The combustion of methane is represented by the equation: \[ \text{CH}_4(g) + 2\text{O}_2(g) \rightarrow \text{CO}_2(g) + 2\text{H}_2\text{O}(l) \] Given: $\Delta H_{\text{CH}_4} = -75 \, \text{kJ/mol}$ $\Delta H_{\text{CO}_2} = -393.5 \, \text{kJ/mol}$ $\Delta H_{\text{H}_2\text{O}} = -285.8 \, \text{kJ/mol}$ What is the enthalpy change ($\Delta H$) for the combustion of 1 mole of methane?
At 700 K, the equilibrium constant $K_e$ for the reaction $ \text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) $ is 0.2 mol L$^{-2}$. What is the value of $K$ for the reverse reaction?
The enthalpy of combustion of methane is 890 kJ/mol. How much heat is released when 8 g of methane is burned completely? (Molar mass of CH\(_4\) = 16 g/mol)
Among the following, choose the ones with an equal number of atoms.
Choose the correct answer from the options given below:
Regarding the molecular orbital (MO) energy levels for homonuclear diatomic molecules, the INCORRECT statement(s) is (are):
For the reaction sequence given below, the correct statement(s) is (are): (In the options, X is any atom other than carbon and hydrogen, and it is different in P, Q, and R.)
