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Guide for optimal pH

Your guide for optimal pH

Identify the ideal buffer formulation to establish and maintain optimal pH conditions for your specific application, ranging from versatile options like Tris-HCl, to highly customized formulations with tight tolerances.

 

INTRODUCTION

 

The importance of maintaining optimal pH

 

Maintaining a controlled pH environment is critical for the support of cellular processes ranging from metabolism to cell growth to membrane potential. The right buffer can help you keep the acidity of a solution within an optimal physiological range, impacting the rate and efficiency of your chemical reactions and the recovery and purity of your products.

 

The charts below can help you find the ideal buffer to establish and maintain optimal pH by taking into account your type of organism, the qualities of different buffer types, the impact of pKa (concentration and temperature), and the complexation of metals in your experimental environment.

You can find the optimal pH range for a selection of common organisms below.

Note: Most bacteria grow over a pH range of approximately 3 units.

Buffers act as a neutralizing agent that helps maintain an optimal pH range. The ideal buffer for the conditions needed to support biological processes should have a pKa close to the optimal pH of the desired cells. The pKa indicates the strength of an acid, with lower pKa indicating a stronger acid. Using the right buffer in the right proportions and concentration can help you maintain a steady pH for longer and prevent pallet formation. 

 

Use the chart below to find the ideal buffer to help maitain your desired pH range throughout your specific application.

When selecting a buffer for your application, it is also important to consider the metal complexation characteristics. When a buffer forms complexes with metal ions, protons are released, which can cause pH to decrease dramatically. In addition, many enzymes need metal ions to sustain their functionality.

 

Buffers with low metal-binding constants are ideal when working with metal-dependent enzymes. For example, HEPES and MOPS are suitable for the use in solutions containing metal ions. Buffers with certain higher metal-binding constants, like Bicine, are less suitable, and should only be used in the absence of specific metals.

 

The chart below provides an overview of the metal complexation of common buffers:

Explore related products and applications

All our products are fully customizable, from formulation to format to fill volume. Quickly adjust the pH of a buffer with our Build•Tek™ Custom Configurator, or consult with our experts to find the best solution for your specific application.

Applications:

Acetate products:

Applications:

Bis-Tris products:

Applications:

Glycine products:

Applications:

HEPES products:

Applications:

MES products:

Applications:

MOPS products:

Applications:

Phosphate products:

Applications:

TAPS products:

Applications:

Tris products:

Frequently asked questions

You can find answers here for our most commonly asked questions. If you don't find the information you are looking for, get in touch with our team. We're here to help.

pH quantifies the measure of the acidity and alkalinity of an aqueous or liquid solution. pH calculation involves relative proportions of hydrogen (H+) and hydroxide (OH-) ions in a solution, such as the concentration of accessible H+.

 

pH = −log[H+]

 

pH = 14.00 – pOH (for aqueous solutions at 25°C)

 

A lower pH indicates a high H+ concentration (i.e., low OH- concentration), while a higher pH indicates a lower H+ concentration (i.e., high OH- concentration).

 

 

For example, a solution with a pH of 7 has a 10 times greater H+ concentration than one with a pH of 8.

Biological buffers contain either a weak acid and salt containing its conjugate base (e.g., acetic acid: CH3COOH and sodium acetate: CH3COONa) or a weak base and salt containing its conjugate acid (e.g., ammonia: NH3 and ammonium chloride: NH4Cl).

 

A biological buffer resists change in pH due to a chemical equilibrium that is established by the donation or acceptance of protons (H+). For instance, acetic acid (CH3COOH) can act as a weak acid and donate H+ to a hydroxide ion in solution (OH-), forming H2O and its conjugate base, CH3COO-. CH3COO- in turn, can accept the H+ from H2O. These interactions are governed by rules known as the Brønsted-Lowry theory—shown below.

 

A solution’s pH is typically controlled by using a buffer system, which must have a suitable pKa. The pKa value of buffers can be influenced by temperature, pressure, ionic strength, and the polarity of the solvent.

