Ag+-ions are weak oxydators that, for example, can oxydise glucose (the aldehyd group in it).
The equation of that oxydation is:
2Ag++ C6H12O6 + H2O
2Ag(s)+ C6H12O7 + 2H+ ΔH > 0 ( eq.1)
To force this equilibrium to the right, you have two options: heating and strengthen the basic environment.
Ag+-ions, brought in basic environment, will form a precipitation: AgOH; that soon will form: Ag2O with a brownish color and being a solid.
But: Ag+ reacts only reasonable well as an oxydator when it is free, dissolved; not when it is connected to OH- or to O2-.
If you want to apply this ion as an oxydator, then free ions must be available, even in basic environment.
There is an option to realise that: adding ammonia (NH3(aq)), that can keep the Ag+-ions dissolved in the following way:
Ag+ + 2NH3
In this way the Ag+-ions are kept in solution, remain dissolved, even in basic environment.
Increase of temperature:
An endothermic reaction costs energy, so: extra energy will support the endothermic reaction.