# 3.2A: Solutions and Freezing Point Depression


Learning Objectives

Goals:

• Collect experimental data and express concentration in different units.
• Determine the Van't Hoff factor for two different concentrations of sodium nitrate

By the end of this lab, students should be able to:

• Use a volumetric flask to make a solution
• Know the precision of volumetric flasks and perform calculations to the correct number of significant digits.
• Demonstrate the colligative property of freezing point depression

Prior knowledge:

## Safety

• Emergency Preparedness
• Eye protection is mandatory in this lab, and you should not wear shorts or open toed shoes.
• NaNO3 PubChem LCSS
• Minimize Risk
• Check the cord on the hotplate, inform the instructor if it is frayed.
• Make sure the electrical cord never touches the surface of the hotplate
• Recognize Hazards
• You will be working with an alcohol thermometer and care must be taken when handling it.  You will be using a water bath to heat your solution and must always take care not to let the thermometer touch the bottom of the glass where it is in contact with the hot plate.
• All waste is placed in the labeled container in the hood and will be recycled when the lab is over.  Contact your instructor if the waste container is full, or about full.

## Equipment and materials needed

 25 mL volumetric flask analytical balance NaNO3(s) Vernier LabQuest Temperature Probe ice eye dropper parafilm 3 100 ml beakers 3 50 mL beakers

## Background

There are two parts to this experiment.  In the first part, you will make two different aqeuous solutions of sodium nitrate and express their concentration in different ways.  Then you will add ice to the solutions and observe its effect on the on the freezing point, by observing how cold they can get, and compare them to a solution of ice and water.  You will then apply the freezing point depression equation and determine the Van't Hoff factor for each of the two solutions.

### Part 1: Solutions

In this part of the experiment you will make two solutions of potassium nitrate solution and report their concentrations in different units.  You will then measure the effect of the salt concentration on the melting point of each solution. Section 13.1: Units of Concentration goes over the various methods of representing solution concentrations.

### Part 2: Freezing Point Depression

A colligavite property is when a solute affects a solute property like its freezing or boiling point, and in this lab we will investigate the effect of the different solute concentrations created in part 1 on the freezing point of water. You should familiarize yourself with sections 3.4.3-3.4.6 of the text as you prepare for this lab.

You should observe that as the salt concentration increases the freezing decreases, and this can be explained by the following equation:

\begin{align} \Delta T & =-ik_{f}m \\ &\text{where} \nonumber \\ i & = \text{Van't Hoff Factor} \nonumber \\ k_{f} & = \text{freezing point point constant} \nonumber \\ m & = \text{molality} \left ( \frac{mole_{solute}}{kg_{solvent}} \right ) \nonumber \\ & and \nonumber \\ \nonumber \Delta T & = T_{fp}\text{(solution)}-T_{fp}\text{(pure solvent)} \nonumber \end{align}

The value of kf for the solvent water can be obtained from table 13.4.2.

Note

If the value if kf in $$\Delta T & =-ik_{f}m$$ is a negative number, remove the negative sign from the equation.

## Experimental Design Considerations

In part 1 you will make two solutions, in the first you will show that volumes are not always additive, that is, 1 mL plus 1 mL may not equal 2 mL. To understand this you should review section 13.2, the Solvation Process from the perspective of the intermolecular forces

### Expressing Solutions

Consideration $$\PageIndex{1}$$

While making the first solution you were instructed to make exactly  25 mL of a heterogenous mixture, that is, to dilute to volume without dissolving all the solid.  So you had exactly 25 mL of material in the volumetric flask, with solid NaNO3(s) at the bottom of the flask and some NaNO3(aq) above it.  You then stirred it so that all the NaNO3(s) disapeared, and then should have noticed the volume shrank.  Can you explain this in terms of intermolecular forces and the concepts of section 13.2, the solvation process.

Initially, most of the volume was occupied by the solvent (solvent-solvent interactions), predeominantly hydrogen bonding of water, with some being the ionic crystal energy (solute-solute interactions) of the NaNO3(s) precipitate. When you dissolved the crystal's ions you had ion-dipole interactions between the ions and the water (solute-solvent interactions) and these were stronger than the solvent-solvent intermolecular forces.  Since they are stronger, the "bond length" became shorter, and so the volume contracted.

Consideration $$\PageIndex{2}$$

When would volumes be additive.  That is, there is no change in volume upon mixing?

If the solute-solvent interaction energies are the same as the solute-solute and solvent-solvent.

### Freezing Point Depression

We are making the assumption that at the instance the ice melts the mixture of ice and water is at the freezing point. There can be multiple reasons this is wrong, and these next experimental considerations will dive into these.

Consideration $$\PageIndex{3}$$

Define Freezing point

The temperature at a given pressure in an isolated systemfor which the solid and liquid phase can co-exist in equilibrium.

Consideration $$\PageIndex{4}$$

Relate the temperature change to the heat transfer between ice and the solution?

The solution cools ( $$m_{H_2O(l)}$$c$$\Delta$$T) as the ice melts (-n$$\Delta H_{fusion}$$)

Consideration $$\PageIndex{5}$$

Why can't we apply the first law and say the heat lost by the ice equals the heat gained by the water?

We do not have an energetically isolated system and need to take into account heat transfer to the room. As both the room and the water are above the temperature of the ice, it will all melt and eventually rise to room temperature, which is being maintained by the heating and cooling system of the building

Consideration $$\PageIndex{6}$$

How could we improve our results and avoid dealing with heat transfer from the room?

