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What is the basic difference between 'q' and 'ΔH' in thermochemistry? I get confused between them!
Is there any criteria for ΔH to become equal to q?
Is there any criteria for ΔH to become equal to q?
Abdul Quadeer said:What is the basic difference between 'q' and 'ΔH' in thermochemistry? I get confused between them! Is there any criteria for ΔH to become equal to q?
Q is the symbol used to represent heat, which is a form of energy. It is measured in Joules (J) or calories (cal). ΔH (delta H) represents the change in enthalpy, which is the total heat content of a system. It is also measured in Joules or calories.
Q and ΔH are related through the equation Q = ΔH + ΔnRT, where Δn is the change in the number of moles of gas molecules, R is the gas constant, and T is the temperature in Kelvin. This equation shows that q is a component of ΔH and that they are both influenced by the same factors.
No, q and ΔH can be positive, negative, or even zero in thermochemistry. A positive value for q or ΔH indicates that the system is gaining heat or increasing in enthalpy, respectively. A negative value indicates that the system is losing heat or decreasing in enthalpy. A value of zero means that there is no change in heat or enthalpy.
The values of q and ΔH can tell us whether a reaction is endothermic or exothermic. An endothermic reaction absorbs heat from the surroundings, resulting in a positive value for q and ΔH. An exothermic reaction releases heat to the surroundings, resulting in a negative value for q and ΔH.
No, q and ΔH cannot be used interchangeably. Q represents the heat exchanged between a system and its surroundings, while ΔH represents the overall change in enthalpy of a system. They have different units and are calculated using different equations, so they cannot be substituted for each other in calculations.