2.2 Le Chatelier's Principle and factors affecting equilibrium
3 min read•july 22, 2024
Chemical equilibrium is a dynamic state where forward and reverse reactions occur at equal rates. Le Chatelier's Principle explains how systems respond to disturbances, shifting to counteract changes in concentration, , or .
Gibbs free energy connects thermodynamics to equilibrium. It determines reaction spontaneity and direction, linking to the equilibrium constant. Understanding these concepts helps predict and control chemical reactions in various applications.
Le Chatelier's Principle and Equilibrium
Le Chatelier's principle and equilibrium shifts
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- Le Chatelier's Principle states that when a system at equilibrium is disturbed by a change in concentration, pressure, volume, or temperature, the system will shift its equilibrium position to counteract the disturbance and re-establish equilibrium (Haber process for ammonia synthesis)
- Changes in concentration affect equilibrium by shifting it towards the side that will reduce the
- Increasing reactant concentration shifts equilibrium towards products to consume the added reactants (adding more CO in the water-gas shift reaction)
- Decreasing reactant concentration shifts equilibrium towards reactants to replenish the consumed reactants (removing H2 in the Haber process)
- Increasing product concentration shifts equilibrium towards reactants to consume the added products (adding more CO2 in the water-gas shift reaction)
- Decreasing product concentration shifts equilibrium towards products to replenish the consumed products (removing NH3 in the Haber process)
- Changes in pressure or volume affect equilibrium in gaseous systems by shifting it towards the side with fewer moles of gas to minimize
- Increasing pressure (decreasing volume) shifts equilibrium towards the side with fewer gas molecules (N2 + 3H2 ⇌ 2NH3)
- Decreasing pressure (increasing volume) shifts equilibrium towards the side with more gas molecules (2SO2 + O2 ⇌ 2SO3)
- Changes in temperature affect equilibrium depending on whether the reaction is exothermic or endothermic
- Increasing temperature shifts equilibrium towards the endothermic direction to absorb the added heat (N2 + O2 ⇌ 2NO)
- Decreasing temperature shifts equilibrium towards the exothermic direction to release heat (2NO2 ⇌ N2O4)
Catalysts in reaction rates
- Catalysts accelerate the rate of a reaction by lowering the activation energy barrier without being consumed in the reaction (enzymes in biological systems, catalytic converters in automobiles)
- Catalysts increase the rate of both forward and reverse reactions equally, thus not altering the equilibrium constant () or the equilibrium position (Haber process using iron catalyst)
- Catalysts help the system reach equilibrium faster without shifting the equilibrium towards reactants or products (catalytic hydrogenation of vegetable oils)
Reactant and product effects on equilibrium
- Adding reactants increases their concentration, shifting the equilibrium towards products to counteract the disturbance (adding more CO in the water-gas shift reaction)
- Removing reactants decreases their concentration, shifting the equilibrium towards reactants to replenish the consumed species (removing H2 in the Haber process)
- Adding products increases their concentration, shifting the equilibrium towards reactants to counteract the disturbance (adding more CO2 in the water-gas shift reaction)
- Removing products decreases their concentration, shifting the equilibrium towards products to replenish the consumed species (removing NH3 in the Haber process)
Gibbs Free Energy and Equilibrium
Gibbs free energy and equilibrium position
- Gibbs free energy () is a thermodynamic quantity that determines the spontaneity and direction of a reaction at constant temperature and pressure
- The relationship between and the equilibrium constant () is given by the equation: , where is the gas constant and is the absolute temperature
- indicates a spontaneous reaction, with equilibrium favoring product formation (formation of rust, Fe2O3)
- indicates a non-spontaneous reaction, with equilibrium favoring reactant formation (decomposition of water into H2 and O2)
- indicates a system at equilibrium, with constant concentrations of reactants and products (saturated salt solution)
- The magnitude of determines the extent of the reaction: a more negative results in a larger equilibrium constant and a greater proportion of products at equilibrium (formation of NaCl from Na and Cl2)
Key Terms to Review (18)
Closed system: A closed system is a physical system that does not exchange matter with its surroundings but can exchange energy. This concept is crucial for understanding chemical reactions and equilibrium, as it allows for the analysis of how changes in energy or conditions affect the system's behavior without the influence of outside matter.
Concentration change: Concentration change refers to the alteration in the amount of reactants or products in a chemical system at equilibrium, impacting the system's position according to Le Chatelier's Principle. When the concentration of either reactants or products is modified, the equilibrium will shift to counteract that change, either favoring the forward or reverse reaction to restore balance. Understanding this concept is crucial for calculating new equilibrium concentrations when disturbances occur.
