How many moles of #Cl^-# ions are needed to completely combine with 0.25 moles of #Mg^(2+)# ions?

1 Answer
Aug 19, 2016

Answer:

To form #MgCl_2#? Clearly we need #0.50*mol# of chloride ion.

Explanation:

In all these sorts of problems, reference to a balanced chemical equation is necessary as a first step:

#Mg^(2+) + 2Cl^(-) rarr MgCl_2#

This equation is balanced with respect to (i) mass, and (ii) charge, and is therefore a reasonable representation of chemical reality, of physical reality if you like. The equation clearly states that each equiv magnesium ion requires 2 equiv of chloride ion to form the salt. Since we started with a given NUMBER of #Mg^(2+)# ions (#0.25*mol#) it follows that we need TWICE this number of chloride ions for reaction.

Why do I refer to the #"mole"# as a number?