Exam 2 lecture 6 part 2

we can practice calculating the standard enthalpy for the combustion of glucose. By applying the formula, using known standard enthalpies of formation for each compound, we can determine the overall energy change. For example, if we consider specific values like -893.51 and -285.83, the calculation would lead to a result of approximately -2808 kJ/mol. This negative value indicates an exothermic reaction, meaning energy is released. Next, we explore bond energies. We understand that the bond energies for most common bonds are already known and tabulated. We can use these values to estimate the enthalpy change, or Delta H, for a reaction. This estimation relies on two fundamental processes: bond breaking and bond forming. Bond breaking always requires energy, making it an endothermic process where Delta H is positive. Conversely, bond forming always releases energy, which is an exothermic process where Delta H is negative. Therefore, the total enthalpy change for a reaction is the sum of the energy needed to break bonds and the energy released from forming new ones. Essentially, we approximate Delta H for the reaction as the sum of the energy of broken bonds minus the sum of the energy of formed bonds. Finally, we consider the determination of lattice enthalpy through the Born-Haber Cycle. It's important to remember that lattice enthalpy of solids cannot be directly measured. However, we can utilize measurable enthalpies and the fact that enthalpy is a state function to determine lattice enthalpy indirectly. For example, charging separation in ionic solids is stabilized by ion-pair interactions, which is known as lattice interaction. We can calculate the change in entropy when the temperature is constant, using the formula ΔS = q_rev/T, where q_rev is the reversible heat transferred. This understanding of entropy helps us predict the direction of spontaneous change, offering insights into why certain reactions proceed naturally while others do not We can see the various phases of matter on this diagram, showing how pressure and temperature influence whether a substance exists as a solid, liquid, or gas. The triple point, also known as the critical point, is where all three phases coexist in equilibrium. We also see the boiling point at 100 degrees Celsius and 1 atm, and the melting point at 0 degrees Celsius. Phase changes are always accompanied by a change in the energy of a system. For instance, processes like vaporization and sublimation require heat, while condensation and freezing release heat. The enthalpy change for a reverse reaction is simply the negative of the enthalpy change for the forward reaction. For example, the enthalpy of fusion is the negative of the enthalpy of freezing. This is because enthalpy is a state function, meaning enthalpy changes can be added to obtain the value for an overall process. We also observe that higher boiling points are associated with higher enthalpies of vaporization.