2.4: Buffering against pH Changes in Biological Systems

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Learning Goals

• Understand pH Homeostasis in Physiology:

  • Describe how humans maintain blood pH within a narrow range (7.35–7.45) and explain the consequences of deviations (acidosis and alkalosis).
  • Compare pH environments in different cellular compartments (e.g., lysosomes at pH ~4.5 vs. blood).

• Apply the Henderson–Hasselbalch Equation to Buffer Systems:

  • Derive and use the Henderson–Hasselbalch equation to predict the ratio of conjugate base to weak acid in buffering systems.
  • Explain why buffering capacity is maximal when pH equals pKa and how small changes in [A⁻] or [HA] impact the pH.

• Analyze the Carbonic Acid/Bicarbonate Buffering System:

  • Describe the sequential reactions involving CO₂, H₂CO₃, H₃O⁺, and HCO₃⁻ that underlie the carbonic acid/bicarbonate buffering system.
  • Calculate the effective pKa for the system and interpret how the CO₂-to-bicarbonate ratio primes the buffer for rapid response to metabolic changes.
  • Explain how the respiratory system and kidneys work in tandem to regulate this buffer system and, hence, blood pH.

• Evaluate Other Biological Buffering Agents:

  • Examine the roles of the phosphate buffering system and protein buffers in maintaining intracellular and extracellular pH.
  • Discuss why certain buffering systems (like the phosphate pair) are less influential in blood compared to the carbonic acid system.

• Develop Laboratory Buffer Preparation Skills:

  • Identify strategies for preparing buffered solutions in the lab (mixing separate acid and conjugate base solutions, titrating with strong acid/base, and adjusting volumes).
  • Understand the importance of selecting appropriate buffers (e.g., dihydrogen phosphate/monohydrogen phosphate, HEPES, Tris) based on their pKa values and potential interactions (e.g., ion binding).

• Link Buffer Systems to Broader Physiological and Environmental Contexts:

  • Discuss how imbalances in buffering systems can lead to clinical conditions like respiratory acidosis or alkalosis, and explain the physiological responses to such imbalances.
  • Evaluate the impact of increased atmospheric CO₂ on ocean pH and consider how buffering reactions relate to carbon capture strategies in climate change mitigation.

These goals encourage a comprehensive understanding of both the quantitative and qualitative aspects of biological buffers, linking fundamental chemistry with physiological regulation and broader environmental challenges.

Introduction

To maintain homeostasis, humans regulate their pH levels between 7.35 and 7.45. (Much lower pH values, ≈ 4.5, are found in the lysosome.) Lower pH values are associated with metabolic and respiratory acidosis, while higher pH values are characteristic of metabolic and respiratory alkalosis. pH is maintained by buffering systems that consist of a weak acid and a base. If you understand the Henderson-Hasselbalch equation from the previous section, buffer systems become easy to understand.

\begin{equation}
\mathrm{pH}=\mathrm{pK}_{\mathrm{a}}+\log \frac{\left[\mathrm{A}^{-}\right]}{\lceil\mathrm{HA}\rceil}
\end{equation}

At the curve's inflection point, pH = pKa; at this pH, the system is most resistant to changes in pH when adding either acid or base. At this pH, [HA]=[A^-].

If a bit of a strong acid is added, it will react with the strongest base in the solution, which would be the conjugate base of the weak acid:

HCl + A^- → HA + Cl^-

The reaction goes from a strong acid, HCl, to a weak acid, HA. Its concentration would increase slightly, but it will only ionize to a small extent since it's a weak acid. The [HA] in the Henderson-Hasselbalch equation increases slightly, but not enough to significantly change the pH. If the same amount of HCl were added to pure water, it would react completely, forming an equal amount of H₃O⁺, significantly altering the pH of pure water (7.0).

If a bit of a strong base is added, it will react with the strongest acid in the solution, which would be HA:

HA + OH^- → H₂O + A^-

The reaction goes from a strong base to a weak base A^-. Its concentration would increase slightly, but it wouldn't significantly affect the pH since it's a weak base. The [A^-] in the Henderson-Hasselbalch equations increases slightly but not enough to change the pH significantly. If the same amount of NaOH were added to pure water, it would react, making the solution basic and significantly altering its pH (7.0).

