Effects of Buffers on pH: Lab Explained

By Dr. Zubair Khalid, DVM, MS, PhD · 2026-09-20

A buffer is a solution that resists a change in pH when you add a strong acid or a strong base. By the end of this article you will be able to prepare a phosphate or acetate buffer, predict its pH with the Henderson-Hasselbalch equation, calculate how far the pH will move after you spike it with a known amount of HCl or NaOH, and judge whether your buffer is actually strong enough for the job. You will also be able to read a titration curve and spot the three mistakes that ruin most buffer preparations: choosing a buffer whose pKa is far from the target pH, confusing buffer capacity with buffer strength, and ignoring temperature or carbon dioxide.

Have these on hand before you start: a calibrated pH meter with at least two fresh standards, monobasic and dibasic phosphate salts (or acetic acid and sodium acetate), a volumetric flask, an analytical balance, a magnetic stirrer, and a burette or micropipette for the acid and base additions. A spreadsheet helps for the arithmetic, but every number below can be done by hand.

What Is a Buffer in Chemistry?

A buffer is a weak acid mixed with its conjugate base (or a weak base mixed with its conjugate acid) in solution. The mixture can neutralize added acid or added base because the two components interconvert. When you add H⁺, the conjugate base absorbs it and becomes the weak acid. When you add OH⁻, the weak acid donates a proton and becomes the conjugate base. The total amount of buffer in the solution does not change, only the ratio between the two forms.

That ratio is what sets the pH. The relationship is the Henderson-Hasselbalch equation:

pH = pKa + log10([A-] / [HA])

where [HA] is the concentration of the weak acid form, [A⁻] is the concentration of the conjugate base form, and pKa is the negative log of the acid dissociation constant of that particular acid. For a weak base buffer you write the same equation with the conjugate acid in the denominator.

This is why the answer to "what are buffers in chemistry" is not simply "solutions that hold pH steady." Buffers hold pH steady near a specific value determined by the pKa of the acid you chose. A buffer cannot be designed to sit at any arbitrary pH. It can only be designed to sit near the pKa of some available weak acid.

The Henderson-Hasselbalch relationship is the workhorse of buffer calculations across chemistry, biology, and clinical laboratory science. It is used to describe acid-base status in blood and other fluids, to calculate urinary bicarbonate, and to model buffer behavior in complex mixtures [1][2][3]. A newer predictive formula based on partitioning of added protons has been developed for multiple-buffer systems, but for a single buffer the Henderson-Hasselbalch form remains the standard tool [4].

The Two Components of a Buffer

Take acetate as the example. Acetic acid (CH₃COOH, written HA) is the weak acid. Sodium acetate (CH₃COO⁻Na⁺, giving A⁻ in solution) supplies the conjugate base. Mix them in water and you have an acetate buffer. The pKa of acetic acid is 4.76 at 25 °C.

Take phosphate as the second example. Monobasic potassium phosphate (KH₂PO₄) supplies H₂PO₄⁻, which acts as the weak acid. Dibasic potassium phosphate (K₂HPO₄) supplies HPO₄²⁻, the conjugate base. The relevant pKa for the H₂PO₄⁻/HPO₄²⁻ pair is 7.20 at 25 °C. That makes phosphate the standard choice for buffers near physiological pH.

The pairs matter. Phosphate has three pKa values (roughly 2.1, 7.2, and 12.3), so it can buffer in three separate pH windows. Acetate has one useful pKa near 4.76, so it buffers well only in the acidic range.

Buffers and pH: Why the pKa Sets the Working Range

A buffer works best when the pH you want is close to the pKa of the acid. At pH = pKa, the concentrations of HA and A⁻ are equal, the log term is zero, and the buffer resists change from both directions equally well. As you move away from pKa in either direction, one component dominates and the buffer becomes lopsided. It can still neutralize added acid or base, but it does so less efficiently.

The practical rule taught in every biochemistry lab is that a buffer is useful within about one pH unit of its pKa. Inside that window, the ratio [A⁻]/[HA] stays between roughly 0.1 and 10, and buffer capacity is at least a meaningful fraction of its maximum. Outside that window, the ratio becomes extreme (100:1 or 1:100) and the solution behaves more like a dilute acid or base than a buffer.

This is why you cannot make a phosphate buffer at pH 4 and expect it to hold. Phosphate's nearest pKa is 7.2, so at pH 4 the H₂PO₄⁻/HPO₄²⁻ pair is almost entirely in the acid form. Adding base will change pH quickly because there is very little conjugate base available to absorb protons in the reverse direction.

Buffer Capacity Versus Buffer Strength

These two terms get used interchangeably in casual conversation, and that causes real errors at the bench.

Buffer capacity (often written β) is a number. It is the amount of strong acid or base, in moles per liter, required to change the pH of one liter of the buffer by one pH unit. Van Slyke defined it as the quantity of base required to change pH by one unit per liter of sample [5]. A buffer with high capacity resists pH change strongly. A buffer with low capacity resists it weakly.

Buffer strength usually refers to the total molar concentration of the buffer species, for example "a 100 mM phosphate buffer." A 100 mM buffer at pH equal to its pKa has high capacity. The same 100 mM buffer at pH two units away from its pKa has much lower capacity, even though the total concentration is identical. So concentration alone does not tell you how well the buffer will hold pH. You need concentration and proximity to pKa.

The two ideas are related but not the same. Capacity is the functional property. Concentration is one of the inputs that determines it. The other input is the ratio of the two buffer forms.

A practical consequence: if you need a buffer to hold pH against a large acid load, increasing the total concentration helps, but only up to the solubility limit of your salts and only within the useful pKa window. Doubling the concentration of a buffer that is already two pH units from its pKa will not rescue it.

Dilution Effects

Diluting a buffer does not change its pH, because the ratio [A⁻]/[HA] is unchanged. Diluting a buffer does reduce its capacity, because capacity scales with total concentration. A 100 mM phosphate buffer diluted tenfold to 10 mM still sits at the same pH, but it now takes one tenth as much acid to move that pH by one unit.

This distinction matters when you prepare working solutions from stocks. If you dilute a 10× phosphate stock into a reaction mix, the pH of the final solution is the pH of the stock (assuming the diluent is neutral and does not itself buffer), but the buffering power in the final tube is ten times lower. Reactions that generate or consume protons will drift further in the diluted condition. When a protocol says "50 mM Tris, pH 7.5," that is the concentration in the final solution, not the stock.

Dilution effects also explain why a buffer can look fine in a beaker and fail in a microplate. The volume is smaller and the absolute amount of buffer is smaller, so a small proton load produces a larger pH shift.

A Worked Example: Phosphate Buffer at pH 7.20

This is the calculation you will do most often. We will prepare a 100 mM phosphate buffer at pH 7.20, then add a known amount of HCl and calculate the new pH.

Step 1: Confirm the pKa and Target pH

Phosphate's relevant pKa is 7.20 at 25 °C. Our target pH equals the pKa, so the ratio [HPO₄²⁻]/[H₂PO₄⁻] should be 1.00. This is the easiest case and the one with maximum capacity.

Step 2: Calculate the Ratio

pH = pKa + log10([A-]/[HA])
7.20 = 7.20 + log10([HPO4^2-]/[H2PO4^-])
log10([HPO4^2-]/[H2PO4^-]) = 0
[HPO4^2-]/[H2PO4^-] = 1.00

Step 3: Split the Total Concentration

Total phosphate = 100 mM = [HA] + [A⁻]. With a 1:1 ratio, each form is 50 mM.

Step 4: Add a Known Amount of Strong Acid

Suppose you add HCl to give a final concentration of 10 mM added H⁺. Strong acid converts conjugate base to weak acid:

HPO4^2- + H+ -> H2PO4^-

New [HPO₄²⁻] = 50 - 10 = 40 mM. New [H₂PO₄⁻] = 50 + 10 = 60 mM.

Step 5: Calculate the New pH

pH = 7.20 + log10(40/60)
pH = 7.20 + log10(0.667)
pH = 7.20 + (-0.176)
pH = 7.02

Adding 10 mM HCl to a 100 mM phosphate buffer at pH 7.20 moves the pH by 0.18 units. That is the buffer doing its job.

Step 6: Compare With Water

If you added 10 mM HCl to unbuffered water at pH 7.0, the pH would drop to about 2.0, a change of roughly 5 units. The buffer reduced that change by more than 25-fold. This comparison is the cleanest way to show what a buffer actually does.

Step 7: Check the Direction With Base

Now add 10 mM NaOH instead. Strong base converts weak acid to conjugate base:

H2PO4^- + OH- -> HPO4^2- + H2O

New [HPO₄²⁻] = 60 mM, new [H₂PO₄⁻] = 40 mM.

pH = 7.20 + log10(60/40)
pH = 7.20 + 0.176
pH = 7.38

The pH moves 0.18 units in the other direction. Symmetric, as expected at the pKa.

Summary Table of the Worked Example

Read the table from top to bottom. The first 20 mM of acid barely moves the pH. The next 20 mM moves it much more. This is the buffer capacity running out as the ratio departs from 1.00. By the time you have added 40 mM HCl, the buffer is nearly exhausted in the acid direction and the pH is falling fast.

The Same Calculation With Acetate

For comparison, an acetate buffer at pH 4.76 (its pKa) behaves identically in shape but at a different pH. A 100 mM acetate buffer at pH 4.76, spiked with 10 mM HCl, gives:

pH = 4.76 + log10(40/60) = 4.76 - 0.176 = 4.58

The pH shift is the same 0.18 units. The absolute pH is different because the pKa is different. This is the key insight: buffer behavior depends on the ratio, not on the identity of the acid, as long as you are near the pKa.

Reading a Titration Curve

A titration curve is a plot of pH on the y-axis against the volume of added strong acid or base on the x-axis. For a buffer, the curve has three regions.

The first region is the flat middle section. This is where the buffer is doing its work. The slope is shallow because added protons are being absorbed by the conjugate base rather than accumulating as free H⁺. The flattest point of this region is at pH = pKa, where the ratio is 1:1 and capacity is maximal.

The second region is the steep section on either side. As you exhaust one component of the buffer, the curve bends upward or downward sharply. The inflection points at the ends of the flat region mark the practical limits of the buffer, roughly one pH unit on either side of pKa.

The third region is the far end of the curve. Once the buffer is fully consumed, the pH is set by the excess strong acid or base alone, and the curve becomes steep again.

A well-designed experiment uses the flat region. If your working pH falls on the steep part of the curve, the buffer is not protecting you.

Titration curves are used routinely to characterize buffer materials. In one study of rumen buffers, three different titration protocols (static pH over 2 and 8 hours, fixed HCl load over 8 hours, and a 3-hour acidotic diet simulation with acetic acid) gave different rankings of the same materials, which shows that the shape of the curve and the method used to generate it both matter [6]. The same principle applies to any buffer: the curve is a property of the solution, and the way you titrate it determines what you measure.

Titration Curve of the Worked Example

If you plotted the phosphate buffer from the table above, you would see a nearly flat line from pH 7.20 down to about pH 6.8, then a gradual bend, then a steep drop below pH 6.3. The flat region corresponds to the first two rows of the table. The bend corresponds to row 4. The steep drop would appear if you continued adding HCl past 40 mM.

Titration Curve of a Bicarbonate Buffer

Bicarbonate buffers are different because they are open systems. Carbon dioxide can escape to the gas phase, and the pH depends on the partial pressure of CO₂ as well as the concentrations of HCO₃⁻ and H₂CO₃. The Henderson-Hasselbalch form for bicarbonate is often written with pKa 6.1 and a solubility coefficient for CO₂ [3][7]. In a closed vessel, a bicarbonate buffer titrates like any other buffer. In an open vessel, CO₂ escapes and the pH drifts upward. This is why bicarbonate buffers are prepared fresh and kept capped.

Choosing a Buffer for a Given pH

The decision tree is short.

First, identify the target pH. Second, look up the pKa of every weak acid you have access to. Third, pick the one whose pKa is closest to your target. Fourth, decide the total concentration based on how much acid or base you expect the system to generate. Fifth, prepare the buffer and verify the pH with a calibrated meter.

Common choices:

Phosphate and bicarbonate are the two most common buffers in pharmaceutical dissolution testing and in biological media [8]. Bicarbonate is preferred when you need to mimic physiological CO₂ conditions, and phosphate is preferred when you need a stable, non-volatile buffer at neutral pH.

One caution from the dissolution literature: even when two buffers have the same initial pH and the same buffer capacity, they can behave differently in the same experiment. In a comparison of high-dose salt-form drugs, the pH at 4 hours was lower in phosphate buffer than in bicarbonate buffer in every case, and the area under the dissolution curve was 1.2 to 2.6-fold lower in bicarbonate for drugs with an acid counterion [8]. The buffer identity, not just the pH and capacity, can change the outcome.

Step-by-Step Lab Procedure

This is the practical sequence for preparing and testing a buffer.

1. Weigh the salts. For 1 L of 100 mM phosphate buffer at pH 7.20, you need approximately 6.8 g of KH₂PO₄ and 8.7 g of K₂HPO₄. Adjust to your target volume and check the calculation with the Henderson-Hasselbalch equation.

1. Dissolve in about 80% of the final volume of deionized water. Do not bring to volume yet. The pH of a concentrated solution is easier to adjust.

1. Calibrate the pH meter with two standards that bracket your target. For a pH 7.20 buffer, use pH 7.00 and pH 10.01, or pH 4.01 and pH 7.00. A three-point calibration with 4.01, 7.00, and 10.01 is better if you work across a wide range.

1. Measure the pH while stirring. Add small volumes of 1 M HCl or 1 M NaOH to bring the pH to 7.20. Wait for the reading to stabilize before each addition.

1. Bring to final volume with deionized water. Measure the pH again. Dilution can shift the reading slightly if the ionic strength changes.

1. Record the final pH, the temperature, and the total buffer concentration. These three numbers define the buffer.

1. Store capped at the working temperature. If the buffer is bicarbonate-based, prepare it fresh and keep it sealed.

Expected Output

A correctly prepared 100 mM phosphate buffer at pH 7.20 should read 7.20 ± 0.05 on a freshly calibrated meter at 25 °C. If it reads 7.35 or 7.05, check the calibration standards first, then check the salt weights, then check the temperature.

Temperature Effects on pKa

The pKa of a weak acid is temperature-dependent. Phosphate's pKa shifts by roughly -0.0028 per °C near 25 °C, so a buffer prepared at 25 °C and used at 4 °C will read slightly higher than expected, and one used at 37 °C will read slightly lower. Tris is the worst offender among common buffers, with a pKa that drops about 0.03 units per °C. Phosphate and acetate are more stable.

The practical rule is to prepare and calibrate at the temperature you will use. If you prepare a Tris buffer at room temperature and then run an enzyme assay at 37 °C, the pH will be lower than the number on the bottle. For Tris, the shift can be 0.3 units or more across a 10 °C change, which is enough to change enzyme activity.

Temperature also affects the pH meter itself. The electrode response is temperature-dependent, and most modern meters have automatic temperature compensation. If your meter does not, or if the temperature probe is not in the solution, the reading will be wrong. Always let the buffer equilibrate to the working temperature before taking the final reading.

