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Buffer Calculator

A buffer recipe is two numbers, not one: how much of the species you weighed has to be converted, and what you add to convert it. Henderson–Hasselbalch gives you the ratio inside a single bracket and leaves the second number to your bottle of HCl; this calculator solves the full charge-and-mass balance — the buffer’s own ionisation plus water’s H⁺ and OH⁻ plus the strong ions the salts drag in — and prints the Henderson–Hasselbalch answer next to it so you can see the size of the shortcut. Choose a system and a starting form (glacial acetic acid, sodium acetate, citric acid monohydrate, Na₂HPO₄·7H₂O, Tris base, HEPES free acid…), give the target pH, the concentration, the final volume, the use temperature and the molarity of the titrant you actually pipette, and it returns the grams to weigh, the millimoles and millilitres of NaOH or HCl that reach the target, the co-ion load the recipe leaves behind, the buffer capacity at that pH and how much acid or base the design can really absorb before it moves. Four modes cover the bench questions in both directions: design a recipe, read back the pH of a pair you already weighed, get capacity and reserve, or ask which of ten systems with cited pKa values buffers your target pH at all. Every constant is the value the cited reference page publishes at 25 °C, and where a page publishes no temperature coefficient the calculator says so instead of inventing one.

Making a buffer looks like a one-line calculation until the meter disagrees with you. The line is Henderson–Hasselbalch — pH equals pKa plus the log of the base-to-acid ratio — and it is a good line: it tells you which ratio to aim for and why the useful window is about one pH unit wide. What it does not tell you is how much strong acid or base that ratio costs once water’s own H⁺ and OH⁻ are counted, how the answer changes when a second ionisation step sits nearby, or what ions you are adding to the beaker while you are adding the pH. Those are the parts that decide whether 10 mM phosphate actually reaches pH 12.3, whether a citrate buffer at pH 4.0 behaves like citrate at all, and whether your Tris reading at 37 °C is a measurement error or just temperature.

This page solves the whole balance instead of the ratio. For a target pH it computes the fraction of the buffer in every ladder form (acid, mono-, di- and tri-deprotonated as the system has them), the mean number of protons the population has given up, and therefore the net strong ion — Na⁺, K⁺ or the Cl⁻ of HCl — the solution has to carry to sit there. The titrant amount is the difference between that requirement and what the form you weighed already supplies. The result is a recipe: grams of the bottle, millimoles of NaOH or HCl, the millilitres that is at the normality you pipette, and the co-ion concentration the finished buffer contains whether or not you wanted it.

The Henderson–Hasselbalch number is printed beside it rather than replaced by it, because the gap is the information. For 50 mM acetate at its pKa the two agree to about a tenth of a percent — the free H⁺ the acid has already released is a fraction of a millimole, and a single bracket holds the design. For 100 mM citrate at pH 4.00 they differ by roughly ten percent of the titrant, because pKa 3.13 and pKa 4.76 are both within a unit of the target and the equation only knows one of them. For 10 mM phosphate pushed to pH 12.3 there is no Henderson–Hasselbalch answer at all: the strong ion needed exceeds what the buffer can carry even fully deprotonated, and the excess is free OH⁻, which the equation has no term for.

Capacity is the same machinery read a different way. β = dn/dpH comes out of the variance of the ladder plus a water term, and the page reports both parts and the share of each — which is what makes dilution legible rather than mysterious. A 100 mM acetate buffer at its pKa absorbs about 41 mmol of acid or base per litre to move a whole pH unit; the same buffer at 1 mM absorbs a hundredth of that and spends a visible part of what remains on its own H⁺ and OH⁻. Nothing in that sentence needs a separate buffer-capacity page, so it is on this one, together with the recipe it belongs to. The guide that walks the same design end to end is how to make a buffer solution, and the arithmetic that turns a bottle’s assay percentage into the titrant volume you pipette lives in the molarity calculator.

Each system carries only the pKa values its reference page publishes.

Hydrates and salts differ in mass and in which titrant you need.

