# Lecture 2: Water, pH, and Buffers in Biological Systems
## Biochemistry
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## Learning Objectives
By the end of this lecture, students will be able to:
1. Describe the unique properties of water that make it essential for life
2. Define pH and calculate the pH of solutions
3. Explain the concept of Ka and pKa for weak acids
4. Apply the Henderson-Hasselbalch equation to buffer problems
5. Describe how biological buffer systems maintain physiological pH
6. Explain the clinical significance of acid-base disturbances
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## Lecture Content
### I. The Structure and Properties of Water
Water is the most abundant molecule in living organisms, constituting 60 to 70% of body mass. Its molecular geometry features a bent shape with a bond angle of 104.5 degrees. Because oxygen is more electronegative than hydrogen, water is a polar molecule with a permanent dipole moment. Each water molecule can form up to four hydrogen bonds -- two as a donor through its hydrogen atoms and two as an acceptor through the lone pairs on oxygen. These hydrogen bonds are dynamic, with an average lifetime of approximately 10 picoseconds in liquid water. In ice, water molecules form a more ordered and less dense hydrogen bonding network, which is why ice floats.
#### Key Properties of Water Relevant to Biochemistry
Water has a **high specific heat capacity** (4.18 J/g/K), meaning it resists temperature changes and acts as a thermal buffer for organisms. Its **high heat of vaporization** (2260 J/g) allows effective cooling through evaporation, as seen in sweating. Water's **high dielectric constant** (approximately 80) makes it an excellent solvent for ionic and polar compounds by reducing the electrostatic force between ions by roughly 80-fold compared to vacuum. Its extensive hydrogen bonding network gives rise to significant **cohesion and surface tension**. Finally, water is **amphoteric**, meaning it can act as both an acid and a base.
A detailed molecular diagram of water's hydrogen bonding network. Panel A: A single water molecule showing the bent geometry, bond angle (104.5 degrees), partial charges (delta+ on H, delta- on O), and the direction of the dipole moment. Panel B: A cluster of five water molecules showing hydrogen bonds as dashed lines, illustrating how each molecule can participate in up to four hydrogen bonds. Panel C: Comparison of hydrogen bonding in liquid water (disordered) versus ice (ordered hexagonal lattice), explaining why ice is less dense than liquid water.
### II. Water as a Solvent
Hydrophilic substances dissolve readily in water. Ionic compounds are solvated by water through ion-dipole interactions, while polar molecules interact via hydrogen bonds and dipole-dipole forces. Hydrophobic substances, such as nonpolar molecules like hydrocarbons and lipids, do not dissolve in water. When nonpolar molecules are introduced into an aqueous environment, water molecules form ordered "cages" called clathrates around them. This ordering decreases entropy, which is thermodynamically unfavorable, and the resulting hydrophobic effect drives nonpolar molecules together to minimize the ordered water interface.
Amphipathic (or amphiphilic) molecules contain both polar and nonpolar regions. Examples include fatty acids, phospholipids, and detergents. In aqueous solution, these molecules form micelles, bilayers, or monolayers, and this behavior is the basis for biological membrane formation.
### III. Ionization of Water
Water undergoes autoionization, a process that can be written in simplified form as H2O reversibly dissociating into H+ and OH-. The equilibrium constant for this reaction gives rise to the ion product of water: **Kw = [H+][OH-] = 1.0 x 10^-14 at 25 degrees C**. Since the concentration of water is essentially constant at 55.5 M, this value is treated as a constant. In pure water, [H+] equals [OH-], and both are 1.0 x 10^-7 M.
### IV. pH: A Measure of Acidity
The pH is defined as **pH = -log[H+]**. On the pH scale, a value of 7.0 is neutral at 25 degrees C, values below 7.0 are acidic (indicating excess H+), and values above 7.0 are basic or alkaline (indicating excess OH-). The complementary scale, pOH = -log[OH-], satisfies the relationship pH + pOH = 14 at 25 degrees C.
Different physiological compartments maintain distinct pH values. Blood plasma is tightly regulated between 7.35 and 7.45. Gastric juice is highly acidic at pH 1.0 to 2.0. The lysosomal interior operates at approximately pH 4.5 to 5.0. The cytoplasm is near neutral at about pH 7.2, and pancreatic secretions are mildly alkaline at approximately pH 8.0. Even small changes in pH can dramatically affect protein structure and enzyme activity, underscoring the critical importance of pH regulation.
### V. Acids, Bases, and Conjugate Pairs
Under the **Bronsted-Lowry definition**, an acid is a proton donor and a base is a proton acceptor. Strong acids and bases dissociate completely in water -- for example, HCl fully dissociates into H+ and Cl-, and NaOH fully dissociates into Na+ and OH-. Weak acids and bases, by contrast, partially dissociate and establish an equilibrium between the protonated form (HA) and the deprotonated conjugate base (A-).
