Question 1: Bond energy is the amount of energy required to break one mole of a specific type of bond in a gaseous molecule. It is usually measured in kilojoules per mole (kJ/mol). When bonds are broken, energy is absorbed; when bonds are formed, energy is released. The difference between these energies determines whether a reaction is exothermic or endothermic.

Answer: Bond energy refers to the energy needed to break a mole of bonds in molecules. Breaking bonds absorbs energy, while forming bonds releases energy. The overall energy change depends on the balance between these two processes, indicating if the reaction is exothermic (releases heat) or endothermic (absorbs heat).

Question 1: Define exothermic and endothermic reactions.

Answer: Exothermic reactions release heat to the surroundings, while endothermic reactions absorb heat from the surroundings.

Question 2: Explain why bond energies are different for different types of bonds (e.g., C-H vs. C=O).

Answer: Bond energies vary because different bonds have different strengths due to factors like bond length, bond order, and the types of atoms involved. Double bonds (like C=O) are generally stronger and require more energy to break than single bonds (like C-H).

Question 1: Given the bond energies: C–H = 412 kJ/mol, C=O = 805 kJ/mol, and O=O = 498 kJ/mol, calculate whether the reaction: CH4 + 2 O2 → CO2 + 2 H2O is exothermic or endothermic based on bond energies.

Answer: Reaction is exothermic because the total energy released when forming bonds in CO2 and H2O exceeds the energy required to break bonds in CH4 and O2. Specific calculations: Bonds broken: 4 C–H bonds in methane + 2 O=O bonds in oxygen. Bonds formed: 2 C=O bonds in CO2 + 4 O–H bonds in water. Total energy absorbed = (4×412) + (2×498) = 1648 + 996 = 2644 kJ/mol. Total energy released = (2×805) + (4×463) [assuming O–H bonds are 463 kJ/mol] = 1610 + 1852 = 3462 kJ/mol. Since energy released > energy absorbed, the reaction is exothermic.

Question 1: Describe a simple experiment to measure bond energies indirectly, including the main variables involved and safety precautions.

Answer: A common approach is to burn a fuel (like methane) and measure the heat released using a calorimeter. Main variables: amount of fuel, temperature change, and specific heat capacity of water. Safety precautions include working in a well-ventilated area, wearing safety goggles, and handling open flames carefully.

Question 1: A laboratory experiment measures the following bond energies: for methane (C–H = 412 kJ/mol), oxygen (O=O = 498 kJ/mol), carbon dioxide (C=O = 805 kJ/mol), and water (O–H = 463 kJ/mol). The measured energy change for the combustion of 1 mole of methane is -890 kJ. Explain whether this reaction is exothermic or endothermic and justify your answer based on bond energies.

Answer: The reaction is exothermic because the measured energy change is negative (-890 kJ), indicating heat is released. This aligns with bond energy calculations showing that the energy released when bonds in products form exceeds the energy needed to break bonds in reactants.

Question 1: Explain how understanding bond energies helps in designing fuels and energy sources.

Answer: Knowing bond energies allows chemists to predict how much energy will be released or absorbed during combustion or other reactions, helping in designing efficient fuels that maximize energy output while minimizing waste and pollution.

Question 2: Describe a practical industry where bond energy calculations are crucial, and explain why.

Answer: In the petrochemical industry, bond energy calculations are essential for refining crude oil into useful products like gasoline and plastics. They help determine the energy required for breaking down complex molecules and optimizing reaction conditions for maximum efficiency and safety.