Acids, Bases, Salts, pH, and Buffers: Foundations for Anatomy & Physiology

Acids, Bases, Salts, pH, and Buffers

Introduction

Understanding acids, bases, salts, pH, and buffers is essential for comprehending the chemical environment of the human body. These concepts are foundational in Anatomy & Physiology, as they influence cellular function, enzyme activity, and homeostasis.

Water Dissociation and Ion Formation

Dissociation of Water

Water molecules can dissociate into ions, a process critical for many physiological reactions.

• Dissociation: The separation of water molecules into hydrogen ions (H+) and hydroxide ions (OH-).

• Equation:

$\mathrm{H_2O} \rightleftharpoons \mathrm{H^+} + \mathrm{OH^-}}$

• In pure water, the concentrations of H+ and OH- are equal.

• Only about 1 in 10 million water molecules dissociates at any given time.

Acids

Definition and Properties

Acids are substances that release hydrogen ions (H+) when dissolved in water. They are also known as proton donors.

• Acidic solutions: Have a high concentration of H+ ions.

• Example: Hydrochloric acid (HCl) dissociates completely in water:

$\mathrm{HCl} \rightarrow \mathrm{H^+} + \mathrm{Cl^-}}$

• Strong acids: Completely dissociate in solution (e.g., HCl).

• Weak acids: Partially dissociate in solution (e.g., acetic acid, carbonic acid).

• Weak acids are represented by reversible reactions:

$\mathrm{CH_3COOH} \rightleftharpoons \mathrm{CH_3COO^-} + \mathrm{H^+}}$

$\mathrm{H_2CO_3} \rightleftharpoons \mathrm{HCO_3^-} + \mathrm{H^+}}$

• Weak acids have a smaller effect on pH compared to strong acids.

Bases

Definition and Properties

Bases are substances that either release hydroxide ions (OH-) or accept hydrogen ions (H+) when dissolved in water. They are also known as proton acceptors.

• Basic solutions: Have a low concentration of H+ ions.

• Example: Sodium hydroxide (NaOH) dissociates completely in water:

$\mathrm{NaOH} \rightarrow \mathrm{Na^+} + \mathrm{OH^-}}$

• OH- ions bind free H+ ions, reducing their concentration.

• Strong bases: Completely dissociate in solution (e.g., NaOH).

• Weak bases: Partially dissociate or accept H+ ions (e.g., ammonia, NH3).

$\mathrm{NH_3} + \mathrm{H^+} \rightarrow \mathrm{NH_4^+}}$

• Weak bases have a smaller effect on pH compared to strong bases.

Salts and Electrolytes

Formation and Importance

Salts are ionic compounds formed when acids react with bases. They are important electrolytes in body fluids.

• General reaction:

$\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$

• Example: Hydrochloric acid reacts with sodium hydroxide:

$\mathrm{HCl} + \mathrm{NaOH} \rightarrow \mathrm{NaCl} + \mathrm{H_2O}$

• Electrolytes: Substances that dissociate into ions in solution, enabling electrical conductivity. Key physiological electrolytes include Na+, K+, Ca2+, Cl-, and HCO3-.

pH Scale

Definition and Calculation

The pH scale measures the concentration of hydrogen ions in a solution, indicating its acidity or basicity.

• pH formula:

$\mathrm{pH} = -\log_{10}[\mathrm{H^+}]$

• The pH scale ranges from 0 (most acidic) to 14 (most basic), with 7 being neutral.

• Each unit change in pH represents a tenfold change in H+ concentration.

• Neutral solution: [H+] = [OH-], pH = 7.

• Acidic solution: [H+] > [OH-], pH < 7.

• Basic solution: [H+] < [OH-], pH > 7.

pH Scale Table

pH Value	Relative [H+]	Solution Type
0-6	High	Acidic
7	Equal to [OH-]	Neutral
8-14	Low	Basic (Alkaline)

Logarithmic Nature of pH

• Each pH unit change = 10x change in [H+].

• Example: pH 5 is 100 times more acidic than pH 7.

Buffers

Definition and Function

Buffers are mixtures of weak acids and their conjugate bases (or weak bases and their conjugate acids) that resist changes in pH when small amounts of acids or bases are added.

• Importance: Buffers maintain stable pH in body fluids, essential for enzyme function and metabolic processes.

• Mechanism: If [H+] increases, the buffer's weak base binds H+ to lower pH. If [H+] decreases, the buffer's weak acid releases H+ to raise pH.

• Major physiological buffer: The carbonic acid (H2CO3)/bicarbonate (HCO3-) system.

$\mathrm{H_2CO_3} \rightleftharpoons \mathrm{HCO_3^-} + \mathrm{H^+}}$

• This system helps maintain blood pH within a narrow range (7.35–7.45).

Summary Table: Strong vs. Weak Acids and Bases

Type	Definition	Example	Dissociation	Effect on pH
Strong Acid	Completely dissociates in water	HCl	100%	Large
Weak Acid	Partially dissociates in water	H2CO3, CH3COOH	Partial	Small
Strong Base	Completely dissociates in water	NaOH	100%	Large
Weak Base	Partially dissociates or accepts H+	NH3	Partial	Small

Key Takeaways

• Acids and bases are crucial for physiological processes and are defined by their ability to donate or accept H+ ions.

• Salts and electrolytes are vital for nerve and muscle function.

• The pH scale is logarithmic; small changes in pH reflect large changes in H+ concentration.

• Buffers maintain pH stability, essential for homeostasis in the human body.