I am having such a hard time approaching this problem. It would be great if you can help me.
1.0 mole of CH4 react with 1.0 mole of O2 in a sealed container. (a) When the reaction is complete, what is the container? (b) How many moles of water are produced?
CH4 + 2O2 ---> CO2 + 2H2O
a) CH4 + O2 ---> CO2 + H2O
Looks like the 1.0 mole of O2 is the limiting reagent. Only 0.5 moles of CH4 can react with it. So, when things are done, you are left with
0.5 mole CH4 (left over)
0.5 mole CO2
1.0 mole H2O
What mass of natural gas (CH4) must you burn to emit 263kJ of heat?
CH4(g)+2O2(g)→CO2(g)+2H2O(g)ΔH∘rxn=−802.3kJ
To determine the container after the reaction is complete, remember that gases, like CH4 and O2, are typically the ones that occupy containers. In this particular reaction, CH4 and O2 are the reactants, and CO2 and H2O are the products.
Since CH4 and O2 are the reactants, they are the ones present in the container at the start of the reaction. After the reaction is complete, the CH4 and O2 have been consumed to form the products CO2 and H2O. Therefore, the container will contain CO2 and H2O as the final substances.
b) To find out how many moles of water are produced, refer to the balanced chemical equation provided:
CH4 + 2O2 ---> CO2 + 2H2O
According to the equation, 1 mole of CH4 reacts with 2 moles of O2 to produce 2 moles of H2O.
Since 1 mole of CH4 and 1 mole of O2 are used in the reaction, we can conclude that the number of moles of water produced will be the same as the number of moles of O2 used. In this case, it is 1 mole.
Therefore, 1 mole of water is produced in the reaction.