Buffer Solutions: Types, Purpose, and Applications in Biology

By Dr. Zubair Khalid, DVM, MS, PhD ·

Buffer Solutions: Types, Purpose, and Applications in Biology

Introduction to Buffer Solutions

Definition and Basic Composition

A buffer solution is an aqueous system that resists changes in pH when small quantities of an acid or a base are added, or when the solution is diluted. This resistance is not absolute—buffers work within a defined pH range and have a finite capacity—but within those limits, they maintain hydrogen ion concentration ([H⁺]) remarkably steady.

Every buffer consists of two components: a weak acid and its conjugate base, or a weak base and its conjugate acid. The weak acid form can donate protons (H⁺) to neutralize added base, while the conjugate base form can accept protons to neutralize added acid. Because both species exist in equilibrium, the system can absorb perturbations from either direction.

The equilibrium for a weak acid buffer is written as:

HA ⇌ H⁺ + A⁻

where HA is the weak acid and A⁻ is its conjugate base. The position of this equilibrium is described by the acid dissociation constant, Ka:

Ka = [H⁺][A⁻] / [HA]

Taking the negative logarithm of both sides yields the Henderson–Hasselbalch equation, which is the central tool for buffer design:

pH = pKa + log([A⁻]/[HA])

Here, pKa = −log(Ka). When [A⁻] = [HA], the pH equals the pKa, and the buffer is at its most effective point.

Why pH Stability Matters in Biological Systems

Living cells operate within narrow pH windows. The cytoplasmic pH of most cells is approximately 7.2, and deviations of more than a few tenths of a pH unit can disrupt metabolism, denature proteins, and arrest growth. This sensitivity arises because pH directly affects:

  • Enzyme catalysis: Most enzymes have an optimal pH at which their active-site residues are in the correct protonation state for substrate binding and catalysis. For example, pepsin works at pH 2 in the stomach, while trypsin works at pH 8 in the small intestine.
  • Protein structure: Ionizable side chains (histidine, cysteine, lysine, aspartate, glutamate, arginine, tyrosine) participate in hydrogen bonds, salt bridges, and electrostatic interactions that stabilize tertiary and quaternary structure. pH changes alter these interactions.
  • Membrane potential and transport: Proton gradients drive ATP synthesis, nutrient uptake, and ion transport. Uncontrolled pH would collapse these gradients.
  • Nucleic acid stability: DNA duplex stability depends on hydrogen bonding between base pairs, which is pH-sensitive. Extreme pH causes depurination or strand separation.

Biological fluids therefore contain buffers—most notably the bicarbonate system in blood, phosphate in intracellular compartments, and proteins themselves, whose ionizable side chains contribute buffering capacity. In the laboratory, we replicate this stability with synthetic buffers to protect biomolecules during isolation, analysis, and manipulation.

The Mechanism of Buffer Action

How Buffers Resist pH Change

Buffer action is a direct application of Le Chatelier's principle: when a system at equilibrium is disturbed, it shifts to counteract the disturbance.

Consider a buffer containing acetic acid (CH₃COOH) and sodium acetate (CH₃COONa). The equilibrium is:

CH₃COOH ⇌ H⁺ + CH₃COO⁻

When acid (HCl) is added: The added H⁺ combines with the acetate ion (CH₃COO⁻), the conjugate base, to form undissociated acetic acid. The equilibrium shifts left, consuming the added protons. The net reaction is:

H⁺ + CH₃COO⁻ → CH₃COOH

The pH barely changes because free H⁺ is removed from solution.

When base (NaOH) is added: The added OH⁻ reacts with H⁺ to form water. This removes H⁺ from the equilibrium, shifting it to the right. Acetic acid dissociates to replenish the consumed H⁺:

OH⁻ + CH₃COOH → CH₃COO⁻ + H₂O

Again, the pH change is minimal because the buffer supplies protons to neutralize the hydroxide.

