6.1 Acid-Base Theories and pH
2 min read•july 25, 2024
Acid-base theories are fundamental to understanding chemical reactions in biological systems. From Arrhenius to Lewis, these concepts explain how substances interact, donate or accept protons, and form bonds in aqueous solutions.
pH and calculations quantify acidity and basicity, crucial for maintaining balance in living organisms. These measurements, along with Ka values and water autoionization, help us grasp the delicate equilibria that sustain life processes.
Acid-Base Theories
Acid-base theories
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- defines acids as substances releasing H+ ions in water, bases release OH- ions (hydrochloric acid, sodium hydroxide)
- Brønsted-Lowry theory broadens definition: acids donate protons, bases accept protons, forming conjugate acid-base pairs (ammonia accepting H+ from water)
- further expands concept: acids accept electron pairs, bases donate electron pairs, forming coordinate covalent bonds (boron trifluoride accepting electrons from ammonia)
pH and pOH calculations
- pH measures acidity: negative log of H3O+ concentration (lemon juice pH ~2)
- pOH measures basicity: negative log of OH- concentration (bleach pOH ~4)
- pH and pOH sum to 14 at 25℃, allowing calculation of one from the other
- Convert between concentration and pH/pOH using and (blood pH 7.4 corresponds to [H3O+] = 3.98 × 10^-8 M)
pH, pOH, and Ka relationships
- Ka (acid dissociation constant) measures acid strength (acetic acid Ka = 1.8 × 10^-5)
- (negative log of Ka) often used instead of Ka for convenience
- relates pH, pKa, and concentrations:
- Buffers resist pH changes, composed of weak acid and conjugate base ( in blood)
Water autoionization in equilibria
- Water molecules react, producing H3O+ and OH- ions:
- Kw (ion product of water) equals [H3O+][OH-], constant 1.0 × 10^-14 at 25℃
- Neutral solutions: [H3O+] = [OH-] = 1.0 × 10^-7 M, pH = 7 (pure water)
- Acidic solutions: [H3O+] > [OH-], pH < 7 (stomach acid pH ~2)
- Basic solutions: [H3O+] < [OH-], pH > 7 (baking soda solution pH ~9)
Key Terms to Review (15)
Arrhenius Theory: The Arrhenius Theory is a fundamental concept in acid-base chemistry that defines acids as substances that increase the concentration of hydrogen ions (H+) in aqueous solution and bases as substances that increase the concentration of hydroxide ions (OH-) in aqueous solution. This theory provides a straightforward way to classify substances as acids or bases based on their behavior in water, establishing a clear connection between acid-base reactions and the concept of pH.
Bicarbonate buffer: The bicarbonate buffer system is a crucial physiological buffer system that helps maintain pH balance in biological fluids. It operates primarily through the equilibrium between carbonic acid (H$_2$CO$_3$) and bicarbonate ion (HCO$_3^-$), allowing for the regulation of acid-base homeostasis in the body. This system is essential for managing changes in pH that can occur due to metabolic processes or the intake of acidic or basic substances.
Bronsted-Lowry Theory: The Bronsted-Lowry Theory defines acids as substances that donate protons (H\(^+\)) and bases as substances that accept protons. This theory expands the concept of acids and bases beyond the limitations of the Arrhenius definition, which focuses on the production of H\(^+\) and OH\(^-\) ions in water, allowing for a broader understanding of acid-base reactions in various solvents.
Buffer systems: Buffer systems are solutions that resist significant changes in pH when small amounts of acids or bases are added. They play a crucial role in maintaining homeostasis in biological systems by stabilizing the pH levels, which is vital for enzyme activity and overall cellular function. Buffer systems primarily consist of a weak acid and its conjugate base, or a weak base and its conjugate acid, which work together to neutralize added acids or bases.
Enzyme activity at different pH levels: Enzyme activity at different pH levels refers to the variation in the rate of enzymatic reactions as influenced by the acidity or alkalinity of the environment. Each enzyme has an optimal pH range where it functions best, and deviations from this range can lead to decreased activity or denaturation. Understanding how pH affects enzymes is crucial because it impacts biological processes and metabolic pathways.
Henderson-Hasselbalch Equation: The Henderson-Hasselbalch equation is a mathematical formula used to estimate the pH of a buffer solution based on the concentration of an acid and its conjugate base. This equation highlights the relationship between pH, pKa, and the ratio of the concentrations of the protonated and deprotonated forms of an acid, making it essential in understanding how buffers work in biological systems and their role in maintaining stable pH levels.
Lewis Theory: Lewis Theory defines acids and bases based on electron pair behavior rather than hydrogen ion transfer. According to this theory, a Lewis acid is an electron pair acceptor, while a Lewis base is an electron pair donor. This perspective broadens the understanding of acid-base reactions beyond traditional definitions, allowing for more complex interactions in chemical systems.
Litmus: Litmus is a water-soluble dye derived from lichens, commonly used as an acid-base indicator to determine the pH of a solution. It changes color in response to acidity or alkalinity, turning red in acidic conditions (pH < 7) and blue in alkaline conditions (pH > 7). This property makes litmus a simple and effective tool for visualizing changes in pH, helping in various scientific and educational applications.
PH meter calibration: pH meter calibration is the process of adjusting a pH meter to ensure accurate readings of hydrogen ion concentration in a solution. This calibration is crucial because it aligns the instrument with standard reference points, usually achieved using pH buffer solutions. Accurate calibration is essential for reliable measurements, as it directly affects the interpretation of acid-base properties in various chemical and biological contexts.
PH Scale: The pH scale is a logarithmic scale used to measure the acidity or basicity of a solution, ranging from 0 to 14. It provides a quantitative basis for understanding how substances interact chemically in biological systems, influencing processes like enzyme activity and cellular functions.
Phenolphthalein: Phenolphthalein is a colorless organic compound often used as a pH indicator in titrations, changing color from colorless to pink as the pH transitions from acidic to slightly basic. This transition occurs around a pH of 8.2 to 10, making it useful for detecting the endpoint of strong acid-strong base titrations and determining pH levels in various chemical solutions.
Phosphate buffer: A phosphate buffer is a solution that resists changes in pH when small amounts of an acid or a base are added, primarily composed of a weak acid (dihydrogen phosphate, H2PO4^-) and its conjugate base (hydrogen phosphate, HPO4^2-). This type of buffer system plays a crucial role in biological systems, maintaining the pH within a narrow range that is vital for various biochemical processes.
PKa: pKa is a quantitative measure of the strength of an acid in solution, defined as the negative logarithm of the acid dissociation constant (Ka). It provides insight into how readily an acid donates protons (H+) to a solution, with lower pKa values indicating stronger acids. Understanding pKa is essential for predicting the behavior of acids and bases in chemical reactions, especially in relation to equilibrium and buffer solutions.
POH: pOH is a measure of the concentration of hydroxide ions (OH⁻) in a solution and is defined as the negative logarithm of the hydroxide ion concentration. It is closely related to pH, which measures hydrogen ion concentration, and together they help to describe the acidity or basicity of a solution. The pOH scale typically ranges from 0 to 14, with lower values indicating more basic conditions and higher values indicating more acidic conditions.
Titration: Titration is a laboratory technique used to determine the concentration of an unknown solution by reacting it with a solution of known concentration. This method is essential for analyzing acid-base reactions, where an acid reacts with a base to form water and a salt, allowing the pH to be measured and controlled throughout the process. By adding titrant gradually until the endpoint is reached, scientists can accurately assess the amount of substance in the solution being tested.