AP Chem Chapter 8
Acids and Bases Overview
Bronsted-Lowry Theory
Acid: A substance that donates protons (H+) in a chemical reaction.
Base: A substance that accepts protons (H+) in a chemical reaction.
Arrhenius Theory
Acid: A substance that produces hydronium ions (H3O+) when dissolved in water.
Base: A substance that produces hydroxide ions (OH-) when dissolved in water.
Key Concepts
Acid-base reactions occur primarily in aqueous solutions.
Strong acids and strong bases completely ionize in water, while weak acids and weak bases ionize only partially.
pH and pOH Calculations
pH: A measure of the acidity or basicity of a solution. It is defined as the negative logarithm (base 10) of the concentration of hydronium ions (H3O+):
pH = -log[H3O+]pOH: A measure of the basicity of a solution, defined as the negative logarithm of the concentration of hydroxide ions (OH-):
pOH = -log[OH-]The relationship between pH and pOH is given by:
pH + pOH = 14at 25°C (298K).
Strong and Weak Acids/Bases
For strong acids, since they completely ionize, the concentration of H3O+ directly corresponds to the concentration of the acid. For example, a 0.1 M HCl solution will have a pH of:
pH = -log(0.1) = 1For weak acids, the pH must be calculated using the acid dissociation constant (Ka) and an equilibrium expression, often requiring the use of an ICE table (Initial, Change, Equilibrium) to determine the concentrations at equilibrium before using the logarithm.
Similarly, for bases, the pOH can be calculated directly from strong bases or using Kb for weak bases.
Equilibrium Constants
The equilibrium constant for the autoionization of water (Kw) at 298K is 1×10^-14.
The relationship between pK and pKw is given by:
pKw = 14 = pKa + pKb.
Mixture Behaviors
Strong Acid + Strong Base: The pH of the solution is determined by the excess reagent present.
Weak Acid + Strong Base: If the weak acid is in excess, the solution acts as a buffer; if the strong base is in excess, the pH is governed by the excess OH- concentration.
Weak Base + Strong Acid: If the weak base is in excess, the solution acts as a buffer; if the strong acid is in excess, the pH is determined by the excess H3O+ concentration.
Weak Acid + Weak Base: The solution reaches an equilibrium state, where both acid and base affect the pH.
Titrations
Equivalence Point: The stage in a titration where the number of moles of titrant added equals the number of moles of the analyte present.
Polyprotic Acid: An acid capable of donating more than one proton; it has multiple dissociation steps, resulting in multiple equivalence points on a titration curve.
Indicators and pH Ranges
Methyl Orange: pH 5-6; turns red in acidic environments and yellow in basic environments.
Methyl Red: pH 6-7; red in acid, yellow in base.
Litmus: pH 7-8; red in acid and blue in base.
Phenolphthalein: pH 9-10; colorless in acid and pink in base.
Structure of Acids and Bases
Strong acids have very weak conjugate bases, making them less stable as they rarely accept protons. Conversely, weak acids have stronger conjugate bases.
The presence of electronegative elements in an acid or base enhances the stability of its conjugate base, thereby increasing the acid strength due to the greater ease of proton donation.
Henderson-Hasselbalch Equations
The pH of the buffer is related to the pKa of the acid and the concentration ratio of the conjugate advice-base pair
pH=pKa + log [salt]/[acid]
pH + pOH = 14
pKa/pKb is fixed for the weak acid or base
Changing concentrations of the components does affect the capacity of the buffer
The Henderson-Hasselbalch equation relates the pH of a buffer solution to the pKa of its acid and the concentration ratio of its conjugate base to the acid. It is formulated as:
[ \text{pH} = \text{pKa} + \log \left( \frac{[\text{salt}]}{[\text{acid}]} \right) ]
Key Points:
The pKa represents the acid dissociation constant, indicating the strength of the acid.
This equation helps calculate the pH of buffer solutions when the concentrations of the acid and its conjugate base are known.
If you change the concentrations of the components in the buffer, it can affect the buffer capacity, but the pKa remains fixed for a given weak acid or base.
Applications:
Used in biochemistry to maintain pH levels in various biological systems and reactions.
Buffer Capacity
Buffer capacity refers to the ability of a buffer solution to resist changes in pH when small amounts of acid or base are added. It is defined as the amount of acid or base that can be added to a buffer solution before a significant change in pH occurs.
Key Factors Influencing Buffer Capacity:
Concentration of Buffer Components: Buffers are typically composed of a weak acid and its conjugate base or a weak base and its conjugate acid. Higher concentrations of these components increase the buffer capacity.
pH Range: Buffer capacity is most effective within the pH range close to the pKa of the weak acid or base.
Calculation of Buffer Capacity: Although there is no single formula for calculating buffer capacity, it can be determined experimentally by measuring the pH change after adding known amounts of strong acid or base to the buffer solution.
Applications: Buffer capacity is essential in biological and chemical processes where maintaining a stable pH is crucial, such as in biochemical reactions, pharmaceutical formulations, and metabolic processes in living organisms.
pH and Solubility
pH and Its Impact on Solubility:
The pH of a solution can significantly affect the solubility of various compounds, particularly salts and acids.
For salts that contain weak acids or bases, changes in pH can shift the equilibrium and thus affect solubility.
Le Chatelier’s Principle:
Le Chatelier’s principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change.
Applying this principle to solubility:
Acidic Conditions: Increasing the concentration of H⁺ ions can shift the equilibrium for weak acids, increasing their solubility.
Basic Conditions: Conversely, for weak bases, increasing OH⁻ concentration can decrease their solubility.
Conclusion:
Understanding the relationship between pH, solubility, and Le Chatelier’s principle is crucial in fields like chemistry, environmental science, and pharmacology, where controlling solubility is often essential.