pH and Buffering

• pH and buffering presented by Despo Papachristodoulou.

Learning Outcomes

• Explain the difference between strong and weak acids in terms of dissociation properties.
• Define and explain pKa and its importance in biological buffering.
• Use the Henderson-Hasselbalch equation.
• List a number of physiological buffers.
• Explain how proteins can be used as physiological buffers.
• Describe the function of Hemoglobin as a buffer for hydrogen ions produced in metabolism; explain what allows this to happen.

What is pH?

• pH measures hydrogen ion concentration, indicating the acidity or alkalinity of a solution.
• Acidity depends only on free hydrogen ions, not those bound to anions.

Blood pH

• Blood pH regulation is critical as blood contacts nearly every body cell.
• Normal blood pH range is very narrow: 7.35 to 7.45; pH levels outside these limits may be fatal.
• The living pH range extends from 7.0 to 7.8, with conditions including acidosis, normal range, and alkalosis.

Acids in the body

• Acids enter the body through foods.
• Acids are generated via breakdown of proteins, incomplete oxidation of fats or glucose, and loading and transport of carbon dioxide in the blood.
• Acid-base balance is regulated by the lungs, the kidneys, and chemical buffers.

Buffering

• Buffers resist abrupt and large pH swings in body fluids
• Buffers release H+ (acting as acids) when pH rises and bind H+ (acting as bases) when pH drops.

How buffers operate?

• Thorough understanding of a strong acid, strong base, weak acid, weak base, and ionization of water is important.

Ionization of water

• Pure water is a 55.6M solution
• Water dissociates to a very small extent: H₂O ↔ H+ + OH-
• [H+] x [OH-] = 10-14 M², which is the ionic product of water and the neutrality, where [H+] = [OH-] = 10-7M

pH Calculation

• pH = -log[H+]
• When [H+] = 10-2M, pH = 2
• At neutrality, where [H+] = [OH-], [H+] = 10-7M, so pH = -log[10-7] = -(-7) = 7
• When [H+] is 10-2, [OH-] is 10-12, and when [H+] is 10-4, [OH-] is 10-10
• Logarithmic values: 10^1 = 10 (log = 1), 10^2 = 100 (log = 2), 10^-2 = 1/100 (log = -2)
• Blood pH is 7.4
• pH calculation: 7.4 = -log[H+]; therefore, [H+] = 3.98 x 10^-8 M

Dissociation of Acids

• Acids are proton (H+) donors, while bases are proton acceptors.
• Strong acids dissociate completely whereas weak acids dissociate incompletely, depending on pH.
• Strong bases are more effective proton acceptors than weak ones.

pKa Explanation

• pKa equals -log Ka, with Ka representing the acid dissociation constant.
• It defines the pH at which an acid is half-dissociated and has equal amounts of undissociated acid and its conjugate base.
• A lower pKa signifies a stronger acid.
• At the pKA: Equal amounts of conjugate base and acid exist, essential for a mixture of weak acids and their conjugate bases to buffer solutions
• Buffering is the ability of a solution to resist pH change upon acid or alkali addition.
• Buffering is most effective at the pKa, where dissociated and non-dissociated forms of the acid (conjugate base and acid) are equal.

How pKa relates to buffering

• If H+ ions are added, conjugate bases pick them up.
• If OH- ions are added, acids donate protons to form H₂O.
• pH remains stable.
• Knowing the pKa of an acid indicates its best buffering pH.

Henderson-Hasselbalch equation

• pH = pKa + log([conjugate base]/[acid])
• Relates pH of a solution to pKa of weak acid and relative amounts of dissociated/non-dissociated forms.

Physiologically Important Buffers

• In blood, saliva, and other body fluids, H₂CO₃↔HCO₃- has a pKa of 6.1.
• H₂PO₄-↔HPO₄²- has a pKa of 6.8.

Henderson-Hasselbalch Equation Application

• (H₂CO₃↔HCO₃-) pKa 6.1
• H₂CO₃ is proportional to pCO₂
• Henderson-Hasselbalch equation may distinguish acidosis causes like metabolic or respiratory issues.

Amino Acids as Buffers

• Amino acids can be used as buffers
• Glycine titration with strong base is the reaction
• Acids present -COOH and -NH3+
• -COOH is the stronger acid, with a pKa of 2.34
• NH3+ is the weaker acid, with a pKa of 9.66
• Best buffering occurs around the two pKa values.
• There is no buffering by the zwitterion.
• Glycine isn't a good candidate for this because it buffers best at pH 2.3 and 9.6.

Haemoglobin's buffering

• Haemoglobin acts as a crucial blood buffer.
• Most amino acid side chains do not buffer within the physiological range.
• A large number of histidine residues contributes to haemoglobin's effectiveness.
• The pKa of histidine in Hb differs from free histidine (6), as neighboring groups affect the pKa.
• Oxyhemoglobin pKa = 6.8, and Deoxyhemoglobin pKa = 7.8.
• Which form of haemoglobin, oxyhemoglobin or deoxyhemoglobin, serves as a better buffer for H+ ions produced in metabolism? The answer depends on which is better able to pick up and carry H ions at the pH of blood. In order to know this, think pKa. What does it mean if the pH of blood is higher or lower than the pKa of Hb? Which form of it would be better able to carry the H ions?
• Hb carries oxygen to cells and carbon dioxide and picks up H ions produced in metabolism
• The different pKas of oxy and deoxy Hb mean that they are protonated to a different degree at the pH of blood

Henderson Hasselbalch equation

• pH of blood is 7.4
• pKa oxyHb 6.8
• pKa deoxyHb 7.8
• OxyHb: 7.4 = 6.8 + 0.6
• Log of base / acid is 0.6
• The ratio of base / acid is 3.98 : 1
• DeoxyHb: 7.4 = 7.8 - 0.4
• Log of base / acid is -0.4
• The ratio of base / acid is 1: 2.51

Free Diffusion & Biological Membranes

• Lipid soluble molecules can diffuse freely across biological membranes
• Substances that interact strongly with water cannot.
• Aspirin diffuses more easily at pH 2 rather than pH 7

Explore pH, strong vs. weak acids, and pKa's role in biological buffering. Use the Henderson-Hasselbalch equation to understand physiological buffers, including proteins and hemoglobin, in maintaining blood pH balance.