Posted: November 9th, 2022

CHM202 H1 and H2 Virtual Lab Activity: Solids and Solutions 2/8/21 and 2/9/21*

CHM202 H1 and H2 Virtual Lab Activity: Solids and Solutions 2/8/21 and 2/9/21*

*Please write a full laboratory report from the background information and data provided below and in the video posted on our Moodle Course Website. Guidelines for laboratory reports are posted in Topic 1 on the course website.

INTRODUCTION AND BACKGROUND INFORMATION:

Purpose:

This laboratory is concerned with the physical properties of solids and solutions. Crystalline solids exist in a lattice of repeating structural units. The student will gain greater appreciation of how atoms are spatially arranged in a solid. The solubility of an ionic compound will be evaluated as a function of temperature. Phase changes associated with super-cooled liquids, supersaturated solutions and sublimation will also be presented.

Introduction:

Everywhere we look we see matter as pure elements, compounds and mixtures. Most of the things that we observe as physical materials are either solids or liquids. Many of the solids we see have a complex structure that is polymeric based from either natural (wood is an example) or synthetic (plastics are an example) sources. Simpler crystalline materials are also observed such as the sugar that we put in our coffee or the salt in the food we eat. We also observe liquid matter normally as water or as solutions of things in water (coffee, tea and soda are examples). It is known that these materials are all comprised of atoms, but how are these atoms arranged? This laboratory session should provide some insight into how atoms and molecules are arranged into the materials that we come into contact with daily.

Crystalline solids exist in an extended network of repeating structural units called a crystalline lattice. This extended network provides additional stability to the solid (related to the lattice energy) and the repeating unit is called a unit cell. We will examine the arrangement of some simple unit cells by making models on a scale we can easily observe. In addition we will calculate the free volume in these structures and calculate the density of some pure metals using these geometric measurements.

Solutions are homogeneous mixtures of one or more compounds in a liquid called the solvent. In our examinations the solvent used will be water and we will determine the solubility of potassium nitrate in water at various temperatures and generate a solubility curve. When determining the solubility of a compound in a solvent, we will be looking for the point at which the solution exists at its maximum concentration that is called the saturation point. We will also examine systems where a compound exists as a liquid below its normal melting point and a solution contains more solute than present in a saturated solution. These systems (supersaturated solutions and super-cooled liquids) are said to be thermodynamically unstable and if given enough time will form a solid but require a seed crystal or shock to initiate crystal formation. Finally a demonstration of sublimation will be observed whereby iodine will be converted from a solid to a gas (sublimation) and then back to a solid (deposition).

Reading Assignments:

You are required to complete the following assignment as part of this laboratory session: Please watch the video from the live session for this virtual laboratory activity. Please review Chapter 10 and 11 from our textbook: Chemistry: The Science in Context 5th Edition Thomas Gilbert.

Grading:
Grading will be as described in the general laboratory handout.

Procedure:

1. Models of Crystalline Unit Cells

1.1 From the Styrofoam balls available in the laboratory take a package of the full and partial balls to prepare models. Record the radius of the balls used. It may be helpful to refer to Figure 11.22

1.2 Using the Styrofoam balls and toothpicks provided prepare a unit cell for the simple cubic, body-centered cubic and face-centered cubic (also called cubic closest-packed; ccp) crystal lattice unit cells.

1.3 Measure the dimensions (unit cell length, l) of each unit cell with a ruler and include this data in the table provided. Also calculate the ratio of the cell length (all are cubes) to the atomic radius (l/r ratio). Draw what you have made to the best of your ability including a representation of the unit cell dimensions. Determine the number of atoms present in each unit cell and enter this data into the table provided.

1.3 Place your finished face centered cubic unit cell at the instructor’s desk along with other student’s unit cells to create a portion of a face- centered cubic lattice structure.