The pH you want the finished solution to sit at, at your use temperature.

Total analytical concentration, counting every ionisation form.

The volume you make up to after the titrant is in.

Sets water’s pKw, and the pKa where the source publishes a coefficient.

The NaOH or HCl you will actually pipette.

pH mode — what you weighed of the more protonated form.

pH mode — the more deprotonated form in the same beaker.

pH mode — leave blank if you weighed only one form.

Capacity mode — the reserve is integrated across this whole step.

NaOH to add24.98 mmol

24.98 mmol NaOH = 24.98 mL of 1 M after dissolving 3002.6 mg of Acetic acid (glacial)

pKa used4.756 pH (25 °C table value)
Distance from pKa0 pH unit
Henderson–Hasselbalch4.755 vs 4.756
Strong-ion gap24.98 mM
Capacity (β)28.82 mmol·L⁻¹·pH⁻¹
Water’s share of β0.14 %
Na⁺ / K⁺ in buffer24.98 mM
Cl⁻ in buffer0 mM

How fast capacity falls as you leave the pKa

For a single ionisation step the capacity is proportional to var(j) = r/(1+r)², with r = 10^(pH − pKa). Dividing by its value at r = 1 gives the share of the peak you still have — pure arithmetic, no constants needed.

|pH − pKa|[A⁻]/[HA]Share of peak capacityWhat it means at the bench
0.01 : 1100%The design point: half and half, maximum resistance in both directions
0.32 : 189%Barely any loss; this is why the window is quoted generously
0.53.2 : 173%Still a buffer; asymmetry starts to show (one direction has less reserve)
1.010 : 133%The edge of the window — a third of the capacity, and only one direction is protected
1.532 : 112%Nominal buffering; the water term is now competitive
2.0100 : 14%Dilute acid or base with a label on it — change system, not concentration

The peak itself is set by concentration alone: ln10·C/4, so any one-step buffer at 100 mM tops out near 57.6 mmol·L⁻¹·pH⁻¹ at its own pKa, whether it is acetate, MES, MOPS, HEPES or Tris. The system decides where the peak is; the concentration decides how high it is; and because capacity is a slope that falls as you move, the reserve over a whole pH unit is always smaller than β × 1 — which is why this page reports both.

How to use the buffer calculator

Pick the mode for the question you have: Recipe (what to weigh and titrate), pH of a pair I already weighed out, Buffer capacity, or Which system for my target pH.

Choose the buffer system. The label carries the pKa values the calculation uses, and they are the values the cited reference page publishes — no fitted or rounded second table.

Choose the starting form — the thing on your balance. Glacial acetic acid and sodium acetate sit at opposite ends of the same ladder, so they need opposite titrants for the identical finished buffer; hydrates change the mass and nothing else.

Enter the target pH, the buffer concentration (total analytical, all forms), the final volume and the use temperature, each with its own unit selector.

Enter the titrant concentration you actually have (1 M NaOH, 5 M HCl, whatever the bottle says) to get millilitres as well as millimoles; if you are starting from a percentage assay, convert the bottle’s mass and purity to molarity first rather than eyeballing the volume.

For the pH mode, weigh what you weighed: enter the mass of the more protonated form and, if you used both, the second form and its mass. The tool reports the pH the mixture actually settles at and the Henderson–Hasselbalch reading of the same beaker.

For capacity, set the pH step you want to survive (0.5 or 1 unit are the bench numbers); the result gives β, its buffer and water components, and the integrated mmol/L that survives that step.

Read the warnings as part of the output: the ±1 window, a second pKa contributing, a water-dominated dilution and a missing temperature coefficient each change what the number means.

Make a concentrated version once and draw it down later with the stock solution calculator and the dilution calculator; keep the co-ion load in view if the buffer touches cells, enzymes or metal ions.

