A constant volume and mass of helium gas at 47°C is heated so that the pressure of the gas doubles. What is the new temperature of the gas in Celsius
degrees?
(P1/T1) = (P2/T2)
Don't forget to convert T to Kelvin.
367°C ? where are the values coming from? (P1/T1) = (P2/T2) ? all I get is the original 47°C. 47°C+273=K, K=320...I'm lost
To solve this problem, we can use the ideal gas law equation, which states that PV = nRT, where P is the pressure, V is the volume, n is the number of moles of gas, R is the ideal gas constant, and T is the temperature.
Since the volume remains constant, V can be considered as a constant value.
We are given that the mass of helium gas and its volume remain constant. Mass remains constant, and volume is constant, so we can say that the number of moles of gas, n, is constant.
Therefore, we can write the equation as P1/T1 = P2/T2, where P1 is the initial pressure, T1 is the initial temperature, P2 is the final pressure, and T2 is the final temperature.
Given that the initial temperature is 47°C and the pressure doubles, we know that P2 = 2P1.
Plugging in the values, we get P1/T1 = (2P1)/T2.
Now we can solve for T2:
P1/T1 = (2P1)/T2
T2 = (2P1 * T1) / P1
T2 = 2 * T1
T2 = 2 * 47°C
T2 = 94°C
Therefore, the new temperature of the helium gas in Celsius degrees is 94°C.