A) We have an experiment , using the solution of KNO3 , measured freezing point of solution is -1.15 degree , and using a sample of pure water the thermometer read o.25 degree as freezing point of the sample , calculate the molal concentration of KNO3 assume that the attraction force between the ions is ignored

B) If we do not ignore the attraction force , and decided to take a sample of that solution 10 mL ,, after boiling and vaporizing the water from the solution , got o.415 g of KNO3 , then find out the real molality of KNO3 , and what is percentage difference between measured concentration and real concentration of KNO3 , if the solution density is 1 g/mL

Thank you a lot

I think you meant to say that the freezing point of the pure water as read on your thermometer was 0.25. Therefore, the freezing point of the original solution KNO2 was -1.40.

Then 1.40 = i*Kf*m
1.40 = 2*1.86*m
m = about 0.376

0.415g KNO3 = 0.415/101.1 = about 0.0041 mols.
m = 0.0041mols/0.010kg solvent = about 0.410m
absolute difference = |0.376-0.410|= 0.034
% diff = (0.034/0.41)*100 = ?

A) To calculate the molal concentration of KNO3 in the solution, we can use the formula:

ΔTf = Kf * molality

where ΔTf is the change in freezing point, Kf is the cryoscopic constant, and molality is the molal concentration.

For pure water (the sample), the freezing point is 0.25 degrees Celsius. The freezing point of the solution is -1.15 degrees Celsius. Therefore, the change in freezing point (ΔTf) is:

ΔTf = -1.15 - 0.25 = -1.40 degrees Celsius

Next, we need to find the cryoscopic constant (Kf) for water. The cryoscopic constant for water is approximately 1.86 degrees Celsius/mole.

Now we can calculate the molal concentration (molality):

-1.40 = 1.86 * molality

molality = -1.40 / 1.86 ≈ -0.75 mol/kg

Therefore, the molal concentration of KNO3 in the solution is approximately -0.75 mol/kg.

B) To calculate the real molality of KNO3, we need to determine the moles of KNO3 present in the solution.

Given that the mass of KNO3 obtained after vaporizing the water is 0.415 g and the density of the solution is 1 g/mL, we can determine the volume of the solution:

Volume of solution = mass of KNO3 / density of solution
= 0.415 g / 1 g/mL
= 0.415 mL

Since the volume of the solution is 10 mL, we can calculate the dilution factor:

Dilution factor = 10 mL / 0.415 mL
= 24.1 approximately

This means that the sample was diluted by a factor of 24.1 during the process.

Now, we can calculate the real molality:

molality = (moles of KNO3) / (mass of solvent in kg)

The moles of KNO3 can be calculated using the molar mass of KNO3 (101.1 g/mol):

moles of KNO3 = mass of KNO3 / molar mass of KNO3
= 0.415 g / 101.1 g/mol
≈ 0.0041 mol

The mass of the solvent can be calculated using the volume and density:

mass of solvent = (volume of solution - volume of KNO3) * density of solution
= (10 mL - 0.415 mL) * 1 g/mL
= 9.585 g

Converting the mass of the solvent to kg:

mass of solvent = 9.585 g / 1000
= 0.009585 kg

Finally, we can calculate the real molality:

molality = 0.0041 mol / 0.009585 kg
≈ 0.428 mol/kg

The percentage difference between the measured concentration and real concentration can be calculated using the formula:

Percentage difference = (|measured concentration - real concentration| / real concentration) * 100

Measured concentration of KNO3: -0.75 mol/kg (based on previous calculation)
Real concentration of KNO3: 0.428 mol/kg

Percentage difference = (|-0.75 - 0.428| / 0.428) * 100
= (|-1.178| / 0.428) * 100
≈ 275.70%

Therefore, the percentage difference between the measured concentration and real concentration of KNO3 is approximately 275.70%.

A) To calculate the molal concentration of KNO3 in the solution, we will use the concept of freezing point depression.

The freezing point depression is given by the formula:

ΔT = Kf * molality

where:
ΔT = change in freezing point
Kf = molal freezing point depression constant
molality = molal concentration of the solute

In this case, we are given the freezing point of the solution (-1.15°C) and the freezing point of pure water (0.25°C). Freeze point depression is the difference between these two values:

ΔT = -1.15°C - 0.25°C = -1.40°C

It is mentioned that we assume the attraction force between the ions is ignored, so we can use the molal freezing point depression constant for water, which is 1.86°C/m.

Therefore, we can rearrange the formula to find the molality:

molality = ΔT / Kf = -1.40°C / 1.86°C/m = -0.75 m

Note: The negative sign signifies a decrease in temperature.

Hence, the molal concentration of KNO3 in the solution is -0.75 m.

B) To calculate the real molality of KNO3 and the percentage difference between the measured and real concentrations, we need to consider the effect of the attraction forces between the ions.

First, let's calculate the number of moles of KNO3 present in the 10 mL solution:

Mass of KNO3 = 0.415 g
Molar mass of KNO3 = 101.11 g/mol

Number of moles of KNO3 = Mass / Molar mass = 0.415 g / 101.11 g/mol = 0.0041 mol

As the solution has a density of 1 g/mL, the mass of the solution is 10 mL * 1 g/mL = 10 g.

Real molality of KNO3 = Number of moles of KNO3 / Mass of solvent in kg
= 0.0041 mol / 0.010 kg = 0.41 m

Now, let's calculate the percentage difference between the measured and real concentrations:

Percentage difference = |Measured concentration - Real concentration| / Real concentration * 100

Measured concentration = -0.75 m (from part A)
Real concentration = 0.41 m (from the current calculation)

Percentage difference = |-0.75 - 0.41| / 0.41 * 100 = 265.85%

Therefore, the percentage difference between the measured concentration and the real concentration of KNO3 is approximately 265.85%.