The key is that both HA and A⁻ are present in significant concentrations. If only the weak acid were present, adding base would simply titrate it, and pH would rise steeply. If only the conjugate base were present, adding acid would consume it rapidly. A buffer works because it has both forms available to absorb the perturbation.

The Henderson–Hasselbalch Equation

The Henderson–Hasselbalch equation is the practical tool for buffer design:

pH = pKa + log([A⁻]/[HA])

This equation has three immediate uses:

  1. Predicting pH from known concentrations of acid and conjugate base.
  2. Calculating the ratio of conjugate base to acid needed to achieve a target pH.
  3. Determining the pH range over which a buffer is effective.

A buffer is most effective when the ratio [A⁻]/[HA] is between 0.1 and 10, which corresponds to pH = pKa ± 1. Outside this range, one species is so dilute that it cannot neutralize significant additions of acid or base. The buffer capacity—the amount of acid or base a buffer can absorb before pH changes by one unit—is maximal when pH = pKa and decreases as you move away from the pKa.

For example, to prepare a phosphate buffer at pH 7.0 using the H₂PO₄⁻/HPO₄²⁻ pair (pKa₂ = 7.21), the required ratio is:

7.0 = 7.21 + log([HPO₄²⁻]/[H₂PO₄⁻]) log([HPO₄²⁻]/[H₂PO₄⁻]) = −0.21 [HPO₄²⁻]/[H₂PO₄⁻] = 0.62

So you would need approximately 1.6 parts H₂PO₄⁻ to 1 part HPO₄²⁻.

Types of Buffer Solutions

Acidic Buffers

Acidic buffers maintain pH below 7. They consist of a weak acid and its salt with a strong base. The weak acid provides the HA form, and the salt provides the A⁻ form.

Common examples include:

  • Acetic acid/sodium acetate: pKa = 4.76, effective range pH 3.8–5.8
  • Citric acid/sodium citrate: pKa₁ = 3.13, pKa₂ = 4.76, pKa₃ = 6.40, effective across pH 2.1–7.4
  • Formic acid/sodium formate: pKa = 3.75, effective range pH 2.8–4.8

These buffers are used when working with acidic conditions, such as in certain enzyme assays, chromatographic separations, or when studying acid-stable proteins.

Basic Buffers

Basic buffers maintain pH above 7. They consist of a weak base and its salt with a strong acid. The weak base provides the B form, and the salt provides the BH⁺ form.

The equilibrium is:

B + H₂O ⇌ BH⁺ + OH⁻

And the Henderson–Hasselbalch form becomes:

pH = pKa + log([B]/[BH⁺])

where pKa refers to the conjugate acid BH⁺.

Common examples include:

  • Ammonia/ammonium chloride: pKa = 9.25, effective range pH 8.3–10.3
  • Tris (tris(hydroxymethyl)aminomethane)/Tris-HCl: pKa = 8.06 at 25°C, effective range pH 7.1–9.0
  • Bicarbonate/carbonate: pKa₂ = 10.33, effective range pH 9.3–11.3

Tris is particularly important in molecular biology because its pKa is close to physiological pH and it is compatible with most enzymes.

Common Examples of Buffer Solutions

Acetate Buffer

The acetic acid/sodium acetate system is the classic acidic buffer. Acetic acid (CH₃COOH) has a pKa of 4.76, making it ideal for pH 3.8–5.8.

Preparation: Mix sodium acetate (CH₃COONa) and acetic acid in the appropriate ratio. For a 0.1 M buffer at pH 5.0:

5.0 = 4.76 + log([CH₃COO⁻]/[CH₃COOH]) log([CH₃COO⁻]/[CH₃COOH]) = 0.24 [CH₃COO⁻]/[CH₃COOH] = 1.74

So for 100 mL of buffer, you would dissolve 0.1 mol of sodium acetate and 0.057 mol of acetic acid in water, then adjust to final volume.

Applications: Acetate buffers are used in DNA extraction protocols, protein precipitation, and certain enzyme assays. They are also used in histology for tissue fixation and staining.

