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Water Structure, pH,
and Buffer Systems

Explore the molecular properties of water, understand the pH scale, and learn how buffers maintain biological homeostasis.

Biochemistry

Why Water Is the Solvent of Life

Water makes up 60-75% of human body mass and is the medium for virtually every biochemical reaction. Its unique physical and chemical properties arise from one simple fact: the water molecule is polar.

The Structure of Water

O H H 104.5° δ− δ+ δ+ Net dipole

Hydrogen Bonds and Water's Properties

High Specific Heat
Many H-bonds must be broken to raise temperature, stabilizing body temperature.
High Heat of Vaporization
Evaporative cooling is effective because breaking H-bonds requires significant energy.
Cohesion and Adhesion
H-bonds hold water molecules together (cohesion) and to other polar surfaces (adhesion).

The Autoionization of Water

Water molecules spontaneously transfer protons to one another in a process called autoionization:

The equilibrium constant for this reaction at 25°C is the ion product of water:

Defining \(\text{pH}\)

The pH scale is a logarithmic measure of hydrogen ion concentration:

0 3 7 10 14 Acidic Neutral Basic ← More H⁺ More OH⁻ →

Strong vs Weak Acids and Bases

Strong Weak
Dissociation Complete (100%) Partial (equilibrium)
Example Acid HCl, HNO₃ CH₃COOH, H₂CO₃
Example Base NaOH, KOH NH₃, HCO₃⁻
Equilibrium Constant Very large (\(K_a \gg 1\)) Small (\(K_a \ll 1\))
Buffer Ability Cannot form buffers Can form buffers

The Henderson-Hasselbalch Equation

For a weak acid \(\text{HA}\) in equilibrium with its conjugate base \(\text{A}^-\):

\(\text{p}K_a\)
The pH at which exactly half of the acid is dissociated. \(\text{p}K_a = -\log K_a\)
\([\text{A}^-]/[\text{HA}]\)
The ratio of conjugate base to acid determines how far pH deviates from \(\text{p}K_a\).

How Buffers Resist pH Change

A buffer is an aqueous solution containing a weak acid and its conjugate base (or vice versa). It resists drastic pH changes when small amounts of strong acid or base are added.

Buffer Solution HA + A⁻ Add H⁺ Add OH⁻ A⁻ + H⁺ → HA Conjugate base absorbs acid HA + OH⁻ → A⁻ + H₂O Weak acid absorbs base pH stays stable

Key Biological Buffer Systems

Bicarbonate
\(\text{H}_2\text{CO}_3 / \text{HCO}_3^-\): the primary buffer in blood plasma. Maintains pH near 7.4. \(\text{p}K_a = 6.1\).
Phosphate
\(\text{H}_2\text{PO}_4^- / \text{HPO}_4^{2-}\): important inside cells. \(\text{p}K_a = 6.86\), close to intracellular pH.
Protein
Histidine residues (\(\text{p}K_a \approx 6.0\)) on hemoglobin and other proteins act as intracellular buffers.

Step-by-Step: Calculating pH

1
Identify the species. Is it a strong acid/base (complete dissociation) or weak acid/base (equilibrium)?
2
Find \([\text{H}^+]\) or \([\text{OH}^-]\). For strong acids: \([\text{H}^+] = C_{\text{acid}}\). For weak acids: use \(K_a\) and ICE table.
3
Apply the formula. \(\text{pH} = -\log[\text{H}^+]\). For buffers, use Henderson-Hasselbalch directly.
4
Check your answer. Acidic solutions: pH < 7. Basic solutions: pH > 7. Buffer pH should be near \(\text{p}K_a\).

Example 1: Strong Acid pH

Calculate the pH of a \(0.01\,\text{M}\) HCl solution.

1
Identify: HCl is a strong acid that dissociates completely.
2
Find \([\text{H}^+]\): Since 100% dissociation: \([\text{H}^+] = 0.01\,\text{M} = 1.0 \times 10^{-2}\,\text{M}\)
3
Calculate pH: \(\text{pH} = -\log(1.0 \times 10^{-2}) = 2.0\)

Example 2: Weak Acid pH

Calculate the pH of \(0.1\,\text{M}\) acetic acid (\(K_a = 1.8 \times 10^{-5}\)).

1
Set up equilibrium: \(\text{CH}_3\text{COOH} \rightleftharpoons \text{CH}_3\text{COO}^- + \text{H}^+\)
2
Approximate: \([\text{H}^+] = \sqrt{K_a \cdot C} = \sqrt{(1.8 \times 10^{-5})(0.1)}\)
3
Solve: \([\text{H}^+] = \sqrt{1.8 \times 10^{-6}} = 1.34 \times 10^{-3}\,\text{M}\)
4
pH: \(\text{pH} = -\log(1.34 \times 10^{-3}) \approx 2.87\)

Example 3: Buffer Calculation

A buffer contains \(0.2\,\text{M}\) acetic acid and \(0.15\,\text{M}\) sodium acetate. What is the pH? (\(\text{p}K_a = 4.74\))

The Titration Curve of a Weak Acid

pH Volume of NaOH added (mL) 0 2 4 6 8 10 12 14 ½ eq pH = pKₐ eq pt Equivalence Buffer Region

Common Mistakes to Avoid

✕ Treating Weak as Strong
Setting \([\text{H}^+]\) equal to the initial concentration of a weak acid. Weak acids only partially dissociate: use \(K_a\).
✕ Forgetting the Log
pH is a logarithmic scale. A pH change of 1 unit = a 10-fold change in \([\text{H}^+]\), not a simple addition.
✕ Swapping Ratio
In Henderson-Hasselbalch, it is \(\frac{[\text{A}^-]}{[\text{HA}]}\), not \(\frac{[\text{HA}]}{[\text{A}^-]}\). Swapping gives the wrong sign on the log term.

Summary of Key Equations

Concept Equation When to Use
pH definition \(\text{pH} = -\log[\text{H}^+]\) Always
pOH relation \(\text{pH} + \text{pOH} = 14\) At 25°C
Ion product \(K_w = 1.0 \times 10^{-14}\) At 25°C
Weak acid \([\text{H}^+] \approx \sqrt{K_a \cdot C}\) Dilute weak acid
Buffer pH \(\text{pH} = \text{p}K_a + \log\frac{[\text{A}^-]}{[\text{HA}]}\) Buffer systems

Question 1 of 5

What is the pH of a \(0.001\,\text{M}\) HNO₃ solution?

Question 2 of 5

Which pair can function as a buffer system?

Question 3 of 5

At the half-equivalence point of a weak acid titration, which statement is true?

Question 4 of 5

A buffer has \(\text{p}K_a = 4.74\) and contains equal concentrations of weak acid and conjugate base. What is the pH?

Question 5 of 5

Water has a \(K_w\) of \(1.0 \times 10^{-14}\) at 25°C. If \([\text{OH}^-] = 1.0 \times 10^{-4}\,\text{M}\), what is the pH?