Cell Calcium 35 (2004) 427–431

Some precautions in using chelators to buffer metalsin biological solutions

Chris Patton∗, Stuart Thompson, David EpelHopkins Marine Station, Stanford University, Pacific Grove, CA 93950, USA

Received 17 October 2003; accepted 22 October 2003

Abstract

Chelators and associated computer programs are commonly used to buffer metal ions in biological experiments. This communicationdiscusses common misunderstandings and pitfalls in use of these buffers and provides information on choosing the best metal buffer fordifferent experimental situations.© 2003 Elsevier Ltd. All rights reserved.

Keywords: Calcium; Metal buffer; Chelators; MaxChelator

1. Introduction

Metal ions such as calcium and magnesium have dramaticeffects on cell metabolism and many experimental protocolstherefore require using metal buffers for precise definitionof the metal ion concentrations. Metal chelators are usedfor such buffers and uses include (1) lowering an existinglevel of metal ions inherent in a tissue or cell extract orresulting from “contamination” in reagents or water supply,(2) preventing changes in metal ion concentration or (3)to set specific metal ion levels by buffering the metal ionconcentration to a desired range of concentrations.

Computer programs are normally employed for the aboveuses since direct measurements using ion-specific electrodesare not available in most labs. Additionally, biological so-lutions often have several metal binding molecules in solu-tion (such as specific chelators, amino acids and nucleotidessuch as ATP) and the free metal and metal associations canbe better estimated with these computer programs than withspecific electrodes. These programs indicate the amount ofchelator and metal ion that needs to be present under theparticular conditions of temperature, pH and ionic strengthused in different experimental settings[1–5].1

There are however, limitations and caveats in the use ofmetal chelators and the computer programs associated with

∗ Corresponding author. Tel.:+1-831-655-6216; fax:+1-831-375-0793.E-mail address: cpatton@stanford.edu (C. Patton).1 Note: The computer programs from references[1–3] are all available

at http://www.stanford.edu/∼cpatton/maxc.html.

their use. A search in Biosis (year 2000) for papers usingmetal chelator buffers found that 35% of the solutions usedwere not buffered properly and the resultant metal ion con-centrations were often far different then the assumed levels.

This letter discusses some of the common problems andmisconceptions about metal buffers and reviews parametersthat need to be considered when using chelators as bufferingagents for metal ions. We also present a simple bar chartthat illustrates proper buffer selection.

2. Results and discussion

2.1. Controlling pH of the buffer

Using a pH buffer along with a metal ion chelator is essen-tial to avoid unpredictable effects of pH changes on metal ionconcentrations. This is because chelators are in equilibriumwith hydrogen as well as with metals, as seen inEq. (1).

Chelator–H+ + metal⇔ Chelator–metal+ H+ (1)

Adding metal to a free chelator will liberate H+thereby lowering the pH whereas removing a metal froma chelator–metal complex will cause H+ to be bound, thusraising the pH. A chelator solution might superficially appearto be well buffered in terms of pH since the H+ in the chela-tor is dissociable and can act as a pH buffer. However, when acation is added to the metal buffer, the exchange of metal forH+ (Eq. (1)) can result in dramatic changes in pH that over-whelm the pH buffering capacity of the remaining chelators.

0143-4160/$ – see front matter © 2003 Elsevier Ltd. All rights reserved.doi:10.1016/j.ceca.2003.10.006

428 C. Patton et al. / Cell Calcium 35 (2004) 427–431

Fig. 1. Chelator–metal, good buffer ranges. As calculated using the Max-Chelator program Winmaxc v. 2.40[1]. Downloadable at:http://www.stanford.edu/∼cpatton/maxc.html.