Buffer formula

Charge-and-mass balance: S = Kw/[H⁺] − [H⁺] + C·(⟨j⟩ − z₀)
Ladder speciation: α_j ∝ Π (Ka_i/[H⁺]) for j = 0…n, ⟨j⟩ = Σ j·α_j
Titrant to add: n = V · ( S(target pH) − C·(j_form − z₀) ) NaOH if positive, HCl if negative
At the pKa: ⟨j⟩ = i + ½, so a neutral acid needs half its moles neutralised
Henderson–Hasselbalch: pH = pKa + log₁₀([A⁻]/[HA]) (one bracket, no water term, no second step)
Buffer capacity: β = dn/dpH = ln10·C·var(j) + ln10·([H⁺] + Kw/[H⁺])
Integrated reserve: Δn = | S(pH ± Δ) − S(pH) | over the whole step, not β × Δ
Co-ion load: [Na⁺] from salts + NaOH, [Cl⁻] from HCl and any chloride salt
Water at temperature: pKw interpolated between the tabulated anchors (0, 25, 50, 75, 100 °C)
Concentrations, volumes and masses are converted to canonical units (mol/L, L, g) through the site’s shared units table before the balance runs, so mM with mL and µM with µL meet inside one equation. The Henderson–Hasselbalch reading is derived from the same speciation, then compared; where its single-bracket assumption cannot represent the mixture, it reports no reading instead of a number.

Variables and units

SymbolMeaningUnit
SStrong-ion gap — net (Na⁺ + K⁺) minus Cl⁻ the solution must carry to sit at the target pHmol/L (reported as mM)
CTotal analytical buffer concentration, counting every ladder formM / mM / µM / nM
z₀Charge of the fully protonated reference form (0 for plain acids, +1 for Tris, HEPES, glycine, ammonium)dimensionless
⟨j⟩Mean number of ladder protons the population has given up at this pHdimensionless (0…number of pKa steps)
α_jFraction of buffer in ladder form j, from the pKa ladder at the use temperaturedimensionless
pKaThe ladder step nearest the target pH; the ±1 window is measured from itpH unit
KwIonic product of water at the use temperature, from the tabulated pKw anchorsM²
βBuffer capacity, the moles of strong acid or base needed per litre per pH unitmol·L⁻¹·pH⁻¹ (shown as mmol·L⁻¹·pH⁻¹)
ΔnIntegrated reserve across the pH step you asked formmol/L
n(OH⁻/H⁺)Titrant the recipe asks for, given the form you weighedmmol (and mL at your titrant normality)

How buffer is calculated

  1. Convert every input to canonical units: concentration to mol/L, volume to L, mass to g, titrant normality to mol/L.
  2. Take the system’s pKa ladder at your use temperature: water’s pKw comes from the tabulated anchors by linear interpolation between them, and a pKa is moved only where the source page publishes a coefficient (Tris, 25→37 °C).
  3. Select the ladder step nearest the target pH; that step is the conjugate pair Henderson–Hasselbalch would use, and the distance to it is what the ±1 window is measured against.
  4. Compute the speciation fractions α_j from the ladder and [H⁺], then the mean dissociation ⟨j⟩ the population actually has.
  5. Evaluate the strong-ion gap S = Kw/[H⁺] − [H⁺] + C(⟨j⟩ − z₀) at the target, and subtract what the form you weighed already supplies; the sign of the difference picks NaOH or HCl and the magnitude is the millimoles.
  6. Fold the millimoles into millilitres at the titrant normality you entered, and report the sodium, potassium and chloride the recipe leaves in solution.
  7. Run the same speciation backward for the Henderson–Hasselbalch comparison, then report β and the integrated reserve over your chosen pH step.
  8. Check the design against the model’s own limits before printing: distance from pKa, a second contributing step, water’s share of capacity, and whether a temperature coefficient exists at all.

