To analyze the change in the equation when you change \( 2O_2 \) to \( 3O_2 \), let's rewrite the modified equation:
\[ CH + 3O_2 \rightarrow CO_2 + 2H_2O \]
Now let's examine what happens:
- On the left side, we have 3 oxygen molecules instead of 2.
- The right side still has only 1 carbon dioxide (\( CO_2 \)) molecule and 2 water (\( H_2O \)) molecules.
Looking at the stoichiometry, the increase in oxygen should ideally allow for more complete combustion products. If we balance everything properly, changing \( 2O_2 \) to \( 3O_2 \) in a combustion reaction typically leads to more carbon dioxide produced if there is sufficient carbon in the hydrocarbon.
However, from what is written in the options, the outcomes you would typically assess from the combustion of carbon and changing oxygen ratio are:
- The number of \( CO_2 \) molecules does not automatically increase due to simply adding the extra \( O_2 \) unless balanced properly.
Thus, with the options provided, considering the fundamental combustion reaction:
- Nothing changes in the equation.
- There are now 3H2O molecules.
- There is an extra O2 molecule left over.
- There are now 3CO2 molecules.
The appropriate answer that summarizes the situation without causing an imbalance is:
There is an extra O2 molecule left over.
The combustion reaction has insufficient carbon and oxygen balance to use all the added \( O_2 \) based on initial hydrocarbons presented, leading to the conclusion that one additional \( O_2 \) remains unreacted for the assumed level of reaction progression.
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