knob until the appropriate volume is displayed. Release the locking button. The set volume is displayed in the window area and is now secured against inadvertent adjustment.

b. The newest pipets have a rotary volume adjustment ring used to select the

volume. The set volume is displayed in the window area and is now secured against

inadvertent readjustment. Never attempt to adjust the volume of the pipet beyond the capacity indicated on the pipet.

2. Apply a clean tip to the instrument. The plunger knob matches the colour of the

required tip. Don’t use your fingers to apply the tip.

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3. Depress the plunger knob to the first stop position before placing the tip into the

sample solution. 4. Now the tip is immersed approximately ≈3 mm into the sample solution. 5. Smoothly return the plunger knob to the ready position allowing the sample to

enter the tip. 6. Remove the pipet from the sample solution. 7. Move the pipette to the receiving vessel. Place the tip against the side of the

receiving vessel. 8. Smoothly depress the plunger knob to the first stop, pause 1 to 3s, then depress

the knob to the second stop to empty the tip completely. 9. Slowly withdraw the tip sliding it against the glass. 10. If, for any reason, an air bubble appears within the tip during the intake, return the

fluid, discard the tip and repeat steps 1 through 8. 11. The tip may be ejected by pressing the plunger knob to the final stop position or

using the releasing knob.

accuracy: 1%

precision: 0.5%

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B. Reverse Pipetting using an Air Displacement Pipet Reverse pipetting is an alternative technique used to dispense a known quantity of liquid when a positive displacement pipet is unavailable. This technique is recommended for solutions with a high viscosity or a tendency to foam. It can help reduces the splashing, foaming or bubble formation and can be more precise in dispensing small volumes of liquids containing proteins and biological solutions compared to forward pipetting, which is the mostly used technique used when dispensing liquids using an air displacement pipet. Reverse pipetting can yield over-delivery.

1. Once the appropriate size tip is attached to the pipette, depress the plunger to the

second stop before immersing it into the sample solution. 2. Dip the tip into the sample solution to a depth of ≈3 mm. 3. Smoothly return the plunger knob to the ready position allowing the sample

solution to enter the tip. 4. Remove the pipet from the sample solution. 5. Move the pipette to the receiving vessel. Place the tip against the side of the

receiving vessel. 6. Dispense the liquid by gently pressing the pipette to the first stop. Withdraw the

tip from the receiving vessel. Some liquid will remain inside the tip. 7. The liquid remaining in the tip can be dispensed back into the original solution or

thrown away. 8. Release the pipette plunger to the ready position.9. If, for any reason, an air

bubble appears within the tip during the intake, return the fluid, discard the tip and repeat steps 1 through 8.

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B. Repetitive Pipetting using an Air Displacement Pipet Repetitive pipetting is useful when the same volume must be added to a number of wells or tubes. It can prevent bubbles or liquid retention in the tip.

1. Once the appropriate size tip is attached to the pipette, depress the plunger to the

second stop before immersing it into the sample solution. 2. Dip the tip into the sample solution to a depth of ≈3 mm. 3. Smoothly return the plunger knob to the ready position allowing the sample to

enter the tip. 4. Remove the pipet from the sample solution. 5. Move the pipette to the receiving vessel. Place the tip against the side of the

receiving vessel. 6. Dispense the liquid by gently pressing the pipette to the first stop. Withdraw the

tip from the receiving vessel. Some liquid will remain inside the tip. 7. Return the pipet to the sample solution and without eliminating the original sample

solution back into the sample container; draw the solution up to the ready position.

8. Move the pipette to the receiving vessel. Place the tip against the side of the

receiving vessel. 9. Dispense the liquid by gently pressing the pipette to the first stop. Withdraw the

tip from the receiving vessel. Some liquid will remain inside the tip. 10. Repeat steps 2 – 9.

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Pre-wetting Pre-wetting is a procedure of aspirating and fully expelling the amount of liquid intended to be pipetted at least three times before actually pipetting the first volume. Pre-wetting will increase the humidity within the tip and may reduce evaporation of the liquid within the tip air space which may cause lower delivery of volumes. Pre-wetting is recommended for volatile liquids. Tips for Accuracy and Reliability Depress the plunger smoothly with a light and consistent force. Immerse the tip into the sample solution at a constant rate and a constant depth. Release plunger slowly to avoid air bubbles. While charging the pipet with the sample solution, hold the pipet vertically and pull the pipette straight out of the centre of the reservoir. Before dispensing examine the tip for and remove any droplets on the outside of the tip. Hold the pipet at a slight angle while dispelling the liquid from the tip into the receiving vessel. Surface tension will help draw the liquid out of the pipet tip, so allow the liquid to run down the sides of the receiving vessel and touch off the tip. Repeatable actions produce repeatable results.

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4. Operating a Positive Displacement Pipet

There are two sizes of positive displacement pipets in this laboratory. The AJ@

size is adjustable for volumes that range between 50 to 250 L, while the AD@ size

is adjustable between the 10.0 and 100.0 L. You will notice that the last digit on the D size pipet is a different colour than the first two digits. The change in colour designates a decimal point. All digits are the same colour on the J size pipet. There is no decimal point on this size.

1. The volume can be set by releasing the locking button by pushing it upward.

Rotate the wheel until the desired volume is reached. Never attempt to adjust the volume of the pipet beyond the capacity indicated on the pipet.

