CH2 CHEM PT 2 - pH &BUFFERS -

pH Basics

• pH is a measure of acidity in a solution, defined by hydrogen ion concentration. An ion is an atom or molecule with a net electric charge; a hydrogen ion is denoted as H^+.

• The more hydrogen ions present, the greater the acidity; fewer hydrogen ions mean less acidity.

• The pH scale runs from 0 to 14.

  • Neutral on the scale is 7.

  • Higher numbers (above 7) indicate basic (alkaline) conditions, i.e., fewer hydrogen ions.

  • Lower numbers (below 7) indicate acidic conditions, i.e., more hydrogen ions.

• In the body, pH is tightly regulated because small changes in hydrogen ion concentration can have large effects on biological systems.

The pH scale and what it means

• The pH value is a negative base-10 logarithm of the hydrogen ion concentration: \mathrm{pH} = -\log_{10}[H^+].

• Equivalently, the hydrogen ion concentration is: [H^+] = 10^{-\mathrm{pH}} \;\text{M} (molar).

• Examples:

  • At pH = 1: [H^+] = 10^{-1} = 0.1\ \text{M}.

  • At pH = 2: [H^+] = 10^{-2} = 0.01\ \text{M}.

  • At pH = 3: [H^+] = 10^{-3} = 0.001\ \text{M}.

• This is a base-10 (common) logarithmic scale.

• Relationship: for each unit change in pH, the hydrogen ion concentration changes by a factor of 10.

  • If pH goes from 2 to 3, \frac{[H^+]{pH=3}}{[H^+]{pH=2}} = \frac{10^{-3}}{10^{-2}} = 10^{-1} = 0.1, i.e., tenfold decrease in acidity (tenfold increase in basicity).

  • If pH goes from 4 to 3, you gain ten times as many hydrogen ions (tenfold increase in acidity): \frac{[H^+]{pH=3}}{[H^+]{pH=4}} = \frac{10^{-3}}{10^{-4}} = 10.

• Magnitude example mentioned in the transcript:

  • The pH range from 7.35 to 7.0 represents a drop of 0.35 pH units. The actual factor in hydrogen ion concentration is 10^{0.35} \approx 2.24\;\text{(about 2.2x more acidic)}. The transcript states “about five times” in this context, which is a rough/illustrative claim rather than a precise calculation.

• The most acidic solution has pH ≈ 0 ( [H^+] = 1\,\text{M} ), and the most basic has pH ≈ 14 ( [H^+] = 10^{-14}\,\text{M} ).

Blood pH and homeostasis

• Normal human blood pH is roughly in the range: 7.35 \leq \text{pH} \leq 7.45.

• This range is considered slightly basic; neutral is 7, so values above 7 are less acidic.

• Why this matters: small deviations can have large biological consequences because proteins rely on a precise three-dimensional shape.

• If pH deviates too far (e.g., dropping toward 7.0 or rising toward 7.8), the risk of coma or death increases due to disruption of cellular processes and protein function.

Why pH affects proteins

• Proteins have specific three-dimensional structures essential for function.

• Changes in acidity (and temperature) can cause proteins to unfold, a process called denaturation, leading to loss of function.

• Denaturation can be life-threatening because many biological processes depend on properly folded proteins.

• This is a core reason why maintaining blood pH within the narrow range is critical for homeostasis.

Buffers and buffering in the body

• Buffers are systems that resist changes in pH by either adding hydrogen ions (H+) or removing them, depending on which direction the pH is moving.

• A buffer consists of two parts:

  • A weak acid that can donate hydrogen ions (H+).

  • The conjugate base (the base form) that can remove hydrogen ions.

• In the presence of excess base (too few H+), the weak acid component donates H+ to bring pH back down.

• In the presence of excess acid (too many H+), the conjugate base removes H+ to bring pH back up.

• The key idea: buffers can both add and remove hydrogen ions to stabilize pH around the normal range.

The bicarbonate-carbonic acid buffer system (a specific buffer)

• The two components:

  • Carbonic acid, \mathrm{H2CO3}, a weak acid (can donate H+).

  • Bicarbonate, \mathrm{HCO_3^-}, its conjugate base (can remove H+).

• The buffer reaction is a reversible equilibrium:
  \mathrm{H2CO3} \rightleftharpoons \mathrm{H^+} + \mathrm{HCO_3^-}

• How it acts:

  • If pH rises (too basic, not enough H+): carbonic acid can dissociate to release a hydrogen ion, forming bicarbonate:
    \mathrm{H2CO3} \rightarrow \mathrm{H^+} + \mathrm{HCO_3^-}

• This adds hydrogen ions to bring the pH back toward normal.

  • If pH falls (too acidic, too many H+): bicarbonate can bind a hydrogen ion to form carbonic acid, removing H+ from solution:
    \mathrm{HCO3^-} + \mathrm{H^+} \rightarrow \mathrm{H2CO_3}

• Important points:

  • Carbonic acid itself does not directly set the pH; it is the free hydrogen ions that determine pH.

  • The buffer system keeps a balance by having both components present in the bloodstream and at any moment one can respond to pH changes.

  • Buffers exist both inside cells and in the blood; they are essential for maintaining pH during daily activities (e.g., eating, drinking, exercising) and during physiological processes like respiration and fluid/electrolyte balance.

Demonstrations and practical notes

• The instructor uses a simple demonstration with an expo marker to illustrate carbonic acid and bicarbonate interconversion in real time:

  • Conceptual reaction: carbonic acid can dissociate to release a hydrogen ion and bicarbonate; bicarbonate can bind a hydrogen ion to reform carbonic acid.

  • This illustrates how the same two species (\mathrm{H2CO3} and \mathrm{HCO_3^-}) continuously interconvert to maintain pH.

• Key takeaway for physiology: maintaining pH via buffers is crucial for maintaining protein structure, enzyme activity, and overall homeostasis; disruptions are linked to adverse outcomes and can be life-threatening.

Connections and implications

• Foundational principle: pH is a logarithmic measure of hydrogen ion concentration; small numerical changes imply large biological changes.

• Buffer systems, especially the bicarbonate-carbonic acid buffer, are central to respiratory and metabolic regulation of pH.

• The same buffering concepts apply across biology: cells, blood, and extracellular fluid rely on buffers to maintain a stable internal environment.

• Ethical/clinical implications: understanding pH and buffering informs treatment of metabolic or respiratory disorders and informs interventions during vomiting, metabolic acidosis/alkalosis, and severe illness.

Quick recap for exam readiness

• pH = -log10([H+]); [H+] = 10^{-pH} M.

• Each unit change in pH represents a tenfold change in hydrogen ion concentration.

• Normal blood pH: approximately 7.35\leq \text{pH} \leq 7.45.

• Elements of buffering: weak acid + conjugate base; capable of both adding and removing H+.

• Carbonic acid (H2CO3) / bicarbonate (HCO3-) buffer system:

  • \mathrm{H2CO3} \rightleftharpoons \mathrm{H^+} + \mathrm{HCO_3^-}

  • If pH rises: dissociation of H2CO3 adds H+.

  • If pH falls: HCO3- binds H+ to form H2CO3.

• Maintaining pH is essential to prevent protein denaturation and maintain homeostasis; disruptions can lead to serious outcomes.