Worked laboratory examples

Example 1 — 50 mM sodium acetate, pH 4.756, one litre

Given: Target pH 4.756 (= acetic acid’s pKa) · C = 50 mM · V = 1 L · 25 °C · Titrant: 1 M NaOH · Starting form: glacial acetic acid

  1. Weigh the acid for the total: 50 mmol × 60.052 g/mol = 3.003 g of glacial acetic acid, dissolved in about 800 mL.
  2. At pH 4.756 exactly half the ladder is deprotonated, so the balance wants 25 mmol of Na⁺ — and one term fewer: 1.75 × 10⁻⁵ M of that H⁺ is already free in solution at this pH, so the strong base needed is 24.98 mmol, not 25.00. That is 24.98 mL of 1 M NaOH, then make up to the mark.
  3. A Henderson–Hasselbalch-only tool asks for 25.00 mmol of the same NaOH. The 0.0175 mmol difference is 0.1% here — which is the honest headline: at the pKa, at 50 mM, the shortcut is safe, and the page shows you the arithmetic that proves it.
  4. Start from sodium acetate instead and the same finished buffer costs 4.102 g (50 mmol × 82.034 g/mol) plus 25.02 mmol of HCl — the titrant flips direction, and the co-ions change: the first route leaves 24.98 mM Na⁺, the second leaves 50 mM Na⁺ and 25.02 mM Cl⁻.

Result: 3.003 g acetic acid + 24.98 mmol (24.98 mL of 1 M) NaOH, made up to 1 L; HH comparison 25.00 mmol (0.1% apart)

Example 2 — citrate at pH 4.00, where one bracket is not enough

Given: Target pH 4.00 · C = 100 mM · V = 1 L · 25 °C · Titrant: 1 M NaOH · Starting form: citric acid, anhydrous

  1. Citric acid publishes three steps: 3.13, 4.76 and 6.39. Two of them sit within a pH unit of 4.00, so both carry charge at the target and the page says so instead of picking one.
  2. The exact balance puts the population at a mean dissociation of 1.03 — a little past the first step — so 100 mmol of citric acid (19.212 g) needs 102.98 mmol of NaOH: 102.98 mL of 1 M, which is 10% of your final volume, so dissolve and titrate in correspondingly less water.
  3. Henderson–Hasselbalch, working only inside the 3.13/4.76 bracket, asks for 114.81 mmol of the same NaOH to reach pH 4.00 — 10.3% more base. Feed the amount the exact model actually prescribed back into the single-bracket equation and it reads the mixture as pH 3.25, three quarters of a unit low.
  4. Neither number is a rounding artefact: the shortfall is the second ionisation the equation cannot see. A pKa table that lists citric acid as 2.79 or phosphoric acid as 2.15 only is quoting one step of a ladder and calling it the constant.

Result: 19.212 g citric acid + 102.98 mmol NaOH; HH asks 114.81 mmol (10.3% apart) and reads the designed mixture as pH 3.25

Example 3 — Tris at 37 °C: the correction that exists, and the ones that do not

Given: Target pH 7.50 at 37 °C · C = 50 mM · V = 1 L · Titrant: 1 M HCl · Starting form: Tris base

  1. Tris buffers through its conjugate acid (pKa 8.07 at 25 °C, the powder sitting one rung up the ladder), and the reference page publishes a temperature coefficient: about 0.025 pH unit down per degree between 25 and 37 °C. At 37 °C the working pKa is therefore 7.77, not 8.07.
  2. pH 7.50 is then 0.27 unit below that pKa, so a little under a third of the amine is protonated and the rest must be protonated by strong acid: 6.057 g of Tris base (50 mmol × 121.136 g/mol) plus 32.53 mmol of HCl — 32.53 mL of 1 M, leaving 32.53 mM Cl⁻ in the buffer.
  3. The practical consequence, straight from the coefficient: a Tris buffer designed to read 7.5 at 25 °C sits near 7.2 when warmed to 37 °C. That is not electrode drift, and adding acid after warming only moves the target.
  4. For every other system in the table the reference page publishes no coefficient, so the calculator holds the 25 °C value and warns you it did — it does not borrow another system’s slope or fit an invented one. Water, though, is corrected for everyone: pKw interpolates from 13.99 at 25 °C to 13.64 at 37 °C, which is part of the balance at low concentration.

