A 1.00% by mass MgSO4(aq) solution has a freezing point of -0.192°C.
(a) Estimate the van't Hoff i factor from the data.
(b) Determine the total molality of all solute species.
(c) Calculate the percentage dissociation of MgSO4 in this solution.
Can anyone help me with this at all?
a) delta T = i*Kf*m
i = delta T/(Kf*m).
Substitute and solve for i. I found 1.23 but you need to confirm that.
(You will need to find m and a 1% solution MgSO4 means 1 g MgSO4 + 99 g water; therefore, molality = moles MgSO4 (which is 1 g/molar mass MgSO4)/kg solvent (and kg solvent is 0.099 kg). I find something like 0.0839m but you need to confirm that.
b).......... MgSO4 ==> Mg^+2 + SO4^-2
start ...0.0839....0........0
change.........-x.....x.........x
end........0.0839-x....x.........x
total = 0.0839-x + x + x = 1.23*0.0839
Solve for x which is the amount ionized.
c) Then percent ion = (amount ionized/starting amount)*100 = ?? and I find something like 23 or 24%. (The starting amount is 0.0839)
Having done all of this I want to state that I STRONGLY disagree with the idea of giving this type problem to students. It fosters the incorrect idea that MgSO4 really has a percent ionization that ISN'T 100%. I have seen similar problems posted on Jishka and in problem texts in which NaCl is treated the same way. Nonsense. NaCl is 100% ionized, MgSO4 is 100% ionized. This 23% for MgSO4 in this problem is nonsense. While it is true that the solution behaves as if it were 23% ionized, it ionizes 100% and any deviation from that is because it is not an ideal solution. Read up on the Debye-Huckel theory and activity coefficients if you want to find the real reason why it behaves this way. And while I'm at it, I should point out that there is a MUCH easier way to go about the fake percent ionization. Note that i = 1.23, and note if it were not ionized at all i = 1.00, then (1.23-1.00)*100 = guess what? 23% and without all of that goop above.
Thanks DrBob, that's what i was thinking about the ionization percentage of the MgSO4 as well but I figured out the problem with a little help from my TA today.
Sure! I can help you with that.
(a) To estimate the van't Hoff i factor, we first need to calculate the change in freezing point (∆Tf). The formula to calculate ∆Tf is:
∆Tf = Kf * m * i
where Kf is the cryoscopic constant (for water it is approximately 1.86 °C/m), m is the molality of the solution, and i is the van't Hoff i factor.
We know that the solution has a freezing point depression of -0.192°C. By substituting this value for ∆Tf and 1.86°C/m for Kf, we can solve for i:
-0.192 = 1.86 * m * i
(b) The total molality of all solute species can be calculated using the equation:
molality (m) = moles of solute / mass of solvent (kg)
To find the moles of solute, we need to know the mass of MgSO4 dissolved in the solution. However, this information is not given in the question. Therefore, we cannot calculate the total molality in this case.
(c) The percentage dissociation of MgSO4 in the solution can be calculated using the equation:
% dissociation = (moles of dissociated particles / moles of initial solute) * 100
To calculate the moles of dissociated particles, we need the van't Hoff i factor from part (a), and to calculate the moles of initial solute, we need the mass of MgSO4. Since the mass of MgSO4 is not given in the question, we cannot calculate the percentage dissociation either.
So, unfortunately, we are unable to determine the total molality of all solute species or the percentage dissociation of MgSO4 in this solution without the necessary information.
Sure! I can help you with these questions.
(a) To estimate the van't Hoff i factor, we need to know the freezing point depression constant, Kf, for the solvent. Without this information, it is not possible to calculate the van't Hoff factor.
(b) To determine the total molality of all solute species, we need to know the freezing point depression constant, Kf, for the solvent, as well as the molality of the MgSO4 solution.
The molality (m) of a solution is defined as the moles of solute per kilogram of solvent. In this case, the solute is MgSO4 and the solvent is water. Let's assume the molality of the MgSO4 solution is "x" mol/kg.
We know the freezing point depression (∆Tf) is -0.192°C. The relationship between ∆Tf, Kf, and m is given by the equation: ∆Tf = Kf * m.
So, by substituting the known values into the equation, we have:
-0.192°C = Kf * x
However, without the value of Kf for the solvent, we cannot calculate the total molality.
(c) To calculate the percentage dissociation of MgSO4 in the solution, we need to know the complete dissociation equation of MgSO4 and its concentration in the solution.
The dissociation equation for MgSO4 is:
MgSO4 → Mg²⁺ + SO₄²⁻
To calculate the percentage dissociation, we need to know the initial concentration of MgSO4 and the concentration of Mg²⁺ or SO₄²⁻ in the solution after dissociation.
However, the question only provides the mass percent of MgSO4 in the solution, which is not sufficient to calculate the percentage dissociation.
In summary, to fully answer these questions, we need additional information such as the freezing point depression constant (Kf) of the solvent and the initial and final concentrations of the solute.