Carbon Dioxide Absorption in Bicarbonate Buffers

Bicarbonate buffers are uniquely sensitive to the atmosphere. Carbon dioxide from room air dissolves into the solution and shifts the equilibrium:

CO2 + H2O <-> H2CO3 <-> H+ + HCO3-

Adding CO₂ pushes the equilibrium to the right, generating H⁺ and lowering the pH. In a closed vessel, the CO₂ that dissolves stays in solution and the pH stabilizes at a new value. In an open vessel, CO₂ escapes and the pH drifts upward as the equilibrium shifts back.

The behavior of bicarbonate buffers under different gas pressures has been studied extensively. In rumen fluid, the buffering capacity under CO₂ was greater at low pH than the same fluid titrated in air, and dilution under constant CO₂ pressure decreased pH as predicted by the model [9]. In whole blood, the buffer value depends on the CO₂ partial pressure and on whether the titration is done under open or closed conditions [5]. These are not quirks of biology. They are direct consequences of the fact that CO₂ is a volatile buffer component.

For bench work, the rule is simple. Prepare bicarbonate buffers fresh, keep them capped, and minimize the headspace. If you need a stable pH at neutral for hours, use phosphate or HEPES instead. If you need to mimic physiological CO₂, use bicarbonate and accept that the pH will drift unless you control the gas phase.

pH Meter Calibration With Two or Three Standards

A pH meter is only as good as its calibration. Two standards are the minimum. Three are better when you work across a wide range.

Two-point calibration uses one standard near pH 7.00 and one at either pH 4.01 or pH 10.01, depending on your working range. This establishes the slope and offset of the electrode response. It is adequate for most routine work.

Three-point calibration adds a second extreme standard, so you bracket the entire range you will use. For a lab that runs both acidic and basic buffers, calibrate with 4.01, 7.00, and 10.01. The meter fits a curve through all three points and reports the slope. A healthy electrode gives a slope between 95% and 105% of theoretical. A slope below 90% means the electrode is aging or contaminated.

Practical points:

• Use fresh standards. An opened bottle of pH 10.01 buffer absorbs CO₂ from air and drifts downward within weeks.
• Rinse the electrode with deionized water between standards and blot gently. Do not wipe, which can scratch the glass.
• Stir the standard gently during measurement. The reading stabilizes faster.
• Check the slope after calibration. If it fails, clean the electrode with pepsin-HCl or a commercial cleaning solution and recalibrate.
• Recalibrate at least daily, and more often if you work at extreme pH or in high-ionic-strength solutions.

A meter that reads 7.00 in pH 7.00 standard but 10.20 in pH 10.01 standard has a slope error. The fix is cleaning or replacement, not adjustment of the buffer recipe.

Common Mistakes and Limitations

The first mistake is choosing a buffer whose pKa is far from the target pH. A 100 mM acetate buffer at pH 7.0 will barely buffer at all, because the ratio [A⁻]/[HA] is over 100:1 and there is almost no acetic acid left to donate protons. The fix is to pick a different buffer or accept a much lower capacity.

The second mistake is confusing buffer capacity with buffer strength. A "strong" 1 M buffer at the wrong pH has less useful capacity than a 50 mM buffer at the right pH. Always check the pKa before you check the concentration.

The third mistake is ignoring temperature. A buffer prepared at 25 °C and used at 4 °C or 37 °C will not have the pH you measured. Tris is the worst case. Prepare and calibrate at the working temperature.

The fourth mistake is leaving bicarbonate buffers open to air. CO₂ absorption or loss will shift the pH. Cap the bottle and prepare fresh.

The fifth mistake is trusting an uncalibrated or aging pH meter. A slope below 90% means the electrode is failing. Recalibrate with fresh standards and check the slope before you blame the buffer.