The **dissociation constant (Ka)** is defined as Ka = [H+][A-]/[HA]. A larger Ka indicates a stronger acid with greater dissociation. The **pKa**, defined as -log(Ka), provides a more convenient scale: a lower pKa corresponds to a stronger acid. Importantly, the pKa is the pH at which the acid is exactly half dissociated, meaning [HA] equals [A-].
### VI. The Henderson-Hasselbalch Equation
Derived from the Ka expression, the Henderson-Hasselbalch equation is **pH = pKa + log([A-]/[HA])**. This equation relates pH, pKa, and the ratio of conjugate base to acid. When [A-] equals [HA], the pH equals the pKa, which defines the half-equivalence point. When [A-] is greater than [HA], the pH is above the pKa, and when [A-] is less than [HA], the pH is below the pKa.
This equation is invaluable for predicting the ionization state of molecules at a given pH, calculating the pH of buffer solutions, and determining the charge on amino acid side chains at physiological pH.
A titration curve for a weak acid (such as acetic acid, pKa = 4.76). The x-axis shows equivalents of OH- added, the y-axis shows pH. Key features labeled: starting pH, buffer region (shaded, centered on pKa +/- 1 pH unit), half-equivalence point where pH = pKa and [HA] = [A-], equivalence point. An inset shows the Henderson-Hasselbalch equation with arrows pointing to where each variable corresponds on the curve. A second inset shows a bar graph of the relative proportions of HA and A- at pH values of pKa - 2, pKa - 1, pKa, pKa + 1, and pKa + 2.
### VII. Buffers
A buffer is a solution that resists changes in pH upon the addition of small amounts of acid or base. Buffers are composed of a weak acid and its conjugate base (or a weak base and its conjugate acid). **Buffering capacity** is maximal within plus or minus 1 pH unit of the pKa, with the greatest capacity at pH equal to the pKa. Outside this range, the solution cannot effectively resist pH changes.
Buffers work by absorbing added H+ through the reaction A- + H+ forming HA, and by neutralizing added OH- through the reaction HA + OH- forming A- + H2O. Because these reactions shift the ratio [A-]/[HA] only slightly, the pH changes minimally.
### VIII. Biological Buffer Systems
#### Bicarbonate Buffer System (blood)
The bicarbonate buffer system is the most important extracellular buffer. Carbon dioxide reacts with water to form carbonic acid, which then dissociates into H+ and bicarbonate (HCO3-). Although the pKa of carbonic acid (6.1) is seemingly far from the blood pH of 7.4, the system is remarkably effective because it is an open system: CO2 is volatile and can be exhaled by the lungs. At pH 7.4, the [HCO3-]/[CO2] ratio is approximately 20:1. The kidneys regulate bicarbonate concentration while the lungs regulate CO2 levels, providing dual control over blood pH.
#### Phosphate Buffer System
The phosphate buffer system is an important intracellular buffer. The equilibrium between H2PO4- and HPO4^2- has a pKa of 6.86, which is close to the intracellular pH of approximately 7.2, making it effective in the cytoplasm and in urine.
#### Protein Buffer System
Histidine residues, with a pKa of approximately 6.0, serve as effective buffers near physiological pH. Hemoglobin is a major blood protein buffer, with its histidine residues buffering H+ ions produced from CO2 metabolism.
### IX. Clinical Acid-Base Disturbances
Normal blood pH is maintained between 7.35 and 7.45. **Acidosis** (pH below 7.35) can be respiratory, caused by CO2 retention from hypoventilation or COPD, or metabolic, caused by excess H+ production or bicarbonate loss as seen in diabetic ketoacidosis, renal failure, or diarrhea. **Alkalosis** (pH above 7.45) can likewise be respiratory, caused by excessive CO2 loss from hyperventilation, or metabolic, caused by excess bicarbonate or H+ loss as seen with vomiting or diuretic use.
The body employs compensatory mechanisms to counteract these disturbances. Respiratory compensation involves the lungs adjusting CO2 levels and occurs rapidly within minutes. Renal compensation involves the kidneys adjusting bicarbonate and H+ excretion and occurs more slowly over hours to days.
A clinical diagram showing the four types of acid-base disturbances. Central panel: a balance beam showing normal pH with CO2 on one side and HCO3- on the other. Four surrounding panels show: respiratory acidosis (elevated CO2 tips balance), metabolic acidosis (decreased HCO3- tips balance), respiratory alkalosis (decreased CO2), and metabolic alkalosis (increased HCO3-). Each panel includes the primary disturbance, the expected compensatory response, and a common clinical cause.
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