Lesson 4 Chemical Equilibrium PDF

Chemical Equilibrium Lesson 4 Introduction Chemical equilibrium occurs when the rate at which the reactants convert into products is equal to the rate at which the products revert back into reactants. In simpler terms, it's a point where the forward and reverse reactions happen at the same...

Chemical Equilibrium Lesson 4 Introduction Chemical equilibrium occurs when the rate at which the reactants convert into products is equal to the rate at which the products revert back into reactants. In simpler terms, it's a point where the forward and reverse reactions happen at the same pace, so the concentrations of all substances remain constant. Reversible Reactions: A Two-Way Street Unlike one-way streets, reversible reactions can go both forward and backward. This means that the reactants can form products, and those products can break back down into the original reactants. Take, for instance, the reaction: 𝑁2(𝑔) + 3𝐻2 → 2𝑁𝐻3 Reversible Reactions: A Two-Way Street 𝑁2(𝑔) + 3𝐻2 → 2𝑁𝐻3 In this equation: The forward reaction: Nitrogen and hydrogen gases react to form ammonia. The reverse reaction: Ammonia can decompose back into nitrogen and hydrogen gases. In a reversible reaction, both these processes happen simultaneously. The rate at which the reactants turn into products is equal to the rate at which the products revert to reactants. This dynamic balance is what we call chemical equilibrium. Reversible Reactions: A Two-Way Street Example: 1. The Dissociation of Hydrogen Iodide 2𝐻𝐼(𝑔) 𝐻2 (𝑔) + 𝐼2 (𝑔) Forward Reaction: Hydrogen iodide decomposes into hydrogen and iodine gases. Reverse Reaction: Hydrogen and iodine gases combine to form hydrogen iodide. At Equilibrium: The rates of the forward and reverse reactions are equal, resulting in constant concentrations of 𝐻𝐼, 𝐻2 , and 𝐼2. Reversible Reactions: A Two-Way Street Example: 2. The Decomposition of Calcium Carbonate C𝑎𝐶𝑂3 𝑠 𝐶𝑎𝑂(𝑠) + 𝐶𝑂2 (𝑔) Forward Reaction: Calcium carbonate decomposes into calcium oxide and carbon dioxide gas upon heating. Reverse Reaction: Calcium oxide reacts with carbon dioxide gas to form calcium carbonate. At Equilibrium: The rates of formation and decomposition of C𝑎𝐶𝑂3 are balanced, with constant amounts of C𝑎𝐶𝑂3 , 𝐶𝑎𝑂, 𝐶𝑂2. Reversible Reactions: A Two-Way Street Example: 3. The Solubility of Gases in Liquids 𝐶𝑂2 𝑔 𝐶𝑂2 (𝑎𝑞) Forward Reaction: Carbon dioxide gas dissolves in water to form aqueous carbon dioxide. Reverse Reaction: Aqueous carbon dioxide escapes back into the gaseous phase. At Equilibrium: The rates of dissolving and escaping are equal, leading to a constant concentration of 𝐶𝑂2 in both phases. Reversible Reactions: A Two-Way Street Concentration: Changing the concentration of either reactants or Factors Influencing Reversible products can shift the equilibrium Reactions: position. For example, adding more reactants will push the equilibrium to form more products. Temperature: For endothermic Pressure: In reactions involving reactions, increasing the gases, increasing pressure by temperature shifts the equilibrium to decreasing volume will shift the produce more products, while for equilibrium towards the side with exothermic reactions, it shifts to fewer gas molecules. produce more reactants. Factors Affecting Rate of Reaction The rate of a chemical reaction is influenced by several key factors. Understanding these factors allows chemists to control and optimize reactions for various applications, from industrial manufacturing to biological processes. Factors Affecting Rate of Reaction 1. Concentration of Reactants Higher Concentration: Increasing the concentration of reactants typically increases the rate of reaction. This is because more reactant particles are available to collide and react. Example: In the reaction between hydrochloric acid (𝐻𝐶𝑙) and sodium thiosulfate (𝑁𝑎2 𝑆2 𝑂3 ), increasing the concentration of 𝐻𝐶𝑙 speeds up the reaction. Factors Affecting Rate of Reaction 2. Temperature Higher Temperature: Raising the temperature increases the kinetic energy of the particles, leading to more frequent and energetic collisions, thus increasing the reaction rate. Example: In the decomposition of hydrogen peroxide (𝐻2 𝑂2 ), heating the solution speeds up the release of oxygen gas. Factors Affecting Rate of Reaction 3. Surface Area of Reactants Larger Surface Area: Breaking down a solid reactant into smaller pieces increases its surface area, allowing more collisions to occur at the surface, which increases the reaction rate. Example: Finely powdered calcium carbonate reacts faster with hydrochloric acid than a single chunk of marble because of the larger surface area. Factors Affecting Rate of Reaction 4. Catalysts Presence of Catalysts: Catalysts are substances that increase the rate of a reaction without being consumed in the process. They work by providing an alternative reaction pathway with a lower activation energy. Example: In the decomposition of hydrogen peroxide (𝐻2 𝑂2 ), the