To review, buffer solutions contain a weak acid and its conjugate base. They have a maximal buffering capacity at a pH equal to the weak acid's pKa. Generally, a buffered solution can best withstand a change in pH of only + 1 pH unit from the pKa.

Biological Buffering Agents

The most relevant biological systems are the carbonic acid/carbonate buffering system, which controls blood and cell pH, and the phosphate buffering system. Proteins with many weak acid and base functional groups can also act as buffering agents.

Carbonic acid/carbonate buffering system: At first glance, the reaction of carbonic acid, H₂CO₃, with water can be written as follows:

\begin{equation}
\mathrm{H}_2 \mathrm{CO}_3(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \leftrightarrow \mathrm{H}_3 \mathrm{O}^{+}(\mathrm{aq})+\mathrm{HCO}_3^{-}(\mathrm{aq}) \quad \mathrm{pKa}=3.6
\end{equation}

where H₂CO₃ (carbonic acid) is the weak oxyacid, and HCO₃^-(aq) (bicarbonate aka hydrogen carbonate) is its weak conjugate base.

However, this system is more complex since we must add to it another reaction for the formation of H₂CO₃ (aq) in the blood:

\begin{equation}
\mathrm{CO}_2(\mathrm{~g}) \leftrightarrow \mathrm{CO}_2(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \leftrightarrow \mathrm{H}_2 \mathrm{CO}_3(\mathrm{aq})
\end{equation}

The [CO₂(aq)] >> [H₂CO₃ (aq)] and their ratio is around 340/1.  This makes sense since CO₂ is a very stable molecule. CO₂ in the aqueous form can be readily transported through the blood.  Combine the reactions to give the equation below:

\begin{equation}
\mathrm{CO}_2(\mathrm{~g}) \leftrightarrow \mathrm{CO}_2(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \leftrightarrow \mathrm{H}_2 \mathrm{CO}_3(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \leftrightarrow \mathrm{H}_3 \mathrm{O}^{+}(\mathrm{aq})+\mathrm{HCO}_3^{-}(\mathrm{aq})
\end{equation}

How can carbonic acid, with a pKa of 3.6, buffer an aqueous solution at pH 7.5 in the blood and cells?  An astute student might have picked up this conundrum. The solution to this problem involves re-examining the complete set of reactions for the components of the buffer system.  Let's simplify Equation 2.3.4 since there would be no free "gas bubbles" in blood, so CO₂ (g) = CO₂(aq):

\begin{equation}
\mathrm{CO}_2(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \leftrightarrow \mathrm{H}_2 \mathrm{CO}_3(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \leftrightarrow \mathrm{H}_3 \mathrm{O}^{+}(\mathrm{aq})+\mathrm{HCO}_3^{-}(\mathrm{aq})
\end{equation}

 

H₂CO_{3(aq) participates in two distinct} reactions.

Rightwards from H₂CO₃ (aq) :

\begin{equation}
\mathrm{H}_2 \mathrm{CO}_3(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \leftrightarrow \mathrm{H}_3 \mathrm{O}^{+}(\mathrm{aq})+\mathrm{HCO}_3^{-}(\mathrm{aq})
\end{equation}

Using the simplified equation with H⁺ gives

\begin{equation}
\mathrm{K}_{\mathrm{a}}=\frac{\left[\mathrm{H}^{+}\right]\left[\mathrm{HCO}_3^{-}\right]}{\left[\mathrm{H}_2 \mathrm{CO}_3\right]}
\end{equation}

Hence,

\begin{equation}
\left[\mathrm{H}_2 \mathrm{CO}_3\right]=\frac{\left[\mathrm{H}^{+}\right]\left[\mathrm{HCO}_3^{-}\right]}{K_a}
\end{equation}

Leftwards from H₂CO₃ (aq) :

\begin{equation}
\mathrm{H}_2 \mathrm{CO}_3(\mathrm{aq}) \leftrightarrow \mathrm{CO}_2(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})
\end{equation}

\begin{equation}
\mathrm{K}_2=\frac{\left[\mathrm{CO}_2\right]}{\left[\mathrm{H}_2 \mathrm{CO}_3\right]}
\end{equation}

so

\begin{equation}
\left[\mathrm{H}_2 \mathrm{CO}_3\right]=\frac{\left[\mathrm{CO}_2\right]}{K_2}
\end{equation}

