Temperature rise without heat energy?

[vcsharp2003]

Homework Statement

Is it possible to raise the temperature of a gaseous mass without having any heat energy flow into this gaseous mass?

Relevant Equations

$\Delta Q = cm \Delta T$ where $\Delta Q$ is the quantity of heat required to produce a temperature change $\Delta T$ in a mass $m$ of substance having a specific heat of $c$

If I look at the specific heat equation mentioned, then I would be inclined to think that without heat energy being added to the gaseous mass its temperature cannot rise. But, if some form of energy like chemical energy in gaseous mass could be directly converted to internal energy of the same gaseous mass without involving heat, then that could also raise the temperature; whether such a transformation is even possible is something I am not sure of.

 

[Steve4Physics]

Homework Statement:: Is it possible to raise the temperature of a gaseous mass without having any heat energy flow into this gaseous mass?
Relevant Equations:: $\Delta Q = cm \Delta T$ where $\Delta Q$ is the quantity of heat required to produce a temperature change $\Delta T$ in a mass $m$ of substance having a specific heat of $c$

If I look at the specific heat equation mentioned, then I would be inclined to think that without heat energy being added to the gaseous mass its temperature cannot rise. But, if some form of energy like chemical energy in gaseous mass could be directly converted to internal energy of the same gaseous mass without involving heat, then that could also raise the temperature; whether such a transformation is even possible is something I am not sure of.

If you compress a gas, the temperature can rise. E.g. look-up adiabatic compression.

The equation $\Delta Q = cm \Delta T$ needs to be used with care when dealing with a gas. You use different values for $c$ depending on the conditions.

The gas could be supplied with thermal energy ($\Delta Q$) while the pressure is kept constant, which means volume is changing. In this case we must use the specific heat capacity at constant pressure ($c_P$).

Or the gas could be supplied with thermal energy while keeping its volume constant, which means its pressure is changing. In this case we must use the specific heat capacity at constant volume ($c_V$).

These two values of specific heat capacity, $c_P$ and $c_V$, are different

(Of course, both pressure and vo;lume might be simultaneously changing, but that's a different problem.)

 

[Chestermiller]

The problem is that the OP is expressing the temperature change and heat capacity in terms of heat. This is the way it was done in freshman physics, but it is no longer correct in thermodynamics. In thermodynamics, the temperature change is related to the change in internal energy U, not heat Q. For an ideal gas, $$\Delta U=mC_v\Delta T$$and, from the first law of thermodynamics, $$\Delta U=mC_v\Delta T=Q-W$$Even if no heat is involved Q=0 (adiabatic system), the temperature can still change if work is involved: $$\Delta U=mC_v\Delta T=-W$$The old freshman physics version of the relationship is obtained only when no work is involved: $$\Delta U=mC_v\Delta T=Q$$(or at constant pressure, and we are using the heat capacity at constant pressure).