1.4 Calculate the following and enter into the table provided:
1.4.1 The dimensions (geometric unit cell length, l) of the unit cell as a function of radius by using geometry (as opposed to direct measurement of the model). Note: For the body and face-centered cubic you will need to apply the Pythagorean theorem (c2 = a2 + b2 for a right triangle).
1.4.2 The volume of the ‘atoms’ used to prepare the models (4/3r3)
1.4.3 The total volume of the unit cell occupied by the atoms (number of atoms x volume of each atom) and the total volume of the cell (volume of the cube, l3).
1.4.4 The percentage of space occupied by atoms (volume of space occupied/volume of unit cell x 100%) and the percentage of free space for each unit cell (100% – Pct. filled).

1.6 Using the information on crystal structures from Figure 1 below and the following atomic radii, calculate the density that you would expect for the metals Cu (r = 128 pm), Fe (r = 124 pm), and Pb (r = 175 pm). Compare the values you obtained with the reported values for density listed in the table. Note: You will need to calculate the mass (using Avogadro’s number) and volume (using the atomic radius and the volume of a sphere) of each metal atom.

2. Determination of the Solubility Curve for Potassium Nitrate.

2.1 Using a heating block set near 90oC (or a water bath) heat the samples described below to up to 90 degrees C. (Note that higher concentrations will require higher temperatures to dissolve.) The instructor may request that you substitute a thermocouple temperature probe for the thermometer.

2.2 Add the following amounts of potassium nitrate to four separate test tubes: 2, 4, 6, and 8 grams. Record the weight to within at least 0.01 grams.

2.3 Pipette 5.00 mL of water into each test tube.

2.4 Heat the test tubes until the samples completely dissolve using a glass stir rod to gently mix the contents. Do not remove the glass stir rod from the mixture until it is completely dissolved.

2.5 Move the test tube containing the KNO3 solution to a test tube rack and quickly place the thermometer into the KNO3 solution. If the sample crystallizes upon the addition of the thermometer then heat the sample and the thermometer together in the heat block until re-dissolved.

2.6 Slowly raise and lower the thermometer in the test tube and record the temperature at which a crystal is first observed. Record this temperature as accurately (within 0.1 degree) as possible. This is the temperature of a saturated KNO3 solution.

2.7 Repeat with the other concentrations of KNO3 recording the temperature for each tube (Note: The temperature required to completely dissolve the sample will be lower at lower concentrations).

2.8 Graph the solubility of KNO3 in g/100 mL as a function of temperature on the graph paper provided.
2.9 Dispose of the KNO3 solutions in the container provided in the laboratory.

3. Super-cooled Liquids and Supersaturated Solutions (Acetic Acid and Sodium Acetate)

3.1 Remove a bottle of pure acetic acid which has been cooled from the instructor’s desk. Record the temperature of the ice/water bath that the acetic acid was held in and compare to the melting point of acetic acid. Handle the bottle carefully and observe if there is any solid present in the solution. Dry off the thermometer and use the thermometer to gently stir the solution and record any observations.

3.2 Take one of the flasks of concentrated sodium acetate solution which is at or below room temperature. Make sure that there are no crystals evident in the solution used (if there are then use another flask of sodium acetate which does not have crystals).

3.3 Observe carefully the beaker containing the sodium acetate solution which has been cooled and confirm that there is no solid present.

1.4 Add a spatula of sodium acetate crystals to the sodium acetate solution and record the results

4. Sublimation (Iodine)

4.1 Observe the demonstration at the instructor’s desk where iodine is being sublimed. What do you observe about the solid in the bottom of the beaker and the solid on the watch glass? What about the “air” in the beaker. Record your observations.

DATA AND RESULTS

Please watch any video provided for each section for analyzing Results and Data below. Also consult with online virtual laboratory session Tuesday 10/13/20 posted on Moodle Course Website with Chapter 10 content.