 

pKa is a dissociation constant measuring the strength of an acid. It tells us about how tightly the acid holds onto its H+. So, a lower pKa value means it is easier for the acid to donate its H+, making it a strong acid. Conversely, a higher pKa value means the acid holds onto its H+ tightly (harder to donate), making it a weak acid.

 

The relationship between pH and pKa is described by the Henderson–Hasselbalch equation:

 

pH = pKa + log10 ([A–]/[HA])

 

 

How do I use pKa when choosing a biological buffer?

 

Any biological buffer's buffering capacity (i.e., ability to maintain a stable pH level) is limited by its capacity to counter the increase in the H+ or OH- in the solution. 

 

According to the Henderson–Hasselbalch equation, when 50% of a weak acid is dissociated (i.e., [conjugate base] = [weak acid]), pH equals pKa—see the example for acetic acid below. If the pH rises, the weak acid will serve as the neutralizing agent; if the pH drops, the conjugate base will serve as a neutralizing agent.

 

 

As a general rule, a buffer whose pKa is a unit above or below the desired pH value would provide optimal buffering (i.e., pKa = desired pH ± 1). For example, if we want to prepare a solution with a pH of 5.0, we’d choose a buffer with pKa of 4.0-6.0.

 

However, if we anticipate a drop in pH (i.e., an increase in H+) during a reaction, choosing a buffer with pKa = desired pH - 1 would be more suitable since the buffer would contain more [conjugate base] to neutralize the H+ influx. Similarly, if we anticipate a rise in pH, choosing a buffer with pKa = desired pH + 1 would be more suitable since the buffer would contain more [weak acid] to neutralize the OH- influx (refer to Henderson–Hasselbalch equation).

 

 

How does the ionic strength and polarity of the solvent influence buffer pKa?

 

The ionic strength and polarity of the solvent can have a minor impact on the pKa of a buffer solution. Higher ionic strength and changes in solvent polarity can cause slight variations in the pKa, but these effects are generally small compared to other factors like temperature and the concentrations of buffer components in buffer systems.

Proteins (e.g., enzymes, building blocks of cells, many hormones, etc.) contain side chains that interact to form secondary and tertiary structures that play a crucial role in the proper functioning of proteins. These interactions are categorized as ionic bonding, hydrogen bonding, disulfide linkages, and dispersion forces. Ionic bonds (e.g., salt bridges) and hydrogen bonding are sensitive to pH change. Therefore, maintaining a stable pH is vital for proteins to fold and function properly.

Nearly all cellular processes require a certain level of pH for optimal functioning.

 

Cells exchange materials (e.g., nutrients, waste, signaling molecules, etc.) with their environment through the plasma membrane comprised of a lipid bilayer. Changes in pH disrupt the fluidity and, thereby, the functioning of the membrane.

 

Proteins (e.g., enzymes) are the workhorses responsible for fundamental cellular processes, including DNA replication and transcription, RNA synthesis and translation, ATP production, etc. Dramatic changes in pH may render proteins dysfunctional, leading to cell death.

 

Other processes and structures influenced by changes in pH are the polymerization of the cytoskeleton, energy metabolism, and functioning of organelles, among others.

Most mammalian cells optimally grow at a pH of 7.4, with little variability. While most bacteria grow at neutral (pH = 7), some prefer acidic (pH<3) or alkaline (pH>9) culture conditions. Most plant cells favor slightly acidic culture conditions (pH = 5.8), yet they can tolerate a wide range (pH between 4.0 and 7.2).

As the cells grow, they consume the nutrients from the culture media and produce waste metabolites and respiration, leading to acidification of the culture media.

 

Effective pH control in bacterial cell culture involves using buffer solutions specific to the desired pH range and employing pH monitoring instruments. You can adjust the pH by adding acid or base solutions to maintain it within the target range.

