Understanding Chemical Equilibrium: Constants, Principles, and Factors
Chemical Equilibrium: Constants and Principles
Kc = [C]c [D]d / [A] [B]b = Q (not express equilibrium concentrations) <kc reaction to right> left = reaction toward equilibrium system
Kp = Pcc Pdd / Paa Pbb (gas only)
Kx = Xcc Xdd / Xaa Xbb
Kp = Kc (RT)W
Kx = Kp · PW
Pparcial = X (n species / total n) · Pt
Degree of dissociation (alpha #) # = x (n-differentiated) / c (initial total n) | A-2B (n (1 – #) / 2n #), A-B + C (n (1 – #) / n # / n #), 2A-B + C (n (1 – #) / n # / 2 / n # / 2), 2A + B-2C ( n (1-2 #) / n (1 – #) / 2n #)
ΔG = ΔH-TΔS =- RTlnKp | ln (KP1 / KP2) = (ΔH / R) (1/T2-1 / T1)
s =
ΔU = Q + W / Qp = ΔU + p · ΔV = ΔH / Qv = ΔU / Qv + Qp = ΔnRT / ΔH (S | G) = ΔHpro-ΔHreac / ΔG = ΔH-TΔS = 0 equilibrium <espontaneo> no spontaneous</espontaneo>
Mass Action Law and Equilibrium Constant
The equilibrium constant plays a crucial role in reversible reactions, similar to the limiting reagent in irreversible reactions, as it affects the concentration of both reactants and products at equilibrium.
For a reversible reaction aA + bB ⇌ cC + dD, the equilibrium constant Kc is defined.
The equilibrium constant Kc has no units and depends on temperature.
The concentrations of both products and reagents are expressed as molar concentrations.
- If Kc >> 1, most reactants have become products.
- If Kc << 1, most of the reagents remain unreacted.
Reaction Quotient
The reaction quotient (Q) is an expression similar to the equilibrium constant, providing information about the reaction’s progress at any given moment. Unlike the equilibrium constant, the reaction quotient uses instantaneous concentrations of reactants and products, not necessarily at equilibrium.
- If Q = Kc, the system is in equilibrium.
- If Q < Kc, the reaction will favor the forward direction, forming products.
- If Q > Kc, the reaction will favor the reverse direction, decomposing products back into reactants.
Value of Kp with Temperature
Van’t Hoff equation: ln (/)….
If the reaction is endothermic (ΔH > 0), an increase in temperature (T1 > T2) favors the reaction. Conversely, if the reaction is exothermic (ΔH < 0), an increase in temperature hinders the reaction.
Le Chatelier’s Principle
If a system in equilibrium is subjected to external changes (pressure, concentration, or temperature), the system will react in a way that counteracts the change and restores equilibrium.
Effect of Temperature
Temperature is the only variable that, in addition to shifting the equilibrium, also changes the equilibrium constant.
Increases in temperature favor endothermic processes, while decreases in temperature favor exothermic processes.
Effect of Pressure
A pressure change only affects equilibria in which Δn ≠ 0.
If pressure increases, the equilibrium shifts toward the side with fewer moles of gas, thus decreasing the volume. In heterogeneous systems, only consider the change in the number of moles of gases.
Effect of Concentration
Increasing the concentration of reactants shifts the equilibrium towards the formation of products. Conversely, increasing the concentration of products shifts the equilibrium towards the decomposition of products and an increased concentration of reactants.
Heterogeneous Equilibria
A saturated solution is one that cannot dissolve any more solute.
Solubility is the concentration of a saturated solution, expressed in grams per liter or moles per liter. Solubility depends on the nature of the substances and temperature.
A substance is considered soluble when the concentration of its saturated solution is greater than 0.01 moles per liter.
Factors Affecting Solubility
These factors primarily apply to ionic compounds in aqueous solutions.
- Temperature: An increase in temperature generally facilitates ion mobility and, consequently, solubility.
- Energy Factor: To dissolve an ionic compound, the lattice energy must be overcome. A compound dissolves when the hydration energy is greater than the lattice energy.
- Entropic Factor: The dissolution of a substance results in an increase in the disorder of the system and, therefore, an increase in entropy.
Factors Affecting the Solubility of Precipitates
A. Common Ion Effect: Increasing the concentration of a common ion decreases the solubility of the precipitate, as the solubility product (Ks) must remain constant.