Since there can be only 1 H₂CO₃ concentration, set Equations  2.3.8 and 2.3.11 equal to each:

\begin{equation}
\left[\mathrm{H}_2 \mathrm{CO}_3\right]=\frac{\left[\mathrm{H}^{+}\right]\left[\mathrm{HCO}_3^{-}\right]}{K_a}=\frac{\left[\mathrm{CO}_2\right]}{K_2}
\end{equation}

Solving for [H⁺] gives:

\begin{equation}
\left[H^{+}\right]=\frac{\left[\mathrm{CO}_2\right]\left(K_a\right)}{\left[H C O_3^{-}\right]\left(K_2\right)}
\end{equation}

Now take the -log of each side to produce an equation similar to the Henderson-Hasselbalch equation.

\begin{equation}
-\log \left[\mathrm{H}^{+}\right]=-\log \left(\frac{\left[\mathrm{CO}_2\right]}{\left[\mathrm{HCO}_3^{-}\right]}\right)-\log \left(\frac{\mathrm{K}_{\mathrm{a}}}{\mathrm{K}_2}\right)
\end{equation}

where

\begin{equation}
K_{a E F F E C T I V E}=\frac{K_a}{K_2}
\end{equation}

This Henderson-Hasselbalch-like equation indicates that pH is determined by the \(K_a/K_2\) ratio. pK_{a EFFECTIVE} = 6.3. This gives a ratio of \(CO_2/HCO_3^{-}\) of 0.08 = 8/100. There is effectively 12-13 times as much HCO₃^-(aq) as CO₂, making the system primed to react with acid produced metabolically. Yet a second conundrum exists. The pH of the blood (7.4) is outside the optimal range for a buffer system (in this case, +1 pH unit from the pKa, which is 6.3). Again, the system is primed to react with acid, as this would move the pH closer to the optimal buffering pH of 6.3. Other biological systems also must be involved in maintaining pH.

The respiratory system can quickly adjust pH by increasing CO₂ exhalation. The kidneys can respond more slowly to remove H₃O⁺ and retain HCO₃^-. The carbonic acid/bicarbonate buffering system can help us understand how shifting equilibria caused by excessive CO₂ released during rapid, deep breathing or decreased CO₂ release associated with pulmonary disease or shallow, rapid breathing can lead to respiratory alkalosis and acidosis, respectively.

• Respiratory alkalosis can be caused by “hyperventilation” or breathing rapidly. This can lead to breathing out (removing) too much CO₂, shifting the above equilibrium to the left, consuming H₃O⁺, and increasing pH, making the blood more alkaline. To increase your CO₂ levels, you could breathe into a bag.
• Respiratory acidosis is caused by increased CO₂, which can occur when the lungs aren’t working well and you can’t get rid of the CO2 you produce during respiration. Respiratory acidosis can occur with asthma, pneumonia, lung disease, or any condition that decreases the respiratory rate.

Inhaling CO₂ can lead to panic. This makes sense as it would mimic suffocation, which is lethal to humans. A suffocation response follows. High CO₂ would drive the equilibrium to the right, leading to H₃O⁺ production. An acid-sensing ion channel-1a (ASIC1a) in the amygdala, a center of emotion regulation in the brain, has been identified and appears to mediate panic. Panic attacks are sometimes associated with hyperventilation, which leads to alkalosis, not acidosis. Less noted is that when some people panic, they take short, shallow breaths (in a way, almost stopping their breath). This would lead to a buildup of CO₂, as it wouldn’t be released during exhalation. The acid channel in the amygdala would be activated, and a panic response would ensue.

Phosphate buffering system: Phosphates, specifically dihydrogen (H₂PO₄^-) and monohydrogen phosphate (HPO₄^2-), are also present in the blood. Given the pKa of HPO₄^2-, why is PO₄^3- not present to any significant degree? Since the concentration of phosphates in blood is low, this system plays a minor role in blood.

Proteins: Proteins are found in all cellular and extracellular fluids and, with their weak acid substituents, act as buffer components. Proteins contain two amino acids, aspartic acid and glutamic acid, that contain carboxylic acid side chains. Each comprises about 6% of the proteins. In blood, hemoglobin is the most abundant protein by far. Its role in buffering and in O₂ and CO₂ will be discussed in a subsequent chapter.

