Guide for optimal pH

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 pK_a. The pK_a value of buffers can be influenced by temperature, pressure, ionic strength, and the polarity of the solvent.

pK_a 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 pK_a value means it is easier for the acid to donate its H+, making it a strong acid. Conversely, a higher pK_a value means the acid holds onto its H+ tightly (harder to donate), making it a weak acid.

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

pH = pKa + log10 ([A–]/[HA])

How do I use pK_a 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 pK_a—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 pK_a is a unit above or below the desired pH value would provide optimal buffering (i.e., pK_a = desired pH ± 1). For example, if we want to prepare a solution with a pH of 5.0, we’d choose a buffer with pK_a 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 pK_a = 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 pK_a = 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 pK_a?

The ionic strength and polarity of the solvent can have a minor impact on the pK_a of a buffer solution. Higher ionic strength and changes in solvent polarity can cause slight variations in the pK_a, 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.

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 pK_a 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 pK_a 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 pK_a 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 pK_a 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 pK_a 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 pK_a is the change in pK_a 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.