💏Intro to Chemistry Unit 14 – Acid–Base Equilibria

Key Concepts and Definitions

• Acids produce hydrogen ions (H+) in aqueous solutions and have a sour taste (lemons, vinegar)
• Bases produce hydroxide ions (OH-) in aqueous solutions and have a bitter taste (soap, baking soda)
• Arrhenius theory defines acids as H+ donors and bases as OH- donors in aqueous solutions
• Brønsted-Lowry theory expands the definition of acids as proton (H+) donors and bases as proton acceptors
  • Allows for the inclusion of substances without OH- (ammonia)
• Lewis theory further broadens the definition of acids as electron pair acceptors and bases as electron pair donors
• Conjugate acid-base pairs consist of a species and its corresponding acid or base formed by the loss or gain of a proton (H+)
• Amphoteric substances can act as both an acid and a base (water, amino acids)

Acids and Bases: Properties and Theories

• Acids have a pH less than 7, while bases have a pH greater than 7
  • Neutral substances have a pH equal to 7 (pure water at 25°C)
• Acids and bases can be classified as strong or weak based on their degree of ionization in aqueous solutions
  • Strong acids and bases completely ionize (hydrochloric acid, sodium hydroxide)
  • Weak acids and bases partially ionize (acetic acid, ammonia)
• Acids and bases can neutralize each other, forming water and a salt (NaOH + HCl -> NaCl + H2O)
• The Arrhenius theory has limitations, as it only applies to aqueous solutions and doesn't account for substances that behave as acids or bases without containing H+ or OH-
• The Brønsted-Lowry theory introduces the concept of conjugate acid-base pairs, allowing for a more comprehensive understanding of acid-base reactions
• The Lewis theory is the most inclusive, as it considers the role of electron pairs in acid-base reactions and can be applied to non-aqueous systems

pH Scale and Calculations

• The pH scale is a logarithmic scale that measures the concentration of H+ ions in a solution, ranging from 0 to 14
  • Each unit change in pH represents a tenfold change in H+ concentration
• pH is calculated using the negative logarithm of the H+ concentration: pH = -log[H+]
• The concentration of OH- ions can be determined using the pOH scale: pOH = -log[OH-]
• The relationship between pH and pOH is: pH + pOH = 14
• The concentration of H+ and OH- ions in pure water at 25°C is 1 x 10^-7 M, resulting in a neutral pH of 7
• Acidic solutions have a higher concentration of H+ ions than OH- ions, while basic solutions have a higher concentration of OH- ions than H+ ions
• The self-ionization of water (Kw) is the equilibrium constant for the reaction: H2O + H2O <-> H3O+ + OH-
  • Kw = [H3O+][OH-] = 1 x 10^-14 at 25°C

Strength of Acids and Bases

• Strong acids and bases completely ionize in aqueous solutions, while weak acids and bases only partially ionize
• The strength of an acid or base is determined by its dissociation constant (Ka for acids, Kb for bases)
  • Higher Ka or Kb values indicate stronger acids or bases, respectively
• The acid dissociation constant (Ka) is the equilibrium constant for the dissociation of a weak acid: HA + H2O <-> H3O+ + A-
  • Ka = [H3O+][A-] / [HA]
• The base dissociation constant (Kb) is the equilibrium constant for the dissociation of a weak base: B + H2O <-> BH+ + OH-
  • Kb = [BH+][OH-] / [B]
• The relationship between Ka and Kb for a conjugate acid-base pair is: Ka x Kb = Kw
• The percent ionization of a weak acid or base can be calculated using the dissociation constant and the initial concentration
• The strength of an acid or base is not the same as its concentration; a strong acid can be present in a low concentration, while a weak acid can be present in a high concentration

Buffer Solutions

• Buffer solutions resist changes in pH when small amounts of acid or base are added or when dilution occurs
• Buffers consist of a weak acid and its conjugate base (or a weak base and its conjugate acid) in approximately equal concentrations
• The Henderson-Hasselbalch equation relates the pH of a buffer solution to the pKa of the weak acid and the ratio of the concentrations of the conjugate base and weak acid:
  • pH = pKa + log([A-] / [HA])
• Buffer capacity is the amount of acid or base that can be added to a buffer solution before a significant change in pH occurs
  • Factors affecting buffer capacity include the concentrations of the weak acid and conjugate base, and the ratio of these concentrations
• Buffers are important in maintaining stable pH in biological systems (blood, cytoplasm) and in various industrial and laboratory applications (fermentation, chemical synthesis)
• Common buffer systems include the carbonic acid-bicarbonate buffer (H2CO3/HCO3-) in blood and the acetic acid-acetate buffer (CH3COOH/CH3COO-) in fermentation processes
• Preparing a buffer solution involves mixing a weak acid with its conjugate base salt (or a weak base with its conjugate acid salt) in the desired ratio and concentration

