Chemistry and Acid-Base Equilibrium for Medicine

General Concepts and Properties of Acids and Bases

According to the chemical principles outlined by González Muradás and Montagut Bosque (2015), acids and bases possess distinct physical and chemical characteristics. Acids are defined by their sour flavor and their behavior in aqueous solutions, where they release hydrogen ions, specifically indicated as $H^+$. Chemically, acids are characterized by their ability to donate protons ($H^+$) and their capacity to accept electrons. Conversely, bases are recognized by their bitter taste and a slippery or soapy consistency. In an aqueous environment, bases release hydroxyl ions, denoted as $OH^-$. From a reactive standpoint, bases can accept protons ($H^+$) and serve as electron donors.

Classification of Acids and Bases by Strength

The strength of an acid or base is determined by its degree of dissociation in an aqueous solution. Strong acids generally exhibit a very low pH because they dissociate $100\%$ into their constituent ions when dissolved in water. Key examples of strong acids include hydrochloric acid ($HCl$), nitric acid ($HNO_3$), sulfuric acid ($H_2SO_4$), hydrobromic acid ($HBr$), hydroiodic acid ($HI$), and perchloric acid ($HClO_4$). Weak acids, such as acetic acid ($CH_3COOH$) and formic acid ($HCOOH$), have a higher pH than strong acids because they only partially dissociate in solution. It is noted that molecules containing carboxyl groups ($-COOH$) typically behave as weak acids.

Similarly, bases are categorized as strong or weak. Strong bases possess a very high pH due to the high concentration of $OH^-$ ions produced through $100\%$ dissociation in aqueous media. These are primarily the hydroxides of alkali and alkaline earth metals, including sodium hydroxide ($NaOH$), potassium hydroxide ($KOH$), lithium hydroxide ($LiOH$), calcium hydroxide ($Ca(OH)_2$), barium hydroxide ($Ba(OH)_2$), and strontium hydroxide ($Sr(OH)_2$). Weak bases, such as ammonia ($NH_3$) and various amines ($R-NH_2$), exhibit a pH that is less basic than strong bases because they only partially dissociate.

Scientific Theories of Acid-Base Equilibria

There are three primary historical and scientific frameworks used to define acid-base chemistry as established by González Muradás, R. M., and Montagut Bosque, P. (2015):

1. Arrhenius Theory: This theory defines an acid as any substance that releases hydrogen ions ($H^+$) when placed in an aqueous solution. A base is defined as a substance that releases hydroxyl ions ($OH^-$) in the same environment.
2. Brønsted-Lowry Theory: This model focuses on the transfer of protons. An acid is characterized as a proton donor ($H^+$), while a base is characterized as a proton acceptor ($H^+$). For example, in a reaction where hydrochloric acid reacts with ammonia, $HCl$ acts as the acid by donating a proton, and $NH_3$ acts as the base by accepting it.
3. Lewis Theory: This theory moves beyond protons and hydroxyl ions to focus on electron pairs. A Lewis acid is defined as a substance that accepts a pair of electrons, such as boron trifluoride ($BF_3$). A Lewis base is a substance that donates a pair of electrons, such as ammonia ($NH_3$).

The Concept of pH and the Potential of Hydrogen

The potential of hydrogen (pH) is a quantitative measure of how acidic or basic a solution is, based on the concentration of hydrogen ions present. The scale is defined as follows:

• pH < 7: Acidic. The closer the value is to $0$, the more acidic the substance.
• pH 7: Neutral.
• pH > 7: Basic or alkaline. The closer the value is to $14$, the more basic the substance.

Mathematically, pH is defined as the negative logarithm (base $10$) of the activity or concentration of hydrogen ions ($H^+$). The concentration is typically measured in moles per decimeter cubed ($mol\,dm^{-3}$). The formula is expressed as: $pH = -\log_{10}([H^+])$

The relationship between hydrogen ion concentration and pH is inversely proportional: as the concentration of $[H^+]$ increases, the pH value decreases (becoming more acidic). Conversely, as the concentration of $[H^+]$ decreases, the pH value increases. The concentration of hydroxyl ions ($[OH^-]$) determines how basic a solution is; higher levels of $[OH^-]$ correlate with higher basicity. The standard scale ranges from $10^0$ ($pH = 0$) to $10^{-14}$ ($pH = 14$) for hydrogen ion concentration.

Conjugate Acid-Base Pairs and Chemical Reactions

According to Timberlake (2011), in any reaction where a substance behaves as an acid by donating a proton ($H^+$), there must be another substance acting as a base to capture that proton. When an acid loses a proton, it transforms into its conjugate base. When a base accepts a proton, it transforms into its conjugate acid. A classic example of this is the reaction between ammonia and hydrogen chloride:

$NH_3 + HCl \rightarrow NH_4^+ + Cl^-$

In this scenario, $NH_3$ is the base and $HCl$ is the acid. Upon reaction, ammonia becomes the ammonium ion ($NH_4^+$), which is the conjugate acid, and hydrogen chloride becomes the chloride ion ($Cl^-$), which is the conjugate base.

