This question tests the understanding of thermodynamic terms, specifically exothermic and endothermic processes, in the context of dissolution. The key is to relate the observed temperature change (solution becoming colder) to the heat exchange between the system (dissolving salt) and the surroundings (water).
Understand the observation: The problem states that the dissolution of NH₄NO₃ in water makes the solution colder. This means the temperature of the surroundings (the water) decreases.
Relate temperature change to heat flow: A decrease in the temperature of the surroundings indicates that the system (the NH₄NO₃ dissolving) has absorbed heat from the surroundings. Heat flows from the warmer surroundings to the colder system.
Define endothermic and exothermic processes:
Apply definitions to the observation: Since the dissolution of NH₄NO₃ absorbs heat from the water (making it colder), the process is endothermic.
B) Endothermic — The dissolution of NH₄NO₃ in water absorbs heat from the water, leading to a decrease in the solution's temperature. This absorption of heat from the surroundings is the defining characteristic of an endothermic process.
A) Exothermic — An exothermic process releases heat to the surroundings, which would cause the solution's temperature to increase, not decrease. For example, dissolving NaOH in water is an exothermic process.
C) Has \(\Delta H = 0\) — If \(\Delta H = 0\), it means there is no net heat exchange with the surroundings, and thus no temperature change would be observed. Since the solution gets colder, there is a significant heat exchange, and \(\Delta H\) is not zero; it is positive.
D) Adiabatic — An adiabatic process is one where no heat is exchanged between the system and its surroundings (\(q = 0\)). In this case, heat is clearly exchanged (absorbed by the dissolving salt from the water), as evidenced by the temperature drop. Therefore, the process is not adiabatic.
The question asks to identify an exothermic reaction among the given options. An exothermic reaction is a chemical reaction that releases energy, typically in the form of heat or light, to its surroundings. This means the products have lower energy than the reactants. Endothermic reactions, on the other hand, absorb energy from the surroundings.
C) Burning of magnesium ribbon. This is a combustion reaction that releases a significant amount of heat and light, characteristic of an exothermic process.
This question asks about the sign convention for heat absorbed by a system, as defined by IUPAC (International Union of Pure and Applied Chemistry). Understanding this convention is fundamental in thermodynamics.
This convention is consistent with the first law of thermodynamics, which states that the change in internal energy (\(\Delta U\)) of a system is equal to the heat added to the system (\(q\)) plus the work done on the system (\(w\)): \(\Delta U = q + w\). For energy to increase the internal energy of the system, it must be positive when added.
B) q is positive. This is the standard IUPAC convention for heat absorbed by a system.
The question asks for the definition of Enthalpy (H). Enthalpy is a thermodynamic property that represents the total heat content of a system at constant pressure. It is defined in terms of internal energy, pressure, and volume.
A) U + PV — This is the standard definition of enthalpy. It represents the sum of the internal energy of the system and the product of its pressure and volume.
This question asks about the relationship between enthalpy change (\(\Delta H\)) and heat transfer under specific conditions. We need to recall the definition of enthalpy and its relation to the First Law of Thermodynamics.
Correct Option: A) qp (heat at constant pressure)
This question asks for the relationship between enthalpy change (ΔH) and internal energy change (ΔU) at constant pressure, specifically for reactions involving gases. We will start with the definition of enthalpy and relate it to internal energy and the work done due to volume change.
Definition of Enthalpy (H): Enthalpy is defined as the sum of the internal energy (U) and the product of pressure (P) and volume (V) of a system:
\[ H = U + PV \]Change in Enthalpy (ΔH): For a process occurring at constant pressure, the change in enthalpy (ΔH) can be expressed as:
\[ \Delta H = \Delta U + P\Delta V \]Work Done (W): For a process at constant pressure, the work done (W) by the system is given by:
\[ W = -P\Delta V \]Relating ΔH and ΔU for ideal gases: For reactions involving gases, if we assume ideal gas behavior, we can use the ideal gas law \(PV = nRT\). Therefore, for a change in volume due to a change in the number of moles of gas (\(\Delta n_g\)) at constant pressure and temperature, we have:
\[ P\Delta V = \Delta n_g RT \]Here, \(\Delta n_g\) represents the change in the number of moles of gaseous products minus the number of moles of gaseous reactants.
Substituting into the ΔH equation: Substitute the expression for \(P\Delta V\) from step 4 into the equation from step 2:
\[ \Delta H = \Delta U + \Delta n_g RT \]This equation relates the change in enthalpy to the change in internal energy, considering the work done due to the change in the number of moles of gaseous species.
Correct Option: D) The relationship \(\Delta H = \Delta U + \Delta n_g RT\) correctly describes the connection between enthalpy change and internal energy change for processes involving ideal gases at constant pressure and temperature, where \(\Delta n_g\) is the change in the number of moles of gaseous substances.