Result: 6.057 g Tris base + 32.53 mmol (32.53 mL of 1 M) HCl, designed against pKa 7.77 at 37 °C

Example 4 — capacity, and what dilution actually costs

Given: Acetate at its pKa 4.756 · Three designs: 100 mM, 20 mM and 1 mM · pH step to absorb: 1 unit

  1. At the pKa a single-step buffer reaches its maximum capacity, β = ln10·C/4 ≈ 0.576 mmol·L⁻¹·pH⁻¹ per mM — so 100 mM acetate gives 57.6, and the same is true of MES, MOPS, HEPES or Tris at their own pKa values. Which buffer you chose barely moves the ceiling; how close you are to it, and how concentrated you are, does.
  2. Capacity is a slope, and it falls as soon as you move, so the honest number is the integrated reserve: pushing 100 mM acetate a whole pH unit needs about 41 mmol per litre rather than the 57.6 that β × 1 would suggest.
  3. At 1 mM the reserve is 0.42 mmol/L going up, and water’s own H⁺/OH⁻ carry 6.6% of the capacity — the point where the Henderson–Hasselbalch shortcut stops being arithmetic and starts being fiction: the same mixture reads 0.03 unit low, and in a real hood the dissolved CO₂ and the rinse water decide the pH long before your ratio does.
  4. Runaway pH is its own failure mode: 100 mM acetate asked to sit at pH 3.0 is 1.76 units from its pKa, the water term takes 37% of the capacity, and the Henderson–Hasselbalch titrant comes out 58% away from the balanced one. The page flags the window, the water share and the mismatch together, because they are the same story.

Result: β 57.6 mmol·L⁻¹·pH⁻¹ at 100 mM (reserve ≈ 41 mmol/L per unit); at 1 mM β 0.62 with 6.6% spent on water

Want the reasoning behind the numbers? Read How to Make a Buffer Solution.

How to interpret the result

The primary card answers the question you asked, and the cards beneath it keep the design consistent: the pKa actually used (with its temperature status named), the distance from your target to it, the strong-ion gap the solution must carry, the co-ion concentrations that come with it, and capacity with its buffer and water components. For a recipe the pair to read together is the titrant and the co-ion load — 25 mM Na⁺ and 25 mM Cl⁻ are not incidental to a cells-and-enzymes experiment, they are the reason the same buffer made two different ways behaves two different ways.

The Henderson–Hasselbalch row is the diagnostic rather than the answer. When it agrees to a tenth of a percent, you are in the regime where the shortcut was always going to work and you can prepare the buffer by ratio alone. When it disagrees by ten percent, the gap is a second ionisation step, and you should read the exact recipe as the one to follow. When it prints no reading, the mixture’s charge cannot be attributed to any single conjugate pair at all — the excess is free H⁺ or OH⁻, and the equation you searched for has no vocabulary for that situation.

The warnings each change the meaning of the numbers above them. A target more than one pH unit from every pKa of the chosen system is not a buffer, and concentration cannot fix it — the recommender mode exists so you can switch systems instead of adding salt. A second pKa within the window means the single-ratio design your protocol quotes is describing a different solution. A water share above about 1% of capacity means the calculation is no longer the dominant term in the beaker, and dissolved CO₂, the rinse water and the carry-over from the previous tube are. And a missing temperature coefficient means the design is stated at 25 °C even though you are working at 4 °C or 37 °C — the page will tell you which of those three cases it is in, in words, under the result.

Capacity and reserve are the quantities behind two common bench decisions. How concentrated should the buffer be: enough that the sample and the reaction’s own acid or base production are small against the reserve, which you can now read in mmol per litre per pH unit rather than guess. And can I dilute this 10× stock straight to 1×: the concentration enters the balance linearly, the water term does not dilute with it, and the share column is exactly the difference those two facts make. If you also need the pH at which a protein in the buffer stops being soluble, that is a different quantity and the protein isoelectric point calculator is where it is computed.

