9.3 Acids, Bases & pH Balance

9.3 Acids, Bases & pH Balance

Acid-base chemistry controls fundamental biological and chemical processes, from digestive gastric acid production to the precise enzymatic regulation of cellular metabolism. For the TEAS 7 Science exam, candidates must understand acid-base definitions, physical and chemical properties, strong vs. weak ionization, the logarithmic pH scale, neutralization reactions, and physiological buffer homeostasis.

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Acid-Base Definitions: Arrhenius vs. Brønsted-Lowry

Two major theoretical models define acid-base chemistry:

1. Arrhenius Theory

Proposed by Svante Arrhenius, this model defines acids and bases by the ions they release in aqueous solution:

• Arrhenius Acid: Any substance that increases the concentration of hydrogen ions ($H^+$) or hydronium ions ($H_3O^+$) when dissolved in water. $$HCl(aq) \longrightarrow H^+(aq) + Cl^-(aq)$$
• Arrhenius Base: Any substance that increases the concentration of hydroxide ions ($OH^-$) when dissolved in water. $$NaOH(aq) \longrightarrow Na^+(aq) + OH^-(aq)$$

2. Brønsted-Lowry Theory

A broader model that applies beyond aqueous solutions, defining acids and bases by proton ($H^+$) transfer:

• Brønsted-Lowry Acid: A proton ($H^+$) donor.
• Brønsted-Lowry Base: A proton ($H^+$) acceptor.

Conjugate Acid-Base Pairs

When a Brønsted-Lowry acid donates a proton, it forms a conjugate base. When a Brønsted-Lowry base accepts a proton, it forms a conjugate acid:

$$\underbrace{NH_3(aq)}_{\text{Base}} + \underbrace{H_2O(l)}_{\text{Acid}} \rightleftharpoons \underbrace{NH_4^+(aq)}_{\text{Conjugate Acid}} + \underbrace{OH^-(aq)}_{\text{Conjugate Base}}$$

In this reaction, $NH_3$ accepts a proton to become $NH_4^+$, while $H_2O$ acts as a Brønsted-Lowry acid by donating a proton.

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Chemical Properties of Acids & Bases

Acids and bases exhibit distinct physical and chemical characteristics:

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Strong vs. Weak Acids and Bases

The strength of an acid or base is determined entirely by its degree of ionization (dissociation) in water, not by its overall concentration.

Strong Acids & Bases

• Strong Acids: Dissociate completely ($100\%$ ionization) in aqueous solution into $H^+$ ions and anions. The reaction arrow goes to completion ($\rightarrow$).
  • Common Strong Acids: Hydrochloric acid ($HCl$), Nitric acid ($HNO_3$), Sulfuric acid ($H_2SO_4$).
• Strong Bases: Dissociate completely ($100\%$ ionization) into metal cations and $OH^-$ ions.
  • Common Strong Bases: Sodium hydroxide ($NaOH$), Potassium hydroxide ($KOH$), Calcium hydroxide ($Ca(OH)_2$).

Weak Acids & Bases

• Weak Acids: Dissociate only partially ($<5\%$ ionization) in water, establishing a dynamic equilibrium between un-ionized molecules and ions ($\rightleftharpoons$).
  • Common Weak Acids: Acetic acid ($CH_3COOH$), Carbonic acid ($H_2CO_3$), Citric acid ($C_6H_8O_7$).
• Weak Bases: Ionize only partially in water to yield small amounts of $OH^-$ ions.
  • Common Weak Base: Ammonia ($NH_3$).

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The Logarithmic pH and pOH Scales

The pH scale measures the acidity or alkalinity of a solution on a logarithmic scale ranging from 0 to 14 at 25°C.

$$\text{pH} = -\log[H^+]$$ $$\text{pOH} = -\log[OH^-]$$ $$\text{pH} + \text{pOH} = 14.0$$

Understanding the Logarithmic Relationship

Because the pH scale is logarithmic (base 10), each 1-unit change in pH represents a 10-fold change in hydrogen ion concentration $[H^+]$:

• A solution with a $\text{pH}$ of 4 has 10 times more $H^+$ ions than a solution with a $\text{pH}$ of 5.
• A solution with a $\text{pH}$ of 3 has 100 times ($10 \times 10$) more $H^+$ ions than a solution with a $\text{pH}$ of 5.
• A solution with a $\text{pH}$ of 2 has 1,000 times ($10 \times 10 \times 10$) more $H^+$ ions than a solution with a $\text{pH}$ of 5.

Solution Classification at 25°C

• Acidic Solution: $[H^+] > 1.0 \times 10^{-7}\text{ M} \implies \text{pH} < 7.0$
• Neutral Solution: $[H^+] = 1.0 \times 10^{-7}\text{ M} \implies \text{pH} = 7.0$
• Basic (Alkaline) Solution: $[H^+] < 1.0 \times 10^{-7}\text{ M} \implies \text{pH} > 7.0$

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Neutralization Reactions

A neutralization reaction occurs when an acid reacts with a base to produce an ionic compound (a salt) and water:

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

Molecular & Net Ionic Equations

When hydrochloric acid ($HCl$) reacts with sodium hydroxide ($NaOH$):

$$\text{Molecular: } HCl(aq) + NaOH(aq) \longrightarrow NaCl(aq) + H_2O(l)$$

Because $HCl$, $NaOH$, and $NaCl$ dissociate completely into aqueous ions ($H^+, Cl^-, Na^+, OH^-$), the spectator ions ($Na^+$ and $Cl^-$) can be canceled out to reveal the net ionic equation for all strong acid-strong base neutralizations:

$$\text{Net Ionic Equation: } H^+(aq) + OH^-(aq) \longrightarrow H_2O(l)$$

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Buffer Systems & Blood pH Homeostasis

A buffer solution is a chemical system that resists drastic changes in pH when small amounts of strong acid or strong base are added. Buffers consist of a weak acid and its conjugate base (or a weak base and its conjugate acid).

The Bicarbonate Buffer System in Human Blood

Human arterial blood plasma must be maintained within a remarkably narrow physiological pH range of 7.35 to 7.45. Deviation below 7.35 causes acidosis, while elevation above 7.45 causes alkalosis. Both conditions disrupt protein structure and enzymatic activity, leading to severe cell damage or death.

The primary extracellular buffer maintaining blood pH is the carbonic acid-bicarbonate buffer system:

$$CO_2(g) + H_2O(l) \rightleftharpoons \underbrace{H_2CO_3(aq)}_{\text{Carbonic Acid (Weak Acid)}} \rightleftharpoons \underbrace{HCO_3^-(aq)}_{\text{Bicarbonate Ion (Conjugate Base)}} + H^+(aq)$$

Physiological Compensation Mechanisms

• Response to Acid Stress (Addition of $H^+$): If metabolic activity increases blood $H^+$ (acidosis), excess $H^+$ combines with $HCO_3^-$ to form $H_2CO_3$. The lungs then hyperventilate to exhale excess $CO_2$, shifting equilibrium to the left and restoring normal pH.
• Response to Base Stress (Loss of $H^+$): If blood $H^+$ drops (alkalosis), $H_2CO_3$ dissociates into $HCO_3^-$ and $H^+$. The respiratory rate decreases (hypoventilation) to retain $CO_2$, raising carbonic acid levels and lowering pH back to baseline.