Phosphate Buffer

The phosphate system is the most important biological buffer. It has three pKa values:

  • pKa₁ = 2.15 (H₃PO₄/H₂PO₄⁻)
  • pKa₂ = 7.21 (H₂PO₄⁻/HPO₄²⁻)
  • pKa₃ = 12.32 (HPO₄²⁻/PO₄³⁻)

The second pKa (7.21) is close to physiological pH, making the H₂PO₄⁻/HPO₄²⁻ pair the dominant buffer in intracellular fluid. Phosphate buffers are prepared by mixing monosodium phosphate (NaH₂PO₄) and disodium phosphate (Na₂HPO₄) in the appropriate ratio.

Applications: Phosphate-buffered saline (PBS) is ubiquitous in biological research. It maintains pH 7.4 and provides physiological osmolarity. It is used for washing cells, diluting samples, and as a base for many assay buffers. Phosphate buffers are also used in protein purification and chromatography.

Bicarbonate Buffer

The bicarbonate system (H₂CO₃/HCO₃⁻) is the primary buffer in blood plasma. The relevant equilibrium is:

CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻

The pKa₁ for carbonic acid is 6.35, but the system is "open" because CO₂ is continuously removed by the lungs. This makes the effective buffering range broader than a closed system would suggest.

Applications: Bicarbonate buffers are used in cell culture media (e.g., DMEM, RPMI) to maintain pH 7.2–7.4. They require a CO₂ incubator (typically 5% CO₂) to maintain the equilibrium. The buffer is also central to blood pH regulation, where the ratio of HCO₃⁻ to CO₂ is approximately 20:1 at pH 7.4.

Purpose and Importance of Buffer Solutions

Role in Biological Systems

Biological organisms face constant acid-base challenges. Cellular metabolism generates CO₂, lactic acid, and other acidic byproducts. Without buffering, pH would plummet and kill the cell. The major biological buffers are:

  • Bicarbonate: Dominant in extracellular fluid and blood
  • Phosphate: Dominant in intracellular fluid and urine
  • Proteins: Hemoglobin in red blood cells, albumin in plasma, and intracellular proteins all contribute buffering capacity through their ionizable side chains

The bicarbonate system illustrates how buffers integrate with physiology. When H⁺ increases in blood, it combines with HCO₃⁻ to form H₂CO₃, which dissociates to CO₂ and H₂O. The CO₂ is exhaled, removing acid from the body. Conversely, when H⁺ decreases, the reaction reverses, and CO₂ is retained. This coupling of chemical buffering with respiratory control provides a powerful homeostatic mechanism.

Role in Laboratory Experiments

In the laboratory, buffers serve several critical functions:

  1. Maintaining enzyme activity: Enzymes have narrow pH optima. Assays must be buffered to keep the pH constant throughout the reaction, even as products accumulate.
  2. Protecting nucleic acids: DNA and RNA are susceptible to acid-catalyzed depurination and alkaline hydrolysis. Buffers keep them stable during storage and manipulation.
  3. Ensuring reproducible results: Many analytical techniques—electrophoresis, chromatography, spectroscopy—depend on pH. A drift of 0.1 pH units can change migration patterns or binding affinities.
  4. Providing ionic strength: Buffers also supply ions that are necessary for biological activity. Many enzymes require specific ions (Mg²⁺, K⁺) for catalysis, and buffers often include these.

Applications of Buffer Solutions in Biology and Biotechnology

Electrophoresis Buffers

Gel electrophoresis separates nucleic acids or proteins based on size and charge. The running buffer must maintain constant pH because the electrical field can generate H⁺ and OH⁻ at the electrodes.

For DNA electrophoresis, the standard buffer is TAE (Tris-acetate-EDTA) or TBE (Tris-borate-EDTA):

  • TAE: 40 mM Tris-acetate, 1 mM EDTA, pH 8.0. Provides good resolution for large DNA fragments (>2 kb) but has lower buffering capacity than TBE.
  • TBE: 89 mM Tris-borate, 2 mM EDTA, pH 8.3. Higher buffering capacity, better for small fragments and long runs.

