Does freezing point depression depend on the charge of the solvated ions?
First let's gather the data from the CRC Handbook of Chemistry and Physics, 92nd edition (mol/L is moles per liter). The data below is reported as molarity rather than molality since the problem is posed in molarity. However in a solution as dilute as 0.1 M the difference shouldn't be large. 

         wt%      mol/L      Delta-F
  HCl     0       0        0.00
          0.5     0.137   -0.49
          1       0.275   -0.99
          2       0.553   -2.08
          3       0.833   -3.28
          4       1.117   -4.58
          5       1.403   -5.98
  least squares fit Delta-f = = -3.4683*[HCL] - 0.5624*[HCL]^2  
  [HCL] =0.1 M, Delta-f = -0.35

         wt%      mol/L      Delta-F
  NaBr    0       0.00     0.00
          0.5     0.049   -0.17
          1       0.098   -0.34
          2       0.198   -0.69
          3       0.301   -1.04
          4       0.405   -1.39
          5       0.512   -1.76
  least squares fit Delta-f = = -3.4823*[NaBr] + 0.0930*[NaBr]^2
  [NaBr] =0.1 M, Delta-f = -0.35        

         wt%      mol/L      Delta-F    
  KNO3   0        0        0.00
         0.5      0.05    -0.17
         1        0.099   -0.33
         2        0.2     -0.64
         3        0.302   -0.94
         4        0.405   -1.22
         5        0.509   -1.50
  least squares fit Delta-f = = -3.4264*[KNO3] + 0.9923*[KNO3]^2
  [KNO3] =0.1 M, Delta-f = -0.33            

         wt%      mol/L      Delta-F    
  MgSO4  0        0        0.00
         0.5      0.042    -0.10
         1        0.084    -0.19
         2        0.169    -0.36
         3        0.257    -0.52
         4        0.346    -0.69
         5        0.437    -0.87
  least squares fit Delta-f = = -2.3466*[MgSO4] + 0.9726*[MgSO4]^2
  [MgSO4] =0.1 M, Delta-f = -0.22   


So there is a difference. 0.1 molar $\ce{MgSO4}$ has a freezing point depression of 0.22 degrees Celsius whereas the other 0.1 molar salts are about 0.35 degrees.

As far as the "answer" goes the problem is very poorly worded. First the problem states to "assume an ideal solution," then poses the problem in such a manor that some sort of non-ideal behavior must be assumed to obtain an answer. "Ideal behavior" for electrolytes only occurs in very dilute solutions. A 0.1 molar solution is much too concentrated for ideal electrolyte behavior. 

Without any "lookup" of data the only solution seems to be some hand-waving about $\ce{MgSO4}$ being the only salt with divalent ions so the effective dissociation is smaller because of forming ion-pairs in solution.