PART 1 Models for Crystalline Unit Cells
Table 1 Building Cubic Unit Cells (Student Data from Previous Semester)

Atomic Radius: r = 3.78 cm

Volume of one atom: V = _

Simple Cubic (SC)
R: 1.25
A: 2.5
Body Centered Cubic (BCC)
R: 1.38
A: 2.75
Face Centered Cubic (CCP)
R: 1.88
A: 3.75
Measured unit cell length (l) 6.6cm
No. Atoms per cell

Ratio of L/r
Geometric unit cell length (l)

Volume of atoms in cell
Total Volume of Unit Cell

Percent volume filled

Percent free space

Table 2 Theoretical Unit Cell Calculations
Copper
Iron
Lead
Crystal Form CCP
BCC
CCP
Radius (pm)

No. atoms per cell

Dimensions of unit cell

Volume of unit cell

Mass of one atom

Total mass of atoms in one unit cell

Mass of atoms per unit cell volume (calculated density)

Reported Density (g/cm3)

Notes on Calculations for Tables 1 and 2:
• Ratio of l/r: Length divided by radius
• Geometric unit cell length:
• Volume of atoms in cell: (4/3πr3) x Number of atoms
Total volume of unit cell: length3
• Percentage of space occupied by atoms: (Volume of atoms / volume of unit cell x 100%)
• Mass of one atom: Atomic mass / Avogadro’s Number
• Dimensions of unit cell: atomic radius x ratio of l/r of that particular unit cell model
• Total mass of the unit cell: Mass of one atom x number of atoms per cell
• Calculated density: Mass of atoms per unit cell / Total volume of unit cell Percentage of free space for each unit cell: (100% – Pct. Filled)
PART 2 Determination of the Solubility Curve for Potassium Nitrate

Potassium nitrate solubility curve

Table 3 Potassium Nitrate Solubility Curve (Please Graph This Data)
Approx g KNO3
per 5 mL water Mass of KNO3
(grams) Volume
Water
(mL) Saturation
Temperature
(oC) Calc.
g KNO3
per 100 mL
2 2.0192g 5.00 mL 34.9
4 4.0091 5.00 mL 51.7
6 6.0046 5.00 mL 69.8
8 8.0020 5.00 mL 80.5

PLEASE GO ON TO NEXT PAGE TO GRAPH THIS DATA (or use excel)

Graph of Results (Solubility/100 mL water vs. Temperature)

Temperature (oC)

Notes for Determination of the Solubility Curve for Potassium Nitrate:
Table 3 Calculated g KNO3 per 100mL: Mass of KNO3 (grams) x 20

PART 3 Super-cooled Liquids and Supersaturated Solutions (Acetic Acid and Sodium Acetate)
Observations (Data) for Supersaturated Solutions (Acetic Acid and Sodium Acetate)
Acetic Acid
Temperature of water bath: -1°C, freezing point of acetic acid = 16.6°C Observations before shaking the bottle: Clear, slightly cloudy liquid Observations after shaking the bottle: No change initially, but then some ice flakes stared to appear in the solution and crystallize on the side of the bottle. Sodium Acetate
Observations before adding seed crystal: Clear liquid
Observations after adding seed crystal: After several scoops of sodium acetate crystal were added, no changes were observed, the powder dissolved or mix in. Finally, after much cooling, we saw the solute come out of solution, as if it suddenly became a snow globe.
Crystallization of glacial acetic acid and sodium acetate

Sodium acetate:

Acetic acid

PART 4 Sublimation of Iodine (Iodine Crystals)

Sublimation iodine:

Questions:

1. How do the densities of the metals calculated in section 1.6 correlate with the reported values for these metals? Could this method be used for determining the density of mercury as well? Briefly explain your answer. (Hint: what do you know about mercury as an element?)

2. Water is a unique material in that the density of the solid is lower than the density of the liquid (which is why ice forms at the top of a pond and why ice floats in our drinks). If the density for ice at 0 oC is 0.917 g/mL and the density for water at 0 oC is 0.999 g/mL, what is the calculated free space (as %) for each of these materials? You will need to estimate the volume of water as the sum of 2 H atoms and 1 O atom with radii of 37 and 66 pm respectively. Note that you will also have to assume a quantity of water to perform this exercise.

3. From the information in the phase diagram of carbon dioxide (Figure 11.41) would the method used to sublime iodine at atmospheric pressure also work to sublime and collect dry ice (solid carbon dioxide)? Explain your answer briefly. (Hint: At what temperatures would dry ice exist at 1 atmosphere? Remember you used this material last semester in a laboratory exercise).

4. Would the method used to determine the solubility curve for potassium nitrate work for sodium acetate as well? Why or why not?

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