acids and bases

Context and Objectives

• Focus: acids and bases in the body and environmental context (drinking water).

• Spotlight on acidified drinking water within a population; reference to the Flint water crisis (started ~10 years ago; resolved legally, etc.).

• Purpose: understand how acids and bases affect water chemistry, buffering, and biological systems.

Water in the body and autoionization

• Water autoionization: $ext{H}_2 ext{O} ightleftharpoons ext{H}^+ + ext{OH}^-$

  • Occurs spontaneously but very infrequently in pure water.

  • In pure water, the hydrogen ion concentration equals the hydroxide ion concentration.

  • This balance underpins the neutral pH of pure water.

• Very small fraction of water molecules ionize at any instant (described as a rare event; the transcript notes it as an infrequent reaction, roughly "one in five hundred million" occurrences in water).

• In solution with other species, acids and bases disrupt this balance by altering the relative amounts of
  $[ ext{H}^+]$ and $[ ext{OH}^-]$, changing the pH.

• Relevance to biology: most body water is water, so acid-base chemistry directly affects biological molecules and processes.

Definitions: acids and bases (as introduced in the transcript)

• Acids:

  • By definition, acids added to water donate hydrogen ions to the solution.

  • This increases the hydrogen ion concentration in the solution, shifting the balance toward more $[ ext{H}^+]$.

  • Example given: Hydrogen chloride, $ext{HCl}$, is a strong acid that dissociates in water to increase $[ ext{H}^+]$.

  • Note from the transcript: HCl is described as a molecule secreted by cells; in physiology, HCl is produced in the stomach (as part of gastric juice).

• Bases:

  • Bases dissolved in water decrease the hydrogen ion concentration, either directly by accepting a hydrogen ion or indirectly by increasing the hydroxide ion concentration, which then binds $ext{H}^+$ to form water.

  • Direct base action: molecules that directly accept $ext{H}^+$ (e.g., ammonia, $ext{NH}<em>3$, which becomes $ext{NH}</em>4^+$ when protonated, in the transcript referred to ammonia as accepting a proton).

  • Indirect base action: molecules that increase $[ ext{OH}^-]$ (e.g., sodium hydroxide, $ext{NaOH}$). The $ext{OH}^-$ can pair with a hydrogen ion to form water, reducing the free $[ ext{H}^+]$.

• Summary from transcript:

  • Acids: increase $[ ext{H}^+]$ in solution.

  • Bases: decrease $[ ext{H}^+]$ directly by accepting a proton or indirectly by increasing $[ ext{OH}^-]$, which consumes $ext{H}^+$.

  • The balance of $[ ext{H}^+]$ and $[ ext{OH}^-]$ defines acidity/basicity of the solution.

Examples discussed in the transcript

• Acid example:

  • $ext{HCl} <br>ightarrow ext{H}^+ + ext{Cl}^-$ in water, increasing $[ ext{H}^+]$ (acid).

• Base examples:

  • Sodium hydroxide: $ext{NaOH} <br>ightarrow ext{Na}^+ + ext{OH}^-$, increasing $[ ext{OH}^-]$ which then binds $ext{H}^+$ to form water, reducing $[ ext{H}^+]$ (base effect).

  • Ammonia: a base that can directly bind a hydrogen ion (the transcript mentions ammonia binding a hydrogen, acting as a base).

• Note on bicarbonate (buffer system) discussed below as a key biological buffer.

pH scale and biological relevance

• pH scale basics (as stated):

  • Range: $1 ext{ to } 14$ with $7$ being neutral (balanced $[ ext{H}^+]$ and $[ ext{OH}^-]$) in pure water.

  • Acidic solutions have pH < 7; basic (alkaline) solutions have pH > 7.

• Biological window for blood pH:

  • Normal human blood pH is roughly between $7.35$ and $7.45$.

  • Dropping below $7.35$ leads toward acidosis; rising above $7.45$ leads toward alkalosis; both can impair physiological functions.

  • The transcript highlights that changes in blood pH can disrupt the bonds in hemoglobin that are essential for oxygen binding and transport.

• Stomach acid and protective mechanisms:

  • The body uses protective mechanisms to prevent damage from stomach acid, including mucus lining and bicarbonate buffering.

  • The question about how the stomach prevents it from digesting itself is tied to buffering and protective layers (not memorized in this section but mentioned as context).

Buffer systems in the body: bicarbonate example

• Bicarbonate buffer system:

  • Bicarbonate is denoted as $ext{HCO}_3^-$ and acts as a base by accepting a hydrogen ion.

  • Key reaction (simplified):
    $ext{HCO}<em>3^- + ext{H}^+ ightarrow ext{H}</em>2 ext{CO}_3$

  • This buffering helps maintain pH in extracellular fluid and blood by neutralizing excess acids.

• The transcript emphasizes that bicarbonate represents an important buffer system; cells can produce bicarbonate both inside and outside cells to help neutralize acidity in the environment around cells.

• Wider context: buffering systems analogous to bicarbonate exist in other environments, e.g., ocean waters (transcript notes a cross-domain analogy).

• Practical significance: buffers help maintain pH within narrow ranges required for enzyme activity, receptor function, and gas transport (e.g., oxygen binding by hemoglobin).

Drinking water, acidification, and monitoring

• Most drinking water is neutral or slightly basic under normal circumstances; acidification is undesirable.

• If drinking water becomes too acidic, buffering agents are used in treatment to raise pH and stabilize water chemistry.

• Water monitoring is essential for communities:

  • Regular testing by water suppliers includes checks beyond pathogens (e.g., chemical balance, contaminants, microplastics).

  • Acid-base balance in drinking water is an important safety consideration.

• Flint water crisis reference:

  • Highlights real-world consequences of improper water chemistry and the importance of maintaining safe pH and buffering in drinking water.

Practical implications and physiological connections

• Reflux and acid exposure:

  • Reflux refers to the regurgitation of stomach acid into the esophagus; this can be painful because the esophagus lacks the same protective mechanisms as the stomach.

  • Protective factors in the stomach include mucus and bicarbonate buffering to limit damage from the highly acidic environment.

• Overall takeaway:

  • Understanding acids, bases, pH, and buffering is crucial for interpreting how body fluids maintain homeostasis, how water quality impacts health, and how disruptions (like acidified drinking water) can have systemic effects.

Quick recap of key equations and numbers from the transcript

• Acid-base definitions:

  • Acids increase $[ ext{H}^+]$ in solution.

  • Bases decrease $[ ext{H}^+]$ directly or increase $[ ext{OH}^-]$ indirectly.

• pH expression:

  • $ext{pH} = -<br>\log_{10}([ ext{H}^+])$

• Pure water autoionization (conceptual):

  • $ext{H}_2 ext{O} <br>ightleftharpoons ext{H}^+ + ext{OH}^-$

• Neutral pH and biological window:

  • $ext{pH} = 7 ext{ is neutral (in pure water).}$

  • Blood pH window: $ext{pH} \, ext{roughly } 7.35 ext{ to } 7.45.$

• Bicarbonate buffering (example):

  • $ext{HCO}<em>3^- + ext{H}^+ ightarrow ext{H}</em>2 ext{CO}_3$

Note: The transcript contains a potential inaccuracy regarding ammonium (NH4+) acting as a base with a negative charge in that context. In physiology, ammonia (NH3) is the typical base that accepts protons, while ammonium (NH4+) is the conjugate acid. The notes above reflect the point as stated in the transcript but you may want to review this detail for accuracy in your course materials.