Identifying spontaneous redox reactions Which one of the processes is NOT spontaneous?
  
  \begin{align}
\ce{Ni + Sn^2+ &amp;-&gt; Ni^2+ + Sn} \label{1} \tag{1} \\
\ce{Cu + Fe^2+ &amp;-&gt; Cu^2+ + Fe} \label{2} \tag{2} \\ 
\ce{Ni + 2Ag+ &amp;-&gt; Ni^2+ + 2Ag} \label{3} \tag{3} 
\end{align}


I failed this class last semester so I'm redoing the tests to see if I can get this over with, however I'm confused with this question. To find out if a process is spontaneous I need to identify the oxidizing and reduction agents. 

\eqref{1} $\ce{Ni}$ is the reduction agent and $\ce{Sn^2+}$ is the oxidizing agent;
 \eqref{2} $\ce{Cu}$ is the reduction agent and $\ce{Fe^2+}$ is the oxidizing agent;
 \eqref{3} $\ce{Ni}$ is the reduction agent and $\ce{Ag+}$ is the oxidizing agent.

For a redox reaction to be spontaneous the reduction potential of the oxidizing agent needs to be higher than the reduction potentials of the reduction agent, therefore \eqref{1} and \eqref{3} are not spontaneous, but only one of the reactions can be not spontaneous. Where did I go wrong?
For starting purposes, you should write the half cells, either with their oxidation/reduction potentials.
Now if you want to compare the oxidizing species then you have to consider the oxidizing potentials of all the species and the one with highest value will be the spontaneous half cell and the reverse of the other one will occur. 
Remember you should consider only one type of potential at a time(i.e. either oxidation or reduction) for comparison.