2. Install the glass capillaries by loosening the knurled tip and slipping the glass

capillary on colour marker first. Tighten the knurled tip. 3. To pipet a volume of liquid, depress the plunger fully and place the capillary end in

the sample. 4. Release the plunger slowly to fill the pipet. 5. Wipe the excess liquid from the outside of the capillary. Be careful to avoid

touching the opening. 6. Expel the liquid down the sides of the container by depressing the plunger. Touch

the tip to remove the last drop. 7. If, for any reason, an air bubble appears, repeat the pipetting.

accuracy: 1%

precision: 0.5%

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5. Operating a Eppendorf Repeater

1. Consult the volume selection chart to determine which dial selection and which Combitip to use for the volume you desire.

2. Some Combitips require an adapter. Insert the Combitip gently by sliding the filling

lever down until it stops, raise the locking clamp upwards, insert the Combitip until it clicks into position and close the locking clamp. The electronic repeater will display the volume selected when it has engaged.

3. To fill the Combitip, immerse the tip into the liquid and fill by slowly sliding the filling

lever upwards. Sliding the filling lever too quickly will cause an excessive vacuum resulting in air bubbles. A small bubble in the combitip beneath the piston does not affect the accuracy of pipetting.

4. The first several steps are discarded or an incorrect volume is pipetted. Reject the

first several volumes pipetted. 5. Pipetting is accomplished by placing the Combitip so that it touches the inner side

of the receiving vessel. Depress the pipetting lever down completely. 6. When ejecting the Combitip, the filling lever must be in the lowest position. Hold

the pipet over a waste container. Either unlock the clamping lock (older) or depress both ejection keys at the same time and carefully remove the Combitip.

accuracy 0.3 - 1.6%

precision 0.5 – 2.5%

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Weighing using the Balance - Performing a simple weighing 1. The balance should always be left on. 2. Ensure the balance is level by checking that the air bubble is in the center of the

level indicator. 3. Place the smallest possible weighing vessel (creased weighing paper or

weighing boat) in the middle of the weighing pan. 4. Push the Tare button on the balance and check that the balance displays exactly

zero before weighing the sample. 5. Use a clean scapula to carefully add the required sample to the weighing vessel.

Be aware of possible electrostatic charging of the sample. 6. Use a brush to clean any spills. Discard disposable weighing paper or weighing

boats. Clean spatula before using again.

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Fisher Scientific pH Meter Model 15 Calibration: 1. Press Ameasure/monitor@ to select continuous pH measurement (no key is

displayed). Immerse the electrode in the first buffer (4.00) and observe the display of the current pH. Wait for the reading to become stable (S).

2. Press AStandardization@ -> A2 - Clear Existing Standards@. The metre returns

to the main screen, with all pH standardization points cleared from memory. 3. Press AStandardization@ -> A1 - Update and add a Standard@.

4. If the meter has been set up for manual buffer entry, the Buffer Value screen

appears. Enter the pH value of the buffer. You will be advised to submerge the electrode and to press AEnter@ when you and the electrode are ready press

AEnter@. Before removing the electrode wait until the display has changed to 4.00.

5. Immerse the electrode in the second buffer (7.00) and observe the display and

current pH. Wait for the reading to become stable (S). 6. Press AStandardization@ -> A1 - Update and add a Standard@. If the metre has

been set up for manual buffer entry, the Buffer Value screen appears. Enter the pH value of the buffer. You will be advised to submerge the electrode and to press AEnter@ when you and the electrode are ready press AEnter@. Before removing

the electrode wait until the display has changed to 7.00. 7. Immerse the electrode in the third buffer (10.00) and observe the display of the

current pH. Wait for the reading to become stable (S). 8. Press AStandardization@ -> A1 - Update and add a Standard@. If the meter has

been set up for manual buffer entry, the Buffer Value screen appears. Enter the pH value of the buffer. You will be advised to submerge the electrode and to press AEnter@ when you and the electrode are ready press AEnter@. Before removing

the electrode wait until the display has changed to 10.00. Measuring pH Procedures: 1. Press AMeas/Monitor@ to select continuous monitoring mode.

2. Immerse the pH electrode into the sample solution and provide gently stirring.

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3. When the stability icon appears (S., the metre beeps and the measurement is complete. Repeat steps 2 through 4 for all samples to be measured.

Equip& Tech-45

ORION SA 720 pH METER - MANUAL CALIBRATION WITH TWO BUFFERS

Select either Buffers 4.00 and 7.00 or 7.00 and 10.00, whichever bracket the expected sample range. The ATC probe automatically senses buffer or sample temperature for use in calculating accurate pH values. 1. Rinse the electrode with dH2O and place the electrodes into the first buffer (7.00). 2. Press ACAL@. The cal 1 LED will light indicating that this is the first buffer and a

value has not been entered. Wait for the Aready@ LED to light. Scroll in the correct

value, using and X10 keys, then press Aenter@. The display will freeze for 3 s

and then the Acal 2" LED will light indicating the metre is ready for the second

buffer. 3. Rinse and place electrode into the second buffer. Wait for the Aready@ LED to

light. Scroll in the correct value, using and X10 keys, then press Aenter@.

4. After the second buffer value has been entered the Asample@ LED will

automatically light and the electrode is ready for sample measurements. 5. Remove the electrode from the second pH buffer and rinse with dH2O. 6. Place the electrode in the sample solution to be measured. Insert a stir bead and

stir gently. Wait until the Aready@ LED lights then before recording the pH directly

from the display.