The EDTA chelates divalent cations (Mg²⁺), which are cofactors for nucleases, thereby protecting the DNA from degradation.

For protein electrophoresis, the standard system is SDS-PAGE using a Tris-glycine buffer. The running buffer is 25 mM Tris, 192 mM glycine, 0.1% SDS, pH 8.3. Glycine acts as a trailing ion that stacks the proteins into a sharp band before separation.

PCR and Enzyme Assays

PCR requires a buffer that maintains pH 8.3–9.0 at the extension temperature (72°C) and provides Mg²⁺ as a cofactor for Taq DNA polymerase. The standard PCR buffer contains:

  • 10–50 mM Tris-HCl, pH 8.3–9.0
  • 50 mM KCl (for ionic strength)
  • 1.5–2.5 mM MgCl₂ (optimal concentration varies by template and primers)
  • 0.1% Triton X-100 or other stabilizers

The buffer must also withstand the thermal cycling (denaturation at 95°C, annealing at 50–65°C, extension at 72°C) without significant pH drift. Tris has a temperature coefficient of −0.028 pH/°C, meaning its pH decreases as temperature increases. A buffer prepared at pH 8.3 at 25°C will be approximately pH 7.9 at 72°C, which is still within the active range of Taq polymerase.

Enzyme assays for kinases, phosphatases, proteases, and other enzymes all use buffers matched to the enzyme's pH optimum. For example, alkaline phosphatase assays use a 0.1 M glycine-NaOH buffer at pH 10.4, while pepsin assays use 0.01 M HCl at pH 2.0.

Cell Culture Media

Mammalian cell culture media must maintain pH 7.2–7.4, provide nutrients, and maintain osmolarity. The two main buffering systems are:

  • Bicarbonate/CO₂: Used in most standard media (DMEM, RPMI 1640). Requires a 5–10% CO₂ incubator. The medium contains NaHCO₃ (typically 3.7 g/L for DMEM), and the CO₂ in the incubator maintains the equilibrium. The phenol red pH indicator turns yellow below pH 6.8 and purple above pH 8.2.
  • HEPES (4-(2-hydroxyethyl)-1-piperazineethanesulfonic acid): A zwitterionic buffer with pKa = 7.55 at 25°C. It is used at 10–25 mM in media for applications where CO₂ control is impractical (e.g., during cell sorting or microscopy). HEPES does not require a CO₂ incubator but is more expensive than bicarbonate.

Protein Purification

Protein purification relies on buffers at every step:

  • Lysis buffers: Contain Tris or phosphate at pH 7.4–8.0, plus salts, detergents, and protease inhibitors to release and stabilize proteins.
  • Chromatography buffers: Ion exchange, affinity, and size exclusion chromatography all require buffers matched to the protein's stability and the column chemistry. For example, Ni-NTA purification of His-tagged proteins uses a 20 mM sodium phosphate, 500 mM NaCl buffer at pH 7.4, with imidazole (20–500 mM) for elution.
  • Storage buffers: Purified proteins are stored in buffers containing glycerol (10–50%) and reducing agents (DTT or β-mercaptoethanol) to prevent oxidation and aggregation.

How to Prepare a Buffer Solution

Calculating Buffer Components

The preparation of a buffer with a target pH requires three pieces of information: the desired pH, the pKa of the buffering species, and the desired molarity.

Step 1: Choose the buffering species. Select a weak acid/base whose pKa is within 1 unit of the target pH. For pH 7.4, phosphate (pKa₂ = 7.21) or HEPES (pKa = 7.55) are appropriate.