The sixth mistake is assuming dilution does not matter. Dilution preserves pH but reduces capacity. If your reaction generates protons, a diluted buffer will drift further than the stock.

The seventh mistake is using a buffer outside its useful range and expecting it to hold. The one-pH-unit rule is not a suggestion. It is a consequence of the Henderson-Hasselbalch equation.

A limitation worth stating plainly: buffer calculations assume ideal behavior. At high ionic strength, activity coefficients deviate from 1 and the measured pH can differ from the calculated value by 0.1 units or more. For critical work, measure the pH rather than trusting the arithmetic. And for any specific experimental system, the right buffer and the right concentration depend on the chemistry of that system. A veterinarian or a qualified lab supervisor should be consulted for work involving biological samples or regulated assays.

Frequently Asked Questions

What is a buffer in chemistry?

A buffer is a solution of a weak acid and its conjugate base that resists pH change when acid or base is added. The two forms interconvert to absorb added protons or hydroxide ions.

What are buffers in chemistry used for?

Buffers are used to hold pH steady in reactions, assays, and biological systems. They are also used in dissolution testing, fermentation, and any process where pH drift would change the outcome.

How do buffers affect pH in the lab?

Buffers reduce the pH change caused by added acid or base. The size of the reduction depends on the buffer concentration and on how close the working pH is to the buffer's pKa.

What is the pH of a buffer?

The pH of a buffer is set by the Henderson-Hasselbalch equation: pH = pKa + log10([A⁻]/[HA]). At equal concentrations of the two forms, pH equals pKa.

Why does buffer capacity peak near the pKa?

Buffer capacity peaks at pH = pKa because the concentrations of the acid and base forms are equal, so the solution can absorb added acid or base from either direction with equal efficiency.

Does diluting a buffer change its pH?

Diluting a buffer does not change its pH because the ratio of the two forms is unchanged. It does reduce the buffer capacity, because capacity scales with total concentration.

Why do bicarbonate buffers drift?

Bicarbonate buffers drift because CO₂ exchanges with the atmosphere. Absorption lowers the pH and loss raises it. Capping the container and preparing fresh minimizes the drift.

How often should a pH meter be calibrated?

Calibrate at least once per day, and more often if you work at extreme pH or in high-ionic-strength solutions. Use fresh standards and check the electrode slope after every calibration.

Related Articles

• DNA Ligase Storage Buffer Composition and Its Effect on Activity
• Reference Intervals in Clinical Chemistry: Establishment and Verification
• What Is a Buffer in Biology? Definition and Examples
• Computational Biology vs. Computational Chemistry
• Buffer Preparation: A Step-by-Step Guide for Students
• How to Prepare Buffer Solutions: A Step-by-Step Guide
• Phenyl Group: Structure, Chemistry, and Examples
• What Does Reduction Mean in Biology and Chemistry?

Sources

1. [[Clinical evaluation of acid-base status: Henderson-Hasselbalch, or Stewart-Fencl approach?].](https://pubmed.ncbi.nlm.nih.gov/27990831/)
2. Clinical assessment of acid-base status: comparison of the Henderson-Hasselbalch and strong ion approaches.
3. Determination of urinary bicarbonate with the Henderson-Hasselbalch equation. Comparison using two different methods.
4. Calculation of the equilibrium pH in a multiple-buffered aqueous solution based on partitioning of proton buffering: a new predictive formula.
5. An estimation of buffer values of human whole blood by titration experiment under the open condition for carbon dioxide gas.
6. Evaluation and Development of Analytical Procedures to Assess Buffering Capacity of Carbonate Ruminant Feed Buffers.
7. In vivo microdialysis of 2-deoxyglucose 6-phosphate into brain: a novel method for the measurement of interstitial pH using 31P-NMR.
8. Dissolution profiles of high-dose salt-form drugs in bicarbonate buffer and phosphate buffer.
9. Calculation of the buffering capacity of bicarbonate in the rumen and in vitro.