More importantly, the resultant pH change can severely re-duce metal buffering capacity. For example, adding 0.5 mMcalcium to 1 mM EGTA at pH 7 to attain a free calciumconcentration of 0.384�M will drive the pH from 7.0 to 4.6.At this pH (the pH after adding calcium with no pH bufferpresent) the binding of EGTA to calcium would be reduced∼1200-fold. It is therefore critical to properly buffer the pHto prevent pH changes and thus changes in the metal ionconcentration. Ionic strength and temperature changes alsoaffect binding. These effects are minimal compared to pH.

2.2. Chelators cannot bring the metal ionconcentration to zero

A common misconception we found was the idea thatraising the chelator concentration high enough will bring themetal ion concentration to zero. This is not the case, since thechelator is always in equilibrium with the metal (seeEq. (1)above). One can significantly lower a metal concentration,but this will never lead to complete elimination of the metal.

2.3. Selecting the proper chelator for the desiredmetal ion concentration

The titration curves for buffering, be it for H+ or for othermetals, can be represented as a sigmoidal curve. The bestbuffering is in the linear portion of the curve with worse orno buffering at the extremes—bottom or top of the curve.The rule-of-thumb is to be within±0.5pKd, as expressed inlog form, or 0.3Kd–3Kd, as expressed in linear form. Be-ing outside this range means that buffering will be reduced.This relationship is shown graphically in the horizontal bar

Table 1Contamination effects 1 mM EGTA, 0.1N ionic, 20◦C, pH 7.0

1/10Kd Kd 10Kd

Initial calcium added 90.9�M total Ca 500�M total Ca 913�M total CaFree calcium 0.0384�M free Ca 0.384�M free Ca 3.84�M free Ca+10�M Ca “contamination” 0.0431�M free Ca 0.399�M free Ca 4.36�M free CaAmount of change (%) 12.2 3.9 13.5

As calculated using the MaxChelator program Winmaxc v. 2.40[1]. Downloadable at:http://www.stanford.edu/∼cpatton/maxc.html.

graph inFig. 1 which shows the concentrations of calciumor magnesium where different chelators have good bufferingcharacteristics at pH 7.0, 20◦C and 0.1N ionic strength. Forexample, it is seen that EGTA is an excellent buffer for cal-cium in the range of 10−6 to 10−7 but is not good in the rangefrom 10−7 to 10−8. If one wants to buffer at 10−8 range, abetter choice would be EDTA. Similarly, to buffer calcium at10−5 M, HEDTA would be preferable whereas EDTA wouldhave no buffering capacity at this concentration.

2.4. Contamination problems

The consequences of using a buffer out of its range is es-pecially problematic with contaminating levels of cations.For example, consider the situation shown inTable 1where asolution is buffered with 1 mM EGTA and differing levels ofcalcium are added so that the calcium level is now adjustedto 384 nM (in range of EGTA buffering), to 38.4 nM (be-low the range for effective EGTA buffering) or to 3840 nM(above the range of effective EGTA buffering). Now considerthat there is 10�M contaminating calcium in the reagents(a reasonable value allowing for the contamination level inACS grade reagents. For example, the amount of calcium inSigma phosphate-buffered saline solution could be as highas 12�M according to the stated purities). As seen in thelast two rows ofTable 1, the change in free calcium con-centration in the good buffering range is about 4%, whereasit is 12–14% off the desired value when the concentrationis out of range of the buffer. This situation will be exac-erbated when using tissue extracts, which typically containmuch higher concentrations of calcium either bound to pro-teins or in organelles such as the endoplasmic reticulum ormitochondria.

2.4.1. Use of different buffers to circumventcontamination problems

There are several alternatives to the problem of contam-ination. The best solution is to use a buffer that is in thecorrect range. For example, to buffer calcium at 38.4 nM,which is below the range of EGTA (as inTable 1), one canuse a buffer that more closely matches this concentration.Referring toFig. 1, it is seen that EDTA is a better buffer forcalcium in this range.Table 2shows that a 10�M calciumcontamination results in a 4% over-estimate of free calciumusing EDTA as a buffer whereas using EGTA results in a12% over-estimate.