Step 2: Calculate the ratio. Use the Henderson–Hasselbalch equation:

pH = pKa + log([A⁻]/[HA])

For a phosphate buffer at pH 7.4:

7.4 = 7.21 + log([HPO₄²⁻]/[H₂PO₄⁻]) log([HPO₄²⁻]/[H₂PO₄⁻]) = 0.19 [HPO₄²⁻]/[H₂PO₄⁻] = 1.55

Step 3: Calculate individual concentrations. If the total phosphate concentration is 0.1 M:

[HPO₄²⁻] + [H₂PO₄⁻] = 0.1 M [HPO₄²⁻] = 1.55 × [H₂PO₄⁻]

Substituting:

1.55[H₂PO₄⁻] + [H₂PO₄⁻] = 0.1 M [H₂PO₄⁻] = 0.039 M [HPO₄²⁻] = 0.061 M

Step 4: Weigh the salts. For a 1 L buffer, weigh 0.039 mol of NaH₂PO₄ (monobasic, MW = 119.98 g/mol) and 0.061 mol of Na₂HPO₄ (dibasic, MW = 141.96 g/mol):

NaH₂PO₄: 0.039 × 119.98 = 4.68 g Na₂HPO₄: 0.061 × 141.96 = 8.66 g

Dissolve both in approximately 800 mL of distilled water, then adjust to 1 L.

Adjusting pH

In practice, buffers are often prepared by dissolving one salt and titrating with acid or base. For example, to prepare 1 L of 0.1 M Tris-HCl at pH 8.0:

  1. Dissolve 12.11 g of Tris base (MW = 121.14 g/mol) in 800 mL of water.
  2. Add concentrated HCl dropwise while stirring, monitoring the pH with a calibrated pH meter.
  3. When the pH reaches 8.0, adjust the volume to 1 L with water.

Critical considerations during pH adjustment:

  • Temperature matters: Tris pH decreases by 0.028 units per °C increase. A buffer adjusted to pH 8.0 at 25°C will be pH 7.7 at 37°C. Decide at which temperature the buffer will be used and adjust accordingly.
  • Dilution effect: Adding acid or base changes the volume. If you adjust pH before bringing to final volume, the pH will shift slightly upon dilution. The standard practice is to bring to ~90% of final volume, adjust pH, then bring to final volume and re-check pH.
  • Equilibration: After adding acid or base, allow the solution to stir for 1–2 minutes before reading the pH, as the electrode response takes time.

Factors Affecting Buffer Capacity and pH

Buffer Capacity

Buffer capacity (β) is the amount of strong acid or base that must be added to change the pH of 1 L of buffer by 1 pH unit. It depends on:

  1. Concentration: Higher concentrations of buffer components give greater capacity. A 0.1 M buffer can absorb ten times more acid than a 0.01 M buffer.
  2. Ratio of [A⁻] to [HA]: Capacity is maximal when [A⁻] = [HA] (pH = pKa). As the ratio deviates from 1, capacity decreases.
  3. Distance from pKa: Buffers are effective within pH = pKa ± 1. Beyond this range, capacity drops sharply.

The relationship is expressed as:

β = 2.303 × C × ([H⁺] × Ka) / ([H⁺] + Ka)²

where C is the total buffer concentration ([HA] + [A⁻]).

Temperature Effects

Temperature affects buffer pH through two mechanisms:

  1. Changes in pKa: The pKa of most buffers changes with temperature. Tris has a large temperature coefficient (−0.028 pH/°C), while phosphate has a smaller one (−0.003 pH/°C). HEPES has a coefficient of −0.014 pH/°C.
  2. Changes in water ionization: The pKw of water changes from 14.00 at 25°C to 13.62 at 50°C, affecting the pH of unbuffered solutions.

For experiments at non-ambient temperatures (e.g., enzyme assays at 37°C, PCR at 72°C), prepare the buffer at the temperature of use, or use a buffer with a small temperature coefficient.

Ionic Strength

Ionic strength affects buffer pH by altering activity coefficients. The Henderson–Hasselbalch equation uses concentrations, but strictly speaking, it should use activities. At high ionic strength, the effective concentration of ions is lower than the nominal concentration, shifting the apparent pKa.

