Chemical Equilibrium for the DAT
/Learn key DAT concepts about chemical equilibrium, plus practice questions and answers
Everything you need to know about chemical equilibrium for the dat
Table of Contents
Part 1: Introduction to chemical equilibrium
Part 2: Reaction equilibrium
a) Equilibrium constant and reaction quotient
b) Le Chatelier’s principle
c) Bicarbonate buffer system
Part 3: Solubility and acid/base equilibria
a) Solubility equilibrium
b) Acid/base equilibrium
Part 4: High-yield terms
Part 5: Questions and answers
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Part 1: Introduction to chemical equilibrium
Chemical equilibrium involves dynamic processes where the rates of forward and reverse reactions are equal, leading to a stable concentration of reactants and products. The concept of equilibrium is pivotal in describing the behavior of chemical systems. Le Chatelier's principle, a fundamental component of chemical equilibria, posits that a system at equilibrium responds to changes in concentration, pressure, or temperature by adjusting its conditions to counteract the disturbance. These topics and more are frequently tested on the DAT. This guide will cover everything you need to know about chemical equilibrium for the exam, including examples with equilibrium constants. Pay attention to the high-yield bold terms and gauge your knowledge with practice questions and answers at the end of the guide.
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Part 2: Reaction equilibrium
a) Equilibrium constant and reaction quotient
Reversible reactions are capable of advancing both forward and backward. This is known as dynamic equilibrium, where the rates of the forward and reverse reactions are equal. Think of this equilibrium as a balanced see-saw with reactants and products on opposing sides. This equilibrium state doesn't imply a lack of reaction progress; instead, it signifies an equilibrium where the rates of the two opposing reactions are identical.
The law of mass action articulates the ratio of reactants to products at equilibrium, quantified by the equilibrium constant (Keq). Represented by the reaction aA + bB ⇋ cC + dD, the Keq expression quantifies the specific point where equilibrium is attained in the dynamic interplay of reactants transforming into products and vice versa.
In the given expression, the square brackets ([ ]) represent the concentration of a specific chemical species, while lowercase letters (a, b, c, d) denote stoichiometric coefficients. Environmental factors like temperature and the presence of catalysts can influence a reaction's equilibrium position, thereby potentially altering Keq. The capital letter "K" in chemistry commonly signifies an equilibrium constant, applicable to various constants such as the acid dissociation constant (Ka), base dissociation constant (Kb), and even the Michaelis constant (Km). For instance, consider the acid dissociation constant, Ka, which reflects the strength of an acid. A higher Ka value indicates a stronger acid. Exploring the Ka expression for the acid dissociation, illustrated as HA ⇋ H+ + A-, provides insight into the quantitative aspects of this equilibrium.
Since stronger acids dissociate more readily, the dissociation reaction of a strong acid favors the products. Thus, the resulting Ka value is higher.
Another important variable is the reaction quotient, Q. The definition of Q is virtually identical to that of Keq:
While Keq relies on the concentrations of species at equilibrium, Q provides information about the current status of the reaction. If Q equals Keq, the reaction is at equilibrium. In contrast, if Q is less than Keq, the reaction is not yet in equilibrium, prompting it to proceed forward towards the products. Conversely, if Q surpasses Keq, the reaction has surpassed equilibrium, leading it to shift in the reverse direction towards the reactants.
b) Le Chatelier’s principle
Le Chatelier's principle illustrates the adjustments in a reaction's favorability towards products or reactants amid varying environmental conditions. This principle explains how equilibrium position shifts. When a system encounters stress, it adapts to alleviate that stress and restore equilibrium. Stress may manifest as alterations in chemical concentrations or environmental factors like temperature, pressure, or volume. Notable considerations include:
If reactant concentration increases or product concentration decreases, a chemical reaction will shift to the right.
Altering temperature impacts the shift direction based on heat's role as a reactant or product. In endothermic reactions, where heat acts as a reactant, an increase in temperature shifts the reaction right. The opposite is true for exothermic reactions, where heat acts as a product.
Adjusting pressure results in a shift direction influenced by gas moles. Increased pressure directs the reaction towards fewer gas moles, while decreased pressure steers it towards more gas moles. For instance, if reactants have more moles than products, increased pressure shifts the reaction right. This shift aligns with the ideal gas law (PV = nRT), where pressure and moles correlate, leading the system to counterbalance by shifting towards fewer moles.
c) Bicarbonate buffer system
The concepts previously examined form the basis for understanding one of the crucial systems in your body: the bicarbonate buffer system. The operation of the bicarbonate buffer system can be described through the following chemical equation:
The bicarbonate buffer system assumes a vital role in maintaining the blood's pH balance. This system operates as a meticulously regulated equilibrium, following Le Chatelier's principle. When hydrogen ion concentration rises, causing a decrease in pH, the body responds by elevating the respiration rate to expel carbon dioxide. This action shifts the reaction leftward, diminishing the hydrogen ion concentration and elevating the pH. Conversely, in response to high pH or a low hydrogen ion concentration, respiration is adjusted downward to retain carbon dioxide within the body. This alteration shifts the equilibrium position of the reaction to the right, increasing the hydrogen ion concentration and lowering the pH.
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