Acid-Base Titrations

• Titration is a technique used to determine the concentration of an unknown acid or base by reacting it with a known concentration of a standard solution
• The equivalence point is reached when the moles of acid and base are equal, and the reaction is complete
  • At the equivalence point, the pH rapidly changes, indicating the end of the titration
• Indicators are substances that change color at specific pH ranges and are used to visually detect the endpoint of a titration
  • The endpoint is the point at which the indicator changes color, ideally close to the equivalence point
• Common indicators include phenolphthalein (colorless in acidic solutions, pink in basic solutions) and bromothymol blue (yellow in acidic solutions, blue in basic solutions)
• Titration curves plot the pH of the solution against the volume of titrant added, providing information about the strength and concentration of the acid or base being titrated
• Titrations can be used to determine the pKa of a weak acid or the pKb of a weak base by identifying the half-equivalence point on the titration curve
• Back titration is a technique used when the direct titration of a substance is not feasible, involving the addition of an excess of a known reagent and then titrating the excess with a standard solution

Real-World Applications

• pH regulation in the human body is crucial for maintaining homeostasis, with buffers like the carbonic acid-bicarbonate system in blood helping to keep the pH within a narrow range (7.35-7.45)
• Acid-base reactions are involved in the formation of caves and sinkholes, as acidic groundwater (carbonic acid) dissolves limestone (calcium carbonate)
• Ocean acidification, caused by the absorption of atmospheric CO2, has negative impacts on marine life, particularly organisms with calcium carbonate shells or skeletons (coral reefs, mollusks)
• Soil pH affects the availability of nutrients for plants, with most plants thriving in slightly acidic to neutral soils (pH 6.0-7.5)
  • Lime (calcium carbonate) can be added to acidic soils to raise the pH, while sulfur can be added to alkaline soils to lower the pH
• Acid-base chemistry is used in the production of various consumer products (soaps, detergents, shampoos) and in industrial processes (wastewater treatment, chemical synthesis)
• Antacids, which are weak bases, are used to neutralize excess stomach acid and relieve symptoms of acid reflux and indigestion (Tums, Rolaids)
• Acidic and basic solutions are used in cleaning products for their ability to dissolve and remove different types of stains and deposits (vinegar for mineral buildup, ammonia for grease)

Common Mistakes and How to Avoid Them

• Confusing the terms strong and concentrated, or weak and dilute
  • Remember that strength refers to the degree of ionization, while concentration refers to the amount of solute per unit volume
• Incorrectly using the pH scale by forgetting that it is logarithmic
  • A change of one pH unit represents a tenfold change in H+ concentration, not a linear change
• Misinterpreting titration curves by incorrectly identifying the equivalence point or endpoint
  • The equivalence point is where the moles of acid and base are equal, while the endpoint is where the indicator changes color
• Neglecting to consider the self-ionization of water (Kw) when calculating the pH of strong acids or bases
  • In dilute solutions, the contribution of H+ or OH- from water can significantly affect the pH
• Misapplying the Henderson-Hasselbalch equation by using the wrong form of the equation or incorrect values for pKa or concentrations
  • Double-check the equation and ensure that the concentrations used correspond to the conjugate base and weak acid
• Overlooking the importance of temperature when working with acid-base equilibria
  • Equilibrium constants (Ka, Kb, Kw) and pH are temperature-dependent, so be sure to use values corresponding to the given temperature
• Confusing the concepts of pH and pOH, or incorrectly calculating their relationship
  • Remember that pH + pOH = 14, and that pH is based on [H+] while pOH is based on [OH-]
• Not properly calibrating pH meters or using expired pH paper, leading to inaccurate measurements
  • Regularly calibrate pH meters using standard buffer solutions and use fresh pH paper within its expiration date