For most biological buffers at concentrations below 0.2 M, this effect is small (<0.05 pH units). However, in high-salt conditions (e.g., 1 M NaCl in chromatography buffers), the pH can shift by 0.1–0.2 units. Always check the pH after adding salts.

Common Pitfalls and Troubleshooting

Mistakes in Buffer Preparation

1. Using the wrong pKa: Selecting a buffer whose pKa is more than 1 unit from the target pH. This produces a buffer with poor capacity. For pH 7.0, using acetate (pKa = 4.76) would require a ratio of [A⁻]/[HA] = 174, meaning the buffer contains almost no HA form and cannot neutralize added base.

2. Ignoring temperature: Adjusting Tris buffer pH at 25°C but using it at 37°C. The pH will be 0.3 units lower than intended, potentially inactivating enzymes.

3. Incorrect salt hydration: Many buffer salts are hydrates. Na₂HPO₄·7H₂O (MW = 268.07) and Na₂HPO₄·12H₂O (MW = 358.14) have very different molecular weights. Using the wrong hydrate gives the wrong concentration.

4. Not accounting for the acid/base used for pH adjustment: When preparing Tris-HCl by titrating Tris base with HCl, the chloride ions from the HCl contribute to the ionic strength. For a 0.1 M Tris buffer at pH 8.0, approximately 0.07 M HCl is needed, giving a final chloride concentration of 0.07 M.

5. Forgetting to filter: Buffers for cell culture or molecular biology should be filtered (0.22 μm) to remove particulates and microorganisms. Autoclaving is appropriate for some buffers but not others (Tris buffers should not be autoclaved as they degrade).

Troubleshooting pH Drift

Symptom: pH drifts upward over time

  • Cause: Absorption of CO₂ from air. This is common for alkaline buffers (pH > 8). CO₂ dissolves to form carbonic acid, which lowers pH. Conversely, loss of CO₂ from bicarbonate buffers raises pH.
  • Solution: Store buffers in tightly sealed containers, minimize exposure to air, and prepare fresh buffers regularly.

Symptom: pH drifts downward over time

  • Cause: Microbial growth. Bacteria and fungi metabolize buffer components, producing acidic byproducts.
  • Solution: Add sodium azide (0.02%) for long-term storage, or prepare fresh.

Symptom: pH changes when the buffer is diluted

  • Cause: The buffer was not at its pKa, or the ionic strength changed significantly. Dilution of a buffer at pH ≠ pKa shifts the ratio [A⁻]/[HA] slightly because activity coefficients change.
  • Solution: Prepare buffers at the concentration of use, or verify pH after dilution.

Symptom: pH meter gives inconsistent readings

  • Cause: Electrode contamination, improper calibration, or temperature mismatch between the buffer and the calibration standards.
  • Solution: Calibrate the pH meter with fresh standards at the same temperature as the buffer. Clean the electrode according to the manufacturer's instructions.

Frequently Asked Questions

What is a buffer solution?

A buffer solution is an aqueous mixture of a weak acid and its conjugate base (or a weak base and its conjugate acid) that resists changes in pH when small amounts of acid or base are added. It works by neutralizing added H⁺ or OH⁻ through the equilibrium between the two buffer components.

What are the types of buffer solutions?

Buffers are classified as acidic (pH < 7) or basic (pH > 7). Acidic buffers consist of a weak acid and its salt with a strong base, such as acetic acid/sodium acetate. Basic buffers consist of a weak base and its salt with a strong acid, such as ammonia/ammonium chloride. Within these categories, buffers are further distinguished by their pKa values, which determine their effective pH range.

What is the purpose of a buffer solution?

The purpose of a buffer is to maintain a stable pH. In biological systems, buffers keep intracellular and extracellular pH within narrow ranges necessary for enzyme function, protein stability, and metabolic processes. In the laboratory, buffers ensure reproducible conditions for reactions, separations, and analyses.