 

pH changes during bacterial growth are influenced by factors like initial pH, culture medium type, and carbon source. For instance, using glucose, glycerol, or octanoate as a carbon source typically lowers pHs, while more oxidized carbon sources like citrate, 2-furoate, 2-oxoglutarate, and fumarate tend to increase medium pH.2 Understanding these factors helps predict pH changes and sheds light on bacterial interactions and fitness in their environment.

 

An excellent open-access publication by Michl et al. offers a detailed flow chart that outlines a step-by-step process for media preparation at a desired pH. Additionally, they provide four specific recommendations aimed at achieving precise pH control in cell culture and enhancing reproducibility.

 

Sources:

Sanchez-Clemente, R., Guijo, M. I., Nogales, J., & Blasco, R. (2020)

Genes 11 (1292)

Good’s buffers are an extensive list of amino acid derivative zwitterionic buffers developed by Norman Good and colleagues from 1966 to 1980. They are the most widely used buffers in biological and biochemical research because they are stable, resist degradation, and tend to not interfere with experimental processes or measurements. They have several characteristics that are crucial in biochemical experiments:

 

·    A pKa value between 6 and 8 (the optimal pH range for most biological reactions.)

·    High-water solubility and minimal solubility in organic solvents.

·    Do not permeate through the cell membrane.  

·    Minimal salt effects.  

·    Minimal influence on buffer concentration, temperature, and ionic composition.

·    Are chemically stable and biochemically inert.

·    Do not form insoluble cationic complexes.

·    Do not absorb visible or UV longer than 230nm.

·    Are easily prepared from affordable materials.

 

Commonly used Good biological buffers:

 

·    HEPES (4-(2-Hydroxyethyl)-1-piperazineethanesulfonic acid)

·    Tris (Tris(hydroxymethyl)aminomethane)

·    MOPS (3-(N-Morpholino)propanesulfonic acid) 

·    MES (2-(N-Morpholino)ethanesulfonic acid)

·    PIPES (Piperazine-N,N'-bis(2-ethanesulfonic acid))

Temperature changes can affect the pKa of buffers. An increase in water temperature leads to the dissociation of H+ concentration and a decrease in the pH of most aqueous biological buffers. The pKa of biological buffers decreases as their concentration and temperature increase.  Another point to remember is that the buffering capacity of biological buffers decreases as the temperature increases. Therefore, measuring the pH of buffers at working concentrations and temperatures is critical for optimal buffering.

 

Fluctuation in pH varies by buffer type. For example, the pKa of buffers containing amino groups tends to decrease as temperature increases, while carboxylic acid moieties are less affected by temperature changes. As you can see below, while the pH of Tris buffers is highly sensitive to temperature change, the pH of MOPS buffers is resistant to temperature change.

 

 

Buffer pKa variations with temperature can be predicted accurately based on changes in enthalpy and heat capacity for the ionization reactions of the buffer components. Specifically, you can use the following equation:

 

ΔpKa = -ΔH / (RT^2)

 

Where:

·    Δ pKa is the change in pKa with temperature. 

·    ΔH is the change in enthalpy (heat content) associated with the ionization reaction.

·    R is the gas constant. 

·    T is the absolute temperature in Kelvin.

 

You can also use this resource to calculate how concentration and temperature change the pH of the commonly used buffer.

Our off-the-shelf buffers and reagents come in a wide variety of pH ranges. We can also easily create custom formulations. Talk to our team to see if we can manufacture whatever you are using, or making, in-house. We’re here to help you find the best solutions for your specific application.

We follow a rigorous quality testing procedure for every lot to ensure variability is less than 0.04 pH from batch to batch.

 

To correctly measure pH value, we carefully establish thermal equilibrium within a controlled environment, in which the product reaches a stable and uniform temperature of 25°C, laying the foundation for precise and consistent pH adjustments.

 

We meticulously manage this controlled environment, relying on precise temperature control equipment to maintain the desired temperature setting. Once the product has fully equilibrated at 25°C, we proceed with the subsequent steps to fine-tune its pH to the desired value, thereby ensuring the accuracy and reliability of our production processes.