C. Patton et al. / Cell Calcium 35 (2004) 427–431 429

Table 2Contamination effects with different buffers

1 mM EGTA 1 mM EDTA

Kd EGTA–Ca= 0.384�M

Kd EDTA–Ca= 0.052�M

Initial calcium added 90.9�M total Ca 427�M total CaFree calcium 0.0384�M free Ca 0.0384�M free Ca+10�M Ca 0.0431�M free Ca 0.0399�M free CaAmount of change (%) 12.2 3.9

Calculations as inTable 1. Using a buffer with aKd closer to the neededfree level of metal gives less problems with contamination.

Table 3Effects of buffer concentration

10 mM EGTA 10 mM EDTA

Initial calcium added 9.09 mM total Ca 4.27 mM total CaFree calcium 0.0384�M free Ca 0.0384�M free Ca+10�M Ca 0.0388�M free Ca 0.0385�M free CaAmount of change (%) 1.04 0.26

Calculations as inTable 1. Raising the buffer concentration, when belowor at range, can help prevent contamination problems. Compare values toTable 2.

2.4.2. Use of higher buffer concentrations to circumventcontamination problems

An alternative approach, if one wants to use a buffer thatis out of its range, is to increase the concentration of metalbuffer. Table 3 shows the consequences of using 10 mMEGTA or EDTA on attaining a level of 38.4 nM in the faceof 10�M contamination by calcium. As seen inTable 3,the consequences of contamination are minimized with bothbuffers when a higher concentration is used. Of course, thisapproach would not be effective if the desired concentrationis far removed from theKd and this needs to be checkedwith a computer program for metal buffers as describedbelow.

Fig. 2. Effects of reagent purity EGTA purity. Calculated as inFig. 1, 1 mM EGTA, pH 7.0, 20◦C, 0.1N ionic strength. Below is 0.1Kd, and above is10Kd. Being above the effective range of a chelator increases the likelihood of error.

2.5. Problems with chelator and standard purity

A major problem in accurate specification of metal con-tent is the purity of the buffer or the metal standard. Thisseems obvious, but a brief recapitulation of the situation andproblems will illustrate the importance of adjusting solu-tions to the known purity. For example, consider that EGTAtypically is sold at 97% purity. If one makes up the solu-tion treating the EGTA as 100% pure, the expected calciumlevels for 38.4, 384 and 3840 nM would respectively be offby +7% below the good buffering range,+4% at the goodbuffering range and a whopping 47% above the good buffer-ing range, as illustrated inFig. 2A.

A related concern is the purity of the calcium used forbuffering. This is a problem since most calcium salts aredeliquescent, making the actual concentration uncertain dueto variable water uptake.Fig. 2B gives an example of thisproblem. This figure depicts a situation similar to that inTables 1–3where calcium is added to achieve different con-centrations of free calcium. Assuming a pure reagent, thedesired free calcium level will be achieved. However, if thereagent has absorbed water so that the level is 90% of thepredicted value, the actual free calcium levels are of coursenow different. The important point here is that the effects ofunknown concentration are seen whether the buffer is in orout of range; however, the consequences of being above thebuffer range are considerably exaggerated.

The best solution to this problem is to use standardizedcalcium solutions such as supplied by Fisher and other com-panies which provide standard solutions dissolved in HCl.The alternatives may not be satisfactory. For example, heat-ing calcium chloride to drive off water is problematic sincethere can be production of calcium oxide if the temperatureor time of heating is not controlled properly. Use of non-hydrated salts such as calcium acetate or calcium nitrate isproblematic since the anions can have physiological effects.

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Table 4Chelator competition

Condition Free magnesium(86.9�M)

ATP–Mg(413�M)

+1 mM EDTA (�M) 3.73 29

Calculations as inTable 1, pH 7.0, 20◦C, 0.1N ionic strength, 1 mM ATP,0.5 mM Mg. Mg is well buffered in both cases, but with the additionof EDTA, the EDTA has removed most of the Mg from the ATP–Mgcomplex and is now buffering the Mg at a much lower range.

