Chemical Equilibrium

Chemical Equilibrium

• A reversible reaction attains a state known as chemical equilibrium when the forward and reverse reaction rates equalize, resulting in no net change in reactant and product concentrations.

• Reversible reaction: Chemical reaction proceeds in both directions, from reactants to products (forward) and products to reactants (reverse).

• Forward reaction: Reaction proceeds from reactants to products.

• Reverse reaction: Reaction proceeds from products to reactants.

• Rate of reaction: Speed at which a chemical reaction occurs, change in concentration of a reactant or product per unit time.

• Equilibrium constant (K): Numerical value expresses the ratio of product concentrations to reactant concentrations at equilibrium, each raised to its stoichiometric coefficient.

• Kc: Equilibrium constant in terms of molar concentrations.

• Kp: Equilibrium constant in terms of partial pressures of gaseous reactants and products.

• Law of chemical equilibrium: For a reversible reaction at constant temperature, the ratio of product to reactant concentrations at equilibrium remains constant.

• Homogeneous equilibrium: Equilibrium where all reactants and products are in the same physical phase.

• Heterogeneous equilibrium: Equilibrium where reactants and products are in two or more different physical phases.

• Reaction quotient (Q): Measure of relative amounts of products and reactants at any given time, same form as K, calculated for systems not at equilibrium.

• Le Châtelier's principle: A system at equilibrium, when subjected to a change (concentration, pressure, temperature), shifts to relieve the stress.

• Endothermic reaction: Reaction absorbs heat (ΔH > 0), heat is considered a reactant.

• Exothermic reaction: Reaction releases heat (ΔH < 0), heat is considered a product.

• Stoichiometric coefficient: Numerical multiplier of each chemical formula in a balanced equation, indicates relative number of moles.

• Molar concentration (molarity): Amount of a substance in moles per liter of solution.

• Partial pressure: Pressure a gas in a mixture exerts if it occupied the entire volume alone.

Electrolytes

• Acids and bases interact in aqueous solutions exhibiting acid-base properties.
• The ion-product constant for water (Kw) relates [H3O+] and [OH-] in aqueous solutions, and can be calculated
• Weak electrolytes only partially dissociate in water, equilibrium exists between undissociated form and ions.
• Ka and Kb quantify ionization extent of weak acids and bases, defined as equilibrium constants.
• Kw = Ka * Kb: Relationship between Ka and Kb for conjugate acid-base pairs.
• In solutions of weak acids, equilibrium expressions can be used to solve for [H3O+], includes using simplifying assumptions and the quadratic equation when necessary.
• Degree/percent ionization: Fraction/percentage of weak acids and bases that ionize.
• Adding a common ion affects the equilibrium of a weak acid/base and its ionization.
• Buffer solutions: Resist pH changes, consist of weak acid/base and its conjugate salt.
• Buffers resist changes in pH upon addition of small amounts of acid or base.
• Buffer capacity: Amount of acid or base a buffer can neutralize before its pH changes significantly.
• pH range: Range of pH values where a buffer effectively resists pH changes.
• pH of a buffer solution can be calculated using concentrations of the weak acid/base and its conjugate salt, along with the Ka or Kb value.
• Henderson-Hasselbalch equation: Used to calculate the pH of buffer solutions.
• pKa and pKb relate to acid and base strength; smaller values indicate stronger acids/bases.
• Buffer systems are important in biological systems such as blood (bicarbonate buffer).
• Buffers maintain pH balance in the body; imbalance leads to acidosis or alkalosis.
• Acid: Donates a proton (H+) or accepts an electron pair. Increases the concentration of H3O+ ions in aqueous solutions.
• Base: Accepts a proton (H+) or donates an electron pair. Increases the concentration of OH- ions in aqueous solutions.
• Weak electrolyte: Partially ionizes/dissociates into ions in a solvent like water.
• Strong electrolyte: Completely ionizes/dissociates into ions in a solvent like water.
• Ionization: An atom/molecule gains charge by losing/gaining electrons, common when acids/bases dissolve in water.
• Dissociation: A compound separates into smaller particles (atoms, ions, radicals) reversibly.
• Acid dissociation constant (Ka): Extent to which a weak acid donates a proton to water, forming a hydronium ion and conjugate base.
• Base dissociation constant (Kb): Extent to which a weak base accepts a proton from water, forming a hydroxide ion and conjugate acid.
• Conjugate acid-base pair: Differ by one proton (H+); acid loses a proton to form conjugate base, base gains a proton to form conjugate acid.
• Hydronium ion (H3O+): Form of a proton in aqueous solution, associated with a water molecule.
• Hydroxide ion (OH-): Diatomic anion of one oxygen and one hydrogen atom with a negative charge.
• Ion-product constant for water (Kw): Equilibrium constant for water's auto-ionization (2H2O ⇌ H3O+ + OH-), equal to [H3O+][OH-] = 1.0 x 10^-14 at 25°C.
• Degree of ionization (α): Fraction of molecules of a weak acid/base that ionize in solution.
• Percent ionization: Percentage of molecules of a weak acid/base that ionize in solution (degree of ionization multiplied by 100%).
• Common ion effect: Equilibrium shifts when a salt containing a common ion is added to a weak electrolyte solution.
• Buffer solution: aqueous solution that resists pH changes by containing a weak acid/base and conjugate salt.
• pH range of a buffer: Values over which a buffer effectively resists pH changes, typically ±1 pH unit of the weak acid's pKa.
• Henderson-Hasselbalch equation for calculating buffer solution pH: pH = pKa + log([A-]/[HA]); for basic buffers: pOH = pKb + log([B+]/[BOH]).
• pKa: -log(Ka), lower value indicates stronger acid.
• pKb: -log(Kb), lower value indicates stronger base.
• Acidosis: Blood pH abnormally low (below 7.35).
• Alkalosis: Blood pH abnormally high (above 7.45).

Stoichiometry

• Stoichiometry: Calculating the amount of each chemical needed or produced in a reaction
• Balanced chemical equation: Atoms on the left (reactants) equal atoms on the right (products), adhering to the Law of Conservation of Mass.
• Mole (mol): Represents a large group of atoms/molecules. 1 mole = 6.022 × 10²³ particles (Avogadro's number).
• Molar mass: Weight of 1 mole of a substance in grams per mole (g/mol).
• Mole ratio: Ratio of reactants and products in a balanced equation.
• Limiting reactant: Reactant that runs out first, stopping the reaction, limiting product amount.
• Excess reactant: Reactant left over after the reaction.
• Theoretical yield: Maximum product amount from perfect reaction.
• Actual yield: Product amount made in lab, usually less than theoretical yield.
• Percent yield compares the actual yield to the theoretical yield. It’s calculated as: (Actual Yield / Theoretical Yield) × 100%.
• Concentration (Molarity, M): Substance amount in solution, moles per liter (mol/L).
• Dilution: Adding solvent to reduce solution concentration.
• Gravimetric stoichiometry: Using mass (grams) to calculate reactants and products.
• Volumetric stoichiometry: Using volume (liters) to calculate reactants and products, especially in solutions and gases.