enzyme catalase acts as a catalyst to speed up the reaction. Factors Affecting Rate of Reaction 5. Pressure (for Gaseous Reactions) Higher Pressure: Increasing the pressure in reactions involving gases increases the concentration of gas molecules, leading to more collisions and a higher reaction rate. Example: In the Haber process for synthesizing ammonia, increasing the pressure of nitrogen and hydrogen gases increases the rate of ammonia formation. Factors Affecting Rate of Reaction 6. Nature of the Reactants Reactive Substances: The intrinsic properties of the reactants, such as bond strengths and the presence of functional groups, affect the reaction rate. Reactive substances with weaker bonds or easily accessible reactive sites will react faster. Example: Sodium reacts rapidly with water due to its highly reactive nature, while iron reacts much more slowly. Factors Affecting Rate of Reaction Practical Applications: Industrial Manufacturing: Optimizing the concentration, temperature, and catalysts can make industrial processes more efficient and cost-effective. Pharmaceuticals: Understanding these factors helps in developing drugs that can be synthesized quickly and with high yield. Everyday Life: Cooking, preserving food, and even digestion in our bodies involve controlling reaction rates. The Law of Mass Action The Law of Mass Action states that the rate of a chemical reaction is proportional to the product of the masses (or concentrations) of the reacting substances, each raised to the power of their stoichiometric coefficients in the balanced chemical equation. In simpler terms, this law helps us understand how the concentrations of reactants influence the rate of a chemical reaction. The Law of Mass Action For a general reaction: 𝑎𝐴+𝑏𝐵→𝑐𝐶+𝑑𝐷 The rate of the forward reaction (R) can be expressed as: 𝑹 = 𝒌 𝑨 𝒂 [𝑩]𝒃 ] where: R is the reaction rate. k is the rate constant, a value specific to each reaction and temperature. [A] and [B] are the molar concentrations of reactants A and B. a and b are the stoichiometric coefficients of the reactants. The Law of Mass Action Explanation: This law is particularly useful in understanding the dynamics of chemical reactions and how changing For a general reaction: the concentration of reactants can affect the overall 𝑎𝐴+𝑏𝐵→𝑐𝐶+𝑑𝐷 reaction rate. The rate of the forward reaction (R) can be expressed as: Concentration Effect: If the concentration of one or more reactants increases, the reaction rate 𝑹 = 𝒌 𝑨 𝒂 [𝑩]𝒃 ] generally increases because there are more where: particles available to collide and react. R is the reaction rate. Stoichiometric Coefficients: The exponents in the rate expression correspond to the coefficients of k is the rate constant, a value specific to each reaction and temperature. the reactants in the balanced chemical equation, reflecting the proportional impact of each [A] and [B] are the molar concentrations of reactants A and B. reactant's concentration on the reaction rate. a and b are the stoichiometric coefficients of the reactants. The Law of Mass Action Application: The Law of Mass Action is also fundamental in describing chemical equilibrium. For a reversible reaction: 𝑎𝐴+𝑏𝐵→𝑐𝐶+𝑑𝐷 At equilibrium, the rate of the forward reaction equals the rate of the reverse reaction. The equilibrium constant (𝐾𝑒𝑞) can be defined using the concentrations of the reactants and products as follows: [𝑪] [𝑫]𝒄 𝒅 𝑲𝒆𝒒 = [𝑨]𝒂 [𝑩]𝒃 This equation allows chemists to predict the concentrations of reactants and products at equilibrium, providing valuable insight into the behavior of chemical systems. The Law of Mass Action Example: Consider the synthesis of ammonia: 𝑵𝟐(𝒈) + 𝟑𝑯𝟐(𝒈) 𝟐𝑵𝑯𝟑(𝒈) According to the Law of Mass Action, the equilibrium constant expression for this reaction would be: [𝑪]𝒄 [𝑫]𝒅 𝑲𝒆𝒒 = [𝑨]𝒂 [𝑩]𝒃 [𝑵𝑯𝟑 ]𝟐 𝑲𝒆𝒒 = [𝑵𝟐 ] [𝑯𝟐 ]𝟑 The Law of Chemical Equilibrium The Law of Chemical Equilibrium states that at a given temperature, a chemical system reaches a state where the ratio of the concentrations of products to the concentrations of reactants, each raised to the power of their respective stoichiometric coefficients, is constant. This constant ratio is known as the equilibrium constant (𝐾𝑒𝑞). For a general reversible reaction: 𝑎𝐴+𝑏𝐵 𝑐𝐶+𝑑𝐷 The Law of Chemical Equilibrium The equilibrium constant expression is: [𝑪]𝒄 [𝑫]𝒅 𝑲𝒆𝒒 = [𝑨]𝒂 [𝑩]𝒃 where: [A], [B], [C], and [D] are the equilibrium concentrations of the reactants and products. a, b, c, and d are the stoichiometric coefficients of the reactants and products. Explanation: At equilibrium, the rates of the forward and reverse reactions are equal, meaning that the concentrations of the reactants and products remain constant over time. The equilibrium constant (𝐾𝑒𝑞) provides a measure of the relative amounts of reactants and products at equilibrium. If 𝐾𝑒𝑞>1: The equilibrium position favors the formation of products, indicating that at equilibrium, the concentration of products is higher than that of reactants. If 𝐾𝑒𝑞