Assumptions

  • Ideal, concentration-based equilibria: activities are not corrected for ionic strength, so a calibrated meter is the release criterion and this page predicts roughly where it will land.
  • A closed system: no CO₂ exchange with the air. That is a fair assumption for a capped tube of HEPES at 50 mM, and a wrong one for an open beaker of dilute bicarbonate, which is why the carbonate system is not offered here.
  • The pKa ladder is the value the cited reference page tabulates, at the temperature that page states; forms not in the table (Tris·HCl, sodium citrate dihydrate, glycine·HCl) were left out rather than guessed at.
  • Volume is final volume, added after the titrant: the recipe’s millilitres of NaOH or HCl are counted against the volume you make up to, not added to it.
  • Temperature affects the model through water’s pKw for every system, and through the pKa only where a published coefficient exists.

Limitations

  • It formulates from published constants; it does not fit a pKa or a capacity from your own titration curve, and it cannot see the impurities, the aged stock or the electrode that decide your reading.
  • No activity or Debye–Hückel correction, so in a 100 mM salt background expect the meter to sit a few hundredths to a tenth of a unit away from the prediction — real, but not this model’s error to remove.
  • Ten systems, and only the pH region their cited pKa values cover: roughly 2.1 to 12.3 across the table, with genuinely useful buffering confined to one unit around each.
  • Temperature correction is honest but sparse: Tris is the only system here whose reference page publishes a coefficient, so a 4 °C phosphate or 37 °C HEPES design is stated at its 25 °C value and flagged.
  • Mixed and non-aqueous systems, protein buffers, liposomes and gradients are outside the model; so is any titrant other than the strong acid or base you name.

Common mistakes

  • Treating a polyprotic acid as one pKa: citrate at pH 4.0 and phosphate near pH 2 or 12 have two steps contributing, and a single-ratio calculator will still hand you a confident answer.
  • Choosing a buffer by the pH you want rather than the pKa you have — pH 3.0 with acetate is not a weak buffer, it is dilute acid with a label.
  • Adding acid or base to a Tris buffer at room temperature and expecting pH 7.5 at 37 °C: the coefficient is about 0.025 unit per degree, and the design temperature is the one that must match.
  • Quoting total protein or a nominal bottle concentration as the active species, then wondering why the recipe is off; and reading hydrate forms as if the water were not on the balance — Na₂HPO₄ anhydrous, dihydrate and heptahydrate are three different masses for the same ladder position.
  • Diluting a 10× stock and assuming the pH is unchanged: the ratio survives dilution, the water term does not, and below a few millimolar the meter is measuring the room, not the buffer.
  • Forgetting the co-ion: a phosphate–NaOH recipe and a phosphate–KOH recipe are the same pH and a different experiment for cells, enzymes and metal-dependent assays.
  • Adjusting pH after making up to final volume, so the titrant you added is not in the concentration you recorded.
  • Trusting a pKa table that gives phosphoric acid 2.15 and citric acid 2.79 as if each were a single number — those are one rung of a ladder, presented without the rest of it.

FAQ

How do I calculate how to prepare a buffer?

Four steps: pick the system whose pKa sits within about one pH unit of your target; decide the total concentration and final volume; compute how many moles of the species you weighed must be converted to reach the target protonation, which is the strong-ion requirement of the charge balance minus what that form already supplies; and add that as strong base or strong acid, then make up to volume. The first and last steps are where a Henderson–Hasselbalch-only tool stops, and the four modes on this page do them in the open: they print the grams, the millimoles, the millilitres at your titrant normality and the ions the recipe leaves behind.

How do I make 50 mM sodium acetate buffer, pH 4.756?

Dissolve 3.003 g of glacial acetic acid (50 mmol) in about 800 mL of water, add 24.98 mmol of NaOH — 24.98 mL of 1 M — and make up to 1.0 L. At the pKa the ladder is half-converted, so the intuitive answer is 25 mmol; the exact balance asks for a fraction of a millimole less because 1.75 × 10⁻⁵ M of the protons are already free in solution. Start from 4.102 g of anhydrous sodium acetate instead and the same buffer takes 25.02 mmol of HCl. Both routes are shown by the recipe mode, with the co-ion load each one leaves in the flask.

