a reference gives the molar enthalpy of combustion for methane as -803 kj/mol. what is the minimum mass of methane must be burned to warm 4.00 L of water from 22.4 C to 87.6 C, assuming no heat loss?

7 years ago

Answered By Vincent E

If we assume no heat loss, then the heat released from the exothermic combustion of methane all goes towards heating the water. We can say: $delta\left(H\right)=-Q$delta(H)=−Q . Where delta(H) is the enthalpy change of the reaction, and Q is the heat gained by the water.

We can calculate Q using the info given in the problem (remember that 1g = 1mL for water!), and then use $delta\left(H\right)=-Q$delta(H)=−Q to calculate the corresponding enthalpy change. Then we can use equation 1 to calculate the moles of methane. Then using $n=\frac{m}{M}$n=mM , (M being the molar mass of methane, which you can calculate using a periodic table) we can find the mass of methane.

7 years ago

Answered By Vincent E

For the answer above, Hc = molar enthalpy of combustion for methane.

7 years ago

## Answered By Vincent E

If we assume no heat loss, then the heat released from the exothermic combustion of methane all goes towards heating the water. We can say: $delta\left(H\right)=-Q$delta(H)=−Q . Where delta(H) is the enthalpy change of the reaction, and Q is the heat gained by the water.

We then have two equations:

1. $delta\left(H\right)=n\cdot delta\left(Hc\right)$delta(H)=n·delta(Hc)

2. $Q=mC\cdot delta\left(T\right)$Q=mC·delta(T)

We can calculate Q using the info given in the problem (remember that 1g = 1mL for water!), and then use $delta\left(H\right)=-Q$delta(H)=−Q to calculate the corresponding enthalpy change. Then we can use equation 1 to calculate the moles of methane. Then using $n=\frac{m}{M}$n=mM , (M being the molar mass of methane, which you can calculate using a periodic table) we can find the mass of methane.

7 years ago

## Answered By Vincent E

For the answer above, Hc = molar enthalpy of combustion for methane.