Dynamic Equilibrium: Dynamic equilibrium is a state in a reversible reaction where the rates of the forward and reverse reactions are equal, leading to constant concentrations of reactants and products over time. This concept is crucial for understanding how systems respond to changes in conditions and is essential for grasping precipitation reactions, solubility product constants, and the influence of external factors on chemical systems.
Exothermic reaction: An exothermic reaction is a chemical process that releases energy, usually in the form of heat, to its surroundings. This type of reaction typically results in a decrease in the enthalpy of the system as reactants transform into products, releasing energy that can be felt as warmth. Understanding exothermic reactions is crucial for analyzing energy changes in chemical reactions and how these reactions affect equilibrium states.
Gas phase: The gas phase refers to the state of matter where substances exist as gases, characterized by low density and high energy. In this phase, molecules are widely spaced apart and move freely, which plays a crucial role in reactions and equilibrium dynamics.
Keq: The equilibrium constant, denoted as keq, is a value that expresses the ratio of the concentrations of products to the concentrations of reactants at equilibrium for a given chemical reaction. This constant helps in predicting the direction in which a reaction will proceed and is influenced by temperature, allowing us to understand how changes in conditions can affect the position of equilibrium.
Law of mass action: The law of mass action states that the rate of a chemical reaction is proportional to the product of the concentrations of the reactants, each raised to a power equal to their stoichiometric coefficients. This principle is fundamental in understanding how changes in concentration can shift the position of equilibrium in a chemical reaction.
Le Chatelier's Principle in Industry: Le Chatelier's Principle states that if an external change is applied to a system at equilibrium, the system will adjust to counteract that change and restore a new equilibrium. In industry, this principle is crucial for optimizing chemical reactions, ensuring maximum product yield, and enhancing reaction efficiency by manipulating conditions such as concentration, temperature, and pressure.
Liquid phase: The liquid phase is one of the states of matter characterized by a definite volume but no definite shape, allowing it to take the shape of its container. In this state, the molecules are closely packed but still have enough energy to move freely, which plays a crucial role in chemical reactions and equilibrium processes.
Open System: An open system is a type of thermodynamic system that can exchange both matter and energy with its surroundings. This means that in an open system, substances can enter and leave, which affects the overall state and equilibrium of the system. The interaction between an open system and its environment is crucial for understanding how changes impact chemical reactions and the position of equilibrium.
Pressure: Pressure is defined as the force exerted per unit area on a surface, commonly measured in units such as atmospheres (atm), pascals (Pa), or mmHg. It plays a crucial role in influencing chemical reactions, state changes, and equilibria by affecting how particles collide and interact, which can ultimately drive the direction of chemical processes and affect their thermodynamic properties.
Pressure change: Pressure change refers to the variation in the pressure of a system, which can influence the position of equilibrium in a chemical reaction. When the pressure of a system is altered, particularly in reactions involving gases, the system will respond by shifting the equilibrium position to counteract that change. This response is outlined by a principle that predicts how systems at equilibrium behave when external conditions are modified.
Reversible reactions: Reversible reactions are chemical processes that can proceed in both the forward and reverse directions, allowing the reactants to form products and the products to convert back into reactants. This dynamic balance is essential for understanding how systems respond to changes and maintain equilibrium. The concept of reversible reactions is critical when discussing factors that affect equilibrium and calculating concentrations of substances at equilibrium.
Solid phase: The solid phase is one of the primary states of matter characterized by its fixed shape and volume, where particles are closely packed together in a structured arrangement. In this state, particles vibrate in place but do not move freely, resulting in a rigid structure. The solid phase plays a crucial role in understanding the behavior of substances under various conditions, especially when discussing how changes in temperature or pressure can affect equilibrium.
Stress on Equilibrium: Stress on equilibrium refers to any change imposed on a chemical system at equilibrium that disrupts its balance, causing the system to respond in a way that re-establishes a new equilibrium. This concept is crucial as it explains how changes in concentration, temperature, or pressure can affect the direction of the reaction and ultimately the concentrations of reactants and products.
Synthesis of Ammonia: The synthesis of ammonia is the chemical process by which nitrogen gas ($$N_2$$) and hydrogen gas ($$H_2$$) react to form ammonia ($$NH_3$$), typically represented by the equation: $$N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$$. This process is crucial in the production of fertilizers and other chemicals, and it exemplifies the principles of chemical equilibrium and the factors that can influence this balance.
Temperature: Temperature is a measure of the average kinetic energy of particles in a substance, which directly influences how substances interact and react with one another. It plays a crucial role in determining reaction rates, the spontaneity of reactions, equilibrium positions, and the behavior of acids and bases.
Temperature change: Temperature change refers to the variation in thermal energy of a system, which can affect the physical and chemical processes occurring within it. In the context of equilibrium, temperature changes can shift the balance between reactants and products, influencing the direction of a reaction. Understanding how temperature impacts equilibrium helps in predicting how a system will respond to external changes.