What is the Henderson–Hasselbalch equation, and when does it fail?

pH = pKa + log₁₀([A⁻]/[HA]) — it converts a target pH into the ratio of a conjugate pair, and it is the right tool inside its bracket. It fails in three specific ways: when a second ionisation step is within about a unit of the target (citrate near pH 4, phosphate near pH 2 or 12), when the buffer is dilute enough that water’s own H⁺ and OH⁻ are a visible share of the charge, and when the required strong ion exceeds what the buffer can carry at all. The calculator reports its answer beside the exact one, and says plainly when there is no Henderson–Hasselbalch reading available.

Why should the pH be within one unit of the pKa?

Because capacity — the thing a buffer is for — is proportional to the variance of the ionisation states, and a single step reaches its maximum when the pair is half and half, at pH = pKa. One unit away the ratio is 10:1 and capacity has fallen to roughly a tenth of the peak; two units away there is no useful resistance left. Concretely, 100 mM acetate at its pKa resists about 41 mmol of acid or base per litre over a whole pH unit, while the same 100 mM asked to sit at pH 3.0 resists far less and spends most of what it has on free H⁺. The distance warning on this page is that arithmetic, not a preference.

How much NaOH or HCl do I add to reach the target pH?

The difference between the strong-ion gap the finished buffer must carry and the gap the weighed form supplies — which is what the recipe mode computes, in millimoles and in millilitres at the normality you enter. Add that as accurately as you can, in a beaker at your use temperature, before making up to final volume; then check the meter, because activity effects are real and the calculation is a concentration model. If you need to convert the bottle’s assay percentage or the mass you weighed into a molarity first, the molarity and stock-solution arithmetic is on its own pages rather than repeated here.

What is buffer capacity and how is it different from just being concentrated?

Buffer capacity is β = dn/dpH, the moles of strong acid or base one litre absorbs per unit of pH change. It has two parts: ln10·C·var(j) from the buffer, which scales with total concentration and collapses as you move away from the pKa, and ln10·([H⁺] + Kw/[H⁺]) from water, which does not care about your buffer at all. That is why 100 mM acetate at its pKa gives β 57.6 mmol·L⁻¹·pH⁻¹, 1 mM gives 0.62, and a dilute design near pH 3 or 12 can be mostly water. Capacity is a readout of this recipe page rather than a separate tool, so it is stated against the same grams you are weighing.

Which buffer should I use at pH 7.4?

At 25 °C, among the systems with cited values here: HEPES is closest (pKa 7.5, one tenth of a unit away), then phosphate and MOPS (7.20, 0.2 away), then PIPES (6.76) and Tris (8.07) at about two thirds of a unit. At 50 mM the ranking also prints β — HEPES 28.4, phosphate and MOPS 27.3, PIPES 17.5, Tris 16.7 mmol·L⁻¹·pH⁻¹ — because distance decides whether a buffer works and capacity decides how many grams it takes. Switch to the recommender mode for any target pH; the same ranking is worked through bench-first in the guide linked above.

Does temperature change a buffer’s pH, and can I calculate it?

Yes, and partially. Two effects always exist: water’s pKw moves from 14.95 at 0 °C to 13.99 at 25 °C to 12.25 at 100 °C, which the calculator interpolates from the tabulated anchors and applies to every system; and each pKa moves by its own thermodynamic amount, which is only calculable where a reference publishes the coefficient. For Tris the published slope is about 0.025 unit per degree between 25 and 37 °C, so a 7.5 design at 25 °C reads near 7.2 at 37 °C. For the rest of the table this page holds the 25 °C value and tells you it did — the alternative is a number nobody can source, which is worse than a warning.

Scientific Sources

External links open in a new tab and carry rel="nofollow". We cite primary references (NIST, NIH/NCBI, peer-reviewed papers) and established manufacturer technical documentation (NEB, Thermo Fisher, Promega, IDT). Calculated results are estimates; your protocol and manufacturer guidance take precedence.

Last reviewed: 2026-10 · Method version: 1.0