1. Define equilibrium and explain the concept of dynamic equilibrium.
Answer: Equilibrium is a state of balance in a system where opposing processes occur at equal rates. Dynamic equilibrium refers to a situation where the rates of the forward and reverse reactions are equal, resulting in no net change in the concentrations of reactants and products over time.
2. What are the types of equilibria involving physical processes?
Answer: Equilibria involving physical processes include solid-liquid, liquid-gas, gas-gas, and solid-gas equilibria, as well as Henry’s law for the equilibrium between a gas and a liquid.
3. Describe Henry’s law and its significance.
Answer: Henry’s law states that the concentration of a gas dissolved in a liquid is directly proportional to the partial pressure of the gas above the liquid. It is significant in understanding gas solubility in liquids, such as the dissolution of oxygen in water.
4. What are the general characteristics of equilibrium involving physical processes?
Answer: In equilibrium involving physical processes, the rates of forward and reverse processes are equal, the concentrations of reactants and products remain constant over time, and the system exhibits dynamic equilibrium.
5. Discuss the equilibria involving chemical processes.
Answer: Equilibria involving chemical processes occur when the rates of the forward and reverse reactions are equal, resulting in no net change in the concentrations of reactants and products over time.
6. State the law of chemical equilibrium and its implications.
Answer: The law of chemical equilibrium states that at a given temperature, the ratio of the product of the concentrations of the products to the product of the concentrations of the reactants is constant. This equilibrium constant (K) provides information about the extent of a chemical reaction at equilibrium.
7. Define equilibrium constants Kp and Kc and explain their significance.
Answer: Kp and Kc are equilibrium constants expressed in terms of partial pressures and concentrations, respectively. They indicate the position of equilibrium and provide information about the relative amounts of reactants and products at equilibrium.
8. Discuss the significance of ΔG and ΔG∘ in chemical equilibrium.
Answer: ΔG is the Gibbs free energy change of a reaction and indicates whether a reaction is spontaneous (ΔG<0) or non-spontaneous (ΔG>0) at a given temperature and pressure. ΔG∘ is the standard Gibbs free energy change and helps predict the direction of a reaction at standard conditions.
9. What factors affect the equilibrium concentrations of reactants and products?
Answer: Factors affecting equilibrium concentrations include changes in temperature, pressure, concentration of reactants or products, and the presence of a catalyst.
10. Explain Le Chatelier’s principle and its application to chemical equilibrium.
Answer: Le Chatelier’s principle states that if a system at equilibrium is subjected to a change, it will adjust to counteract the change and restore equilibrium. This principle helps predict the direction of shift in equilibrium in response to changes in temperature, pressure, or concentration.
11. Differentiate between weak and strong electrolytes in ionic equilibrium.
Answer: Weak electrolytes partially ionize in solution, while strong electrolytes completely ionize. Weak acids and bases are examples of weak electrolytes, while strong acids and bases dissociate completely in solution.
12. Discuss the ionization of electrolytes and its significance in ionic equilibrium.
Answer: The ionization of electrolytes refers to the dissociation of ions in solution. It is significant in determining the conductance and chemical behavior of electrolyte solutions.
13. Explain the concepts of acids and bases according to Arrhenius, Bronsted-Lowry, and Lewis.
Answer: According to Arrhenius, acids produce hydrogen ions (H+) in solution, while bases produce hydroxide ions (OH−). Bronsted-Lowry acids donate protons, while bases accept protons. Lewis acids accept electron pairs, while bases donate electron pairs.
14. Describe acid-base equilibria and ionization constants in ionic equilibrium.
Answer: Acid-base equilibria involve the transfer of protons between acids and bases. Ionization constants, such as Ka for acids and Kb for bases, indicate the extent of ionization of weak acids and bases in solution.
15. What is the pH scale and how is it calculated?
Answer: The pH scale is a measure of the acidity or basicity of a solution, ranging from 0 to 14. It is calculated as the negative logarithm of the hydrogen ion concentration ([H+]) in moles per liter:
16. Explain the common ion effect in ionic equilibrium.
Answer: The common ion effect is the suppression of ionization of a weak electrolyte by the addition of another electrolyte containing a common ion. It reduces the solubility of sparingly soluble salts and affects the pH of solutions.
17. Discuss the hydrolysis of salts and its impact on solution pH.
Answer: Hydrolysis of salts occurs when ions in a salt react with water to produce acidic or basic solutions. It can lead to changes in solution pH depending on the nature of the ions involved.
18. What factors affect the solubility of sparingly soluble salts and solubility products?
Answer: Factors affecting solubility include temperature, pressure, and the presence of common ions. Solubility products (Ksp) indicate the extent of dissolution of sparingly soluble salts in solution.
19. Define buffer solutions and discuss their significance in maintaining pH.
Answer: Buffer solutions resist changes in pH when small amounts of acid or base are added. They consist of a weak acid and its conjugate base or a conjugate base or a weak base and its conjugate acid. Buffers are important in maintaining the pH of biological systems and in various chemical processes where pH stability is required.
20. How does the solubility product (Ksp) relate to the solubility of a sparingly soluble salt?
Answer: The solubility product (Ksp) is the equilibrium constant for the dissolution of a sparingly soluble salt in water. It indicates the extent to which the salt dissociates into its ions in solution, thereby determining its solubility. Higher Ksp values correspond to greater solubility of the salt in water.