In chemical reactions, identifying these pairs requires observing proton movement:

• Acid: The species that loses protons.
• Base: The species that gains protons.
• Example: In the reaction of acetic acid ($CH_3COOH$) with water ($H_2O$), $CH_3COOH$ loses a proton to become the conjugate base acetate ($CH_3COO^-$), while water gains a proton to become the conjugate acid hydronium ($H_3O^+$).

Buffering Systems and Amortiguadores

Buffer solutions, or "sustancias amortiguadoras," are substances capable of maintaining the pH of a solution at a specific, stable value even when small amounts of acids or bases are added. For illustration, if lemon juice (an acid) is added to plain water, the pH drops significantly. However, if that water contained a buffer, the pH change would be negligible.

Buffers are typically composed of a weak acid and its conjugate base, or a weak base and its conjugate acid. Common chemical examples include:

• Acetic acid ($CH_3COOH$) and sodium acetate ($CH_3COO^-$).
• Formic acid ($HCOOH$) and sodium formate ($HCOONa$).
• Carbonic acid ($H_2CO_3$) and bicarbonate ($HCO_3^-$).

Clinical Applications: pH in the Human Body

Maintaining pH balance is critical for human physiology. The standard pH range for human blood is strictly regulated between $7.35$ and $7.45$. Deviations from this range lead to distinct clinical conditions:

• Alkalosis (pH > 7.45): Symptoms include dizziness, nausea, irritability, convulsions, confusion, muscle spasms, and hand tremors.
• Acidosis (pH < 7.35): Symptoms include fatigue, weakness, drowsiness, confusion, headache (cefalea), nausea, vomiting, hyperventilation, and potentially coma.

Different biological compartments maintain specific pH levels to facilitate their functions:

• Oral Cavity: $pH \, 6.8 - 7.5$
• Stomach Cavity: $pH \, 1.5 - 2.0$
• Duodenum: $pH \, 5.6 - 8.0$
• Small Intestine: $pH \, 7.2 - 7.5$
• Colon: $pH \, 7.9 - 8.5$

Physiological Buffer Systems in the Organism

The body utilizes physiological buffers (tampones fisiológicos) to keep the pH within normal ranges. The three primary systems are:

1. Bicarbonate System ($H_2CO_3 / HCO_3^-$): This is the most important system for regulating blood pH. In cases of acidosis (excess $H^+$), the bicarbonate neutralizes the hydrogen ions, forming $CO_2$ which is then expelled by the lungs through respiration. In cases of alkalosis (low $H^+$), carbonic acid releases hydrogen ions to lower the pH, while excess bicarbonate can be excreted via the kidneys in urine.
2. Phosphate System ($H_2PO_4^- / HPO_4^{2-}$): This buffer is present in all cells and is the most significant regulator of urinary pH. Under acidic conditions, hydrogen phosphate ($HPO_4^{2-}$) accepts excess hydrogen ions to form dihydrogen phosphate ($H_2PO_4^-$). Under alkaline conditions, the dihydrogen phosphate releases protons to neutralize excess hydroxyl ions entering the cell.
3. Protein Buffers: Proteins contain amino acids with amino groups ($-NH_2$) and carboxylic acid groups ($-COOH$). At physiological pH, these exist as carboxylate ions ($COO^-$) and ammonium ions ($NH_3^+$). If the environment becomes acidic, the carboxylate groups capture excess $H^+$ to reform the carboxylic acid. If the blood becomes alkaline, the $NH_3^+$ ion releases a proton to become $NH_2$.

Questions & Discussion

Question 1 (Case Study): A 28-year-old female patient presents with drowsiness and weakness. Arterial blood gas analysis reveals a pH of $7.23$. What is the diagnosis and is the $H^+$ load high or low? Response: The diagnosis is Acidosis because the pH is below the normal range of $7.35$. The load (concentration) of $H^+$ ions is high, as pH is inversely proportional to the concentration of hydrogen ions.

Question 2 (Identification): Which of the following could function as buffer substances: Sodium chloride and potassium chloride; Hydrochloric acid and sodium sulfate; Sodium hydroxide and potassium nitrate; or Acetic acid and sodium acetate? Response: Acetic acid and sodium acetate, because a buffer must consist of a weak acid and its conjugate base.

Question 3 (Case Study): A 30-year-old patient has a blood gas pH of $7.25$. What does this indicate? What will be the concentration of $H^+$ in the blood? Will any buffer system act in this case? Response: The pH indicates Acidosis ($pH < 7.35$). The concentration of $H^+$ will be high ($10^{-7.25}\,mol\,dm^{-3}$). In this case, the bicarbonate system would act by neutralization, and the protein and phosphate systems would also engage to try to restore equilibrium.

Question 4 (Reaction Identification): Identify the donors and acceptors in the reaction of $NH_3 + H_2O$. Response: In the reaction $NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$, $H_2O$ loses a proton (Acid) to become the conjugate base $OH^-$, and $NH_3$ gains a proton (Base) to become the conjugate acid $NH_4^+$.