Making Buffers in the Lab

When studying biomolecules, such as proteins and nucleic acids, in the lab, the pH of the solution is typically maintained under physiological conditions. These macromolecules are either dissolved in or diluted into a buffer solution. Sometimes, it's essential to study their properties and activities as a function of pH. A wide variety of buffer systems have been developed for the lab study of these molecules. The dihydrogen phosphate (H₂PO₄^-)/monohydrogen phosphate (HPO₄^2-) pair is commonly used as the pKa of H₂PO₄^- is 7.21, making them physiologically relevant. Care must be taken when selecting buffer systems, as some might bind calcium ions. The pKa of some weak acids also varies significantly with temperature. The following are some standard biological buffers.

Buffers	pKa (at 25°C)
MES	6.10
Bis-Tris	6.50
ACES	6.78
PIPES	6.76
MOPSO	6.90
MOPS	7.20
HEPES	7.48
Tris	8.06
Tricine	8.05
Gly-Gly	8.20
Bicine	8.26
TAPS	8.40
AMPSO	9.00
CAPS	10.40

There are three general ways to make a buffered solution:

1. Make separate and equal concentration solutions for both the weak acid (for example, NaH₂PO₄) and its conjugate base (for example, Na₂HPO₄₎. Use the Henderson-Hasselbalch equation to calculate how much of each should be added to give the correct [A^-]/[HA] ratio (in the case [HPO₄^2-]/[H₂PO₄^-1]) to give the correct pH.
2. Use a pH meter to monitor the pH as you add both solutions together until the desired pH is achieved.
3. Prepare a solution of one of the components (weak acid or its conjugate base) to the correct volume for the desired molarity. Monitor the pH as you add a concentrated solution of either HCl or NaOH to get the desired pH. Then, bring the solution to the correct volume in a volumetric flask. With this method, you will add counter ions (Cl^- with HCl and Na⁺ with NaOH), which you may not want in your buffer solution. Often, it is not a problem.

Climate Change and Carbon Capture 

Perhaps humanity's greatest challenge is the effects of global warming and climate change on the biosphere and its health. CO₂  in the atmosphere, HCO₃^-(aq), its soluble form, and CO₃^2-, which is mainly precipitated as insoluble carbonate, are key players in the Carbon Cycle and climate change.  Follow the link below to see how increasing CO₂ in the atmosphere due to the production and use of fossil fuels affects the pH of the oceans and their health, and how CO₂ captured by oceans could be used to decrease atmospheric CO₂ and help stop and reverse global warming. 

32.3: Climate Change - Part 1 - The Carbon Cycle and Carbon Chemistry

Summary

Chapter Summary

This chapter explores the critical role of buffering systems in maintaining pH homeostasis within the human body and in laboratory settings. It begins by discussing how blood pH is tightly regulated between 7.35 and 7.45, with deviations leading to metabolic or respiratory acidosis and alkalosis. The chapter emphasizes that this regulation is achieved through buffer systems—combinations of weak acids and their conjugate bases—that resist significant pH changes when small amounts of strong acids or bases are added.

A central tool introduced is the Henderson–Hasselbalch equation, which quantitatively relates pH, pKa, and the ratio of the concentrations of the conjugate base and the weak acid. The equation explains the buffering capacity at pH = pKa and underpins the behavior of titration curves and the concept of buffering ranges.

The chapter then focuses on the carbonic acid/bicarbonate buffering system, a primary mechanism for pH regulation in blood and cells. Through a series of interrelated reactions involving CO₂, H₂CO₃, H₃O⁺, and HCO₃⁻, the system is shown to rapidly adjust pH. The effective pKa of the system, derived from the ratio of CO₂ and bicarbonate, explains how the buffer is primed to neutralize metabolic acids. At the same time, respiratory and renal adjustments fine-tune the pH balance.

Additional biological buffers, including the phosphate system and proteins, are discussed. Although the phosphate buffer plays a minor role in blood due to its low concentration, it is vital in other cellular contexts. Proteins, especially abundant ones like hemoglobin, also contribute to buffering through their ionizable side chains.

Finally, the chapter provides practical insights into buffer preparation in the laboratory. Various strategies for creating buffered solutions are presented, with attention to selecting buffers based on their pKa values and potential interactions with other ions. The discussion concludes with a broader perspective on how buffer chemistry connects to global issues, such as the role of CO₂ in climate change and carbon capture efforts.

The chapter integrates fundamental chemical principles with biological applications, enabling students to understand and predict how buffer systems stabilize pH in physiological and experimental contexts.