Can you give an example of a buffer solution?

The acetic acid/sodium acetate buffer is a classic example. Acetic acid (CH₃COOH) has a pKa of 4.76. A buffer prepared by mixing equal molar amounts of acetic acid and sodium acetate has a pH of 4.76. This buffer effectively maintains pH in the range 3.8–5.8 and is used in DNA extraction and protein precipitation protocols.

Why is a buffer solution used in biological experiments?

Biological experiments use buffers to protect biomolecules from pH-induced damage. Enzymes lose activity outside their optimal pH range, DNA can depurinate at low pH, and proteins can denature or aggregate. Buffers also provide the ionic environment required for biological activity and ensure that experimental results are reproducible.

What are the applications of buffer solutions?

Buffer solutions are used in gel electrophoresis (TAE, TBE, Tris-glycine), PCR (Tris-HCl with MgCl₂), cell culture media (bicarbonate/CO₂, HEPES), protein purification (lysis, chromatography, and storage buffers), enzyme assays, and countless other molecular biology techniques. They are also central to physiological function, including blood pH regulation.

How does a buffer solution work?

A buffer works through the equilibrium between a weak acid (HA) and its conjugate base (A⁻). When acid is added, the conjugate base neutralizes it: H⁺ + A⁻ → HA. When base is added, the weak acid neutralizes it: OH⁻ + HA → A⁻ + H₂O. The buffer continues to resist pH change until one of the two components is depleted.

Key Takeaways

  • A buffer solution consists of a weak acid and its conjugate base (or weak base and conjugate acid) and resists pH change through equilibrium-based neutralization of added H⁺ or OH⁻.
  • The Henderson–Hasselbalch equation (pH = pKa + log([A⁻]/[HA])) is the fundamental tool for calculating buffer composition and predicting pH.
  • Buffers are effective within approximately pH = pKa ± 1, with maximum capacity at pH = pKa.
  • Common biological buffers include acetate (pKa 4.76), phosphate (pKa₂ 7.21), bicarbonate (pKa₁ 6.35), and Tris (pKa 8.06).
  • Buffers are essential in biological systems for maintaining enzyme activity, protein stability, and membrane function, and in the laboratory for electrophoresis, PCR, cell culture, and protein purification.
  • Buffer preparation requires attention to pKa selection, concentration calculations, temperature effects, and proper pH adjustment.
  • Common errors include using the wrong pKa, ignoring temperature coefficients, using incorrect salt hydrates, and failing to account for pH drift from CO₂ absorption or microbial growth.

Further Reading

  • Safa BN et al. Exposure to buffer solution alters tendon hydration and mechanics. Journal of biomechanics. 2017. PubMed 28720201
  • Wang P et al. Buffer solution induced highly crystalline sodium-rich Prussian blue for sodium storage. Chemical communications (Cambridge, England). 2024. PubMed 38230427
  • Ali M et al. Effect of plasma-activated water and buffer solution combined with ultrasound on fungicide degradation and quality of cherry tomato during storage. Ultrasonics sonochemistry. 2023. PubMed 37269690
  • Ali M, Cheng JH, Sun DW. Effect of plasma activated water and buffer solution on fungicide degradation from tomato (Solanum lycopersicum) fruit. Food chemistry. 2021. PubMed 33636620
  • Jörres A. Is introduction of the bicarbonate buffer in PD solution unavoidable?. Peritoneal dialysis international : journal of the International Society for Peritoneal Dialysis. 1996. PubMed 8914173
  • Featherstone RL et al. Hypothermic preservation of isolated rat lungs in modified bicarbonate buffer, EuroCollins solution or St Thomas' Hospital cardioplegic solution. European journal of cardio-thoracic surgery : official journal of the European Association for Cardio-thoracic Surgery. 1998. PubMed 986020800188-2)

Related Topics

Related Clinical & Scientific Guides