There are three lessons from the above examples. First ofcourse is the need to account for the actual purity of reagents.Secondly, the examples depict how the buffering capacity ofthe chelators can get around impurity problems in the areasof the buffer range that are below or in the good bufferingrange. The third lesson is that the effects of impurities (orcontamination) are more exacerbated in the area above thegood buffer range.

2.6. Competition among chelators and metals canlead to unexpected effects

Changing chelators or metal concentrations can affectother metals or other binding molecules in the mixture. Forexample, Mg2+ and ATP are in dynamic equilibrium andindeed the active form of ATP in cells is the Mg2+–ATPsalt. Adding an additional chelator to mixtures of Mg2+ andATP can have profound effects on the Mg2+–ATP equilib-rium. An example is shown inTable 4which depicts thedual effects of EDTA in reducing free Mg2+ but also reduc-ing the amount of Mg2+–ATP. In this case, the EDTA haschelated Mg2+ previously present as free Mg2+ as well asMg2+ bound to ATP. The problem here then is the chelatorcompetition between ATP and EDTA. This type of compe-tition will take place in any mixture in which there is morethan one chelator.

Metals can also compete with each other for the chela-tor. For example, adding Ca2+ to a mixture of EDTA andMg2+ will lead to the “bumping” off of the Mg2+ and anincrease in the free Mg2+ concentration, as well as reduc-tion in the amount of free EDTA. An example is depicted inTable 5which shows the consequences of adding calciumto a Mg2+–EDTA solution. The Ca2+ binds the remainingEDTA and also displaces Mg2+ from EDTA, so that thereis a 10-fold increase in free Mg2+ as well as a diminutionof available EDTA.

Table 5Metal competition

Condition Free magnesium(4.19�M)

Free EDTA(504�M)

+0.5 mM calcium (�M) 44 44

Calculations as inTable 1, pH 7.0, 20◦C, 0.1N ionic strength, 1 mMEDTA, 0.5 mM Mg. Ca has effectively “bumped” most of the Mg fromthe EDTA–Mg complex and eliminated EDTA–Mg buffering capacity.After the addition of the calcium neither metal is well buffered.

2.7. Using the right dissociation constants

The computer programs available in the literature or on theInternet use similar algorithms for determining metal con-centrations with specific buffers under defined conditions ofionic strength, temperature and pH. The major variable inthese programs is choosing the metal dissociation constants.Most programs provide these constants internally but theseconstants are often updated and/or revised and it can be diffi-cult to change the programs constants. Some programs, suchas MaxChelator, attempt to keep up with these revisions andimportantly MaxChelator has editable files so that changesin the constants can be easily entered. New chelators or theuse of biological molecules that have intrinsic metal buffer-ing activity also have to be accounted for by using the properdissociation constants. Constants, where they are available,may be obtained from NIST[7] or the Martell and Smithvolumes[6].

2.8. Solutions to the problems brought up in this letter

No computer program can ever replace a properly cali-brated metal-specific electrode, but they can serve very wellwhere an electrode is impractical or not available. These pro-grams are particularly useful in biological situations wherethere are multiple cations and multiple molecules capable ofchelating these metals. Taking care of the factors discussedabove when using these programs can insure that the cor-rect level of metal has been achieved in the experiment. Insummary, it is important to:

1. Insure that the pH is properly buffered with a pH buffer(that is also within its effective range,±0.5 pH units ofpKa).

2. Choose the proper buffer/chelator and buffer/chelatorconcentration.

3. Be aware of the problem of contamination and how thisproblem can be diminished by using the right chelatorand chelator concentration and using standard solutionsof the buffered metal.

4. Be aware of the problem of competition between metaland chelators in the particular solution that is being used.

5. Use the proper dissociation constants for the species ofchelators that are present in the mixture.

Acknowledgements

Supported in part by NIH grant 1 RO1 HD 38488, andNS 43462.

References

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