Decomposition vs Combustion reaction

I'm balancing chemical equations, and I came across the following:  

$$\ce{O2 + C4H9NH2 -> CO2 + H2O + N2}\tag{unbalanced}$$


My initial reaction was that it is a combustion reaction. According to my workbook, the definition is "an organic compound reacts with oxygen to form carbon dioxide and water." This equation seems to fit the bill. However, the solution states that in fact this is a decomposition reaction. In a YouTube video I saw, it said that decomposition reactions are a single reactant to multiple products. Where is the misunderstanding? Is it possible the workbook is wrong?




You're right, I'd consider that a decomposition reaction is primarily a single chemical species breaking into other chemical species. Here is the IUPAC Gold Book definition for [decomposition](https://goldbook.iupac.org/html/D/D01547.html). 


The IUPAC definition specifically mentions two phases, which implies that the decomposition is "irreversible." In other words, there isn't a significant backwards reaction. 


The molecular formula $\ce{C4H9NH2}$ could be a number of different aliphatic primary amines, but all of them are stable under ordinary conditions. In other words, you'd have to "burn" them in a combustion reaction with oxygen to get $\ce{CO2}$ and $\ce{H2O}$. In such a reaction I would absolutely not expect to get $\ce{N2}$ as a product. Rather I'd expect some sort of nitrogen oxide $\ce{NO\_x}$. Which one would depend on temperature, pressure, and the proportion of oxygen. 


So the unbalanced reaction would better be written as:
$$\ce{O2 + C4H9NH2 ->[\Delta] CO2 + H2O + NO}\_x$$


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