Determining order of reaction when concentration of reactant is not given Suppose reactant $\ce{A}$ decomposes in an aqueous solution to give $\ce{P}, \ce{Q}$:
  $$
\ce{A(aq) -&gt; P(aq) + Q(aq)}
$$
  
  Suppose further that $[\ce{A}]_0=0.42\text{M}$, and that $[\ce{Q}]$ was measured as follows (in M):
  
  \begin{align}
t&amp;=0 &amp;[\ce{Q}]&amp;=0 \\
t&amp;=100 \quad &amp;[\ce{Q}]&amp;=0.101\\
t&amp;=200 \quad &amp;[\ce{Q}]&amp;=0.165\\
t&amp;=300 \quad &amp;[\ce{Q}]&amp;=0.206\\
t&amp;=400 \quad &amp;[\ce{Q}]&amp;=0.239\\
t&amp;=500 \quad &amp;[\ce{Q}]&amp;=0.263
\end{align}
  
  Using the graph of $\ln [\ce{A}]$ and $\frac 1 {[\ce{A}]}$ vs $t$, determine if the reaction is of order $1$ or $2$.


I know how to solve these kind of problems when I'm given the concentration of $[\ce{A}]$, I just plot the graph and check which one resembles a line more, and the slope gives me the velocity constant $k$.

But here $[\ce{A}]$ is not given. I supposed it would be $[\ce{A}](t)=[\ce{A}]_0-[\ce{Q}](t)$, but I was told that was wrong. Could someone help me out?
If the stoichiochemistry is as you present it:

$$\ce{A -&gt; P + Q}$$

Then if at $t=0$, $[\ce{A}]=[\ce{A}]_0$ and $[\ce{Q}]=0$, then because one equivalent of $\ce{Q}$ is formed for each equivalent of $\ce{A}$ consumed, your method of determining $[\ce{A}]_t$ must be correct. 

$$[\ce{A}]_t = [\ce{A}]_0 - [\ce{Q}]_t$$

Perhaps the error your instructor noted came in applying your relationship or interpreting your graph.