By running the experiment in a colorimeter, like we did with the enthalpy of neutralization experiment in general chemistry 1.

Consideration $$\PageIndex{7}$$

We are making the assumption that the once the temperature stabilizes, it represents the freezing point.  This implies that the rate the ice melts equals the rate heat is transferred to the room and water, and as long as there is ice, it will maintain that temperature.  Yet according to the freezing point depression equation, the freezing point will rise during the experiment, why?

As ice melts it turns to liquid water and the molaity goes down as the mass of liquid water goes up. As $$\Delta T$$ is proportional to molality, there is less of a freezing point depression

Consideration $$\PageIndex{8}$$

Why do we use the temperature when the last piece of ice melts, even if it is not the lowest temperature?

As only then do we know the mass of liquid water, which is the total mass of water and ice originally added, and so only at that point can we calculate the molaity.

Consideration $$\PageIndex{9}$$

What could be the mistake if the temperature never flattens out?

There was not enough ice to cool the system to the freezing point

Consideration $$\PageIndex{10}$$

Why might the temperature hit a bottom and start to rise, and why would we not want to use the lowest temperature?

If the temperature hits a bottom and starts to rise it means we have hit the freezing point, but as ice melts the solution becomes more dilute and the freezing point  depression is reduced.  So even if the lowest temperature did represent the freezing point at the concentration, we do not know that concentration, as we do not know how much ice remains unmelted.

Consideration $$\PageIndex{11}$$

What would you expect to happen to the temperature once the last piece of ice melts?

The temperature will now rise to room temperature

So there are a lot of flaws in this experimental design, but these flaws can give us a bearing on the natue of running an experiment.

## Experimental Procedures

### Part 1: Preparing solutions

#### Solution 1

1. Weigh an empty 25 mL volumetric flask and record the mass in your data sheet to the precision of your instrument
2. Label 3 empty 100 mL beakers, weigh and record mass in your data sheet to the precision of your instrument.
3. Weigh approx 1.5 g sodium nitrate and add to the 25 mL volumetric flask.
4. Weigh the flask plus solid solute and record the mass in your data sheet to the precision of your instrument
5. Carefully add water to the volumetric flask so the bottom of the meniscus lines up with the calibration mark.  Try to do this without dissolving all the salt (ie., you want a heterogenous mixture with undissolved solid salt in the presence of water).  Use an eye dropper for the last few drops
6. Cover the top of the volumetric flask with a cover or Parafilm and swirl and invert several times until all the solid is dissolved.
7. Look at the meniscus and record in your data sheet your observation.
8. Using an eyedropper add additional water until the meniscus aligns with the calibration mark
9. Weigh the volumetric flask with dissolved salt and record the mass in your data sheet to the precision of your instrument
10. Pour the solution from the volumetric flask into the 100 mL beaker labeled #1 and set aside.  You will use it in part 2.
11. Clean the volumetric flask with DI water

#### Solution 2

1. Weigh approx 4.5 g sodium nitrate and add to the clean volumetric flask from step 10 above
2. Weigh the volumetric flask plus solid solute and record the mass in your data sheet to the precision of your instrument
3. Carefully add water until the flask is 2/3rd full (no fluid is going up the neck)
4. Swirl the flask until all the solid is dissolved.
5. Place your hand on the flask and write on the data sheet if the process is endothermic or exothermic
6. Once all the solute is dissolved, dilute to volume, using an eyedropper for the last few drops.
7. Weigh the volumetric flask with dissolved salt and record the mass in your data sheet to the precision of your instrument
8. Pour the solution from the volumetric flask into the 100 mL beaker labeled # 2.  You will use it in part 2.

### Part 2: Freezing point determination

1. Obtain a Vernier LabQuest, hook up the temperature probe and set it to "meter" with the scale reading in oC.  Instructions for running the LabQuest are in the Instrumentation section of your lab manual (sections 0.4.1 & 0.4.2).
2. Using a clean volumetric flask pour 25 mL of pure water into the 100 mL beaker labeled #3
3. Place temperature probe into pure water
4. Add around 20 mL of fresh crushed ice to a 50 mL beaker.
5. Quickly add the ice to the 100 mL beaker, gently swirl the temperature probe and record the temperature just as the last bit of ice melts.
6. Repeat steps 3 & 4 for the two salt solutions, and record the temperatures as the last bit of ice melts. Try to use identical amounts of ice for each iteration.
7. Weigh the mass of each beaker with the solution and melted ice. Record on your data sheet to the precision of the instrument
8. Pour the pure water down the sink and pour the sodium nitrate solutions into a waste jar, and it will be recovered for a future class.

## Data Analysis

In this lab you will work your data up in a Google Spreadsheet, the template of which can be downloaded on the report page (3.3A: Solutions and Freezing Point Depression Report). There are three sheets in the Google Workbook

### Cover page tab

The first tab is always the cover page.  Please fill it out.

### Temperature tab

The freezing point depression constant can be found in the Resources tab of any LibreText page (Resources/Reference tables/reference tables/bulk properties/cryoscopic/Melting Point Depression constants) or section 13.4.3.1.  But you should become familiar with the blue resources tab on the left of all LibreText pages, as the function like the appendices in a normal textbook.

3.2A: Solutions and Freezing Point Depression is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts.