Absorbance Spectrum of a Substance Data Sheet
Copper Cycle Reaction
Copper Cycle Reaction Data Sheet
Density Determination
Density Determination Data Sheet
Determination of a Chemical Formula
Determination of a Chemical Formula Data Sheet
Determination of an unknown concentration of nickel nitrate
Determination of an unknown concentration of nickel nitrate Data Sheet
Determination of the Gas Constant, R
Determination of the Gas Constant, R Data Sheet
Determining the Size of Zinc and Iron Atoms
Determining the Size of Zinc and Iron Atoms Data Sheet
Effect of Temperature on Solubility of Potassium Nitrate
Effect of Temperature on Solubility of Potassium Nitrate Data Sheet
Gas Laws Problem Set
Heat of Solution
Heat of Solution Data Sheet
Identification of a Substance Experiment
Identification of a Substance Experiment Data Sheet
Lecture Review Problem Set
Molecular Geometry
Oxidation Numbers
Oxidation Numbers Fillable Worksheet
Preparation and Properties of Oxygen
Preparation and Properties of Oxygen Data Sheet
Preparation and Properties of Oxygen Data Sheet Word Document
Stoichiometry Problem Set
General Chemistry I Laboratory Manual
Overview
This General Chemistry I laboratory manual consists of 11 laboratory wet experiments and 6 activities. The documents within were evaluated for accessibility using NVDA screenreader software.
Table of Contents
Table of Contents:
1. Experiment 1 - Density Determination
2. Experiment 2 - Identification of a Substance
3. Experiment 3 - Heat of Solution
4. Experiment 4 - Determining the Size of Zinc and Iron Atoms
5. Experiment 5 - Determination of a Chemical Formula
6. Experiment 6 - Absorbance Spectrum of a Substance
7. Experiment 7 - Effect of Temperature on Solubility of Potassium Nitrate
8. Experiment 8 - Preparation and Properties of Oxygen
9. Experiment 9 - Copper Cycle Reaction
10. Experiment 10 - Determination of an Unknown Concentration of Nickel (II) Nitrate
11. Experiment 11 - Determination of the Gas Constant, R
12. Problem Set 1 - Nomenclature
13. Problem Set 2 - Oxidation Numbers
14. Problem Set 3 - Molecular Geometry
15. Problem Set 4 - Stoichiometry
16. Problem Set 5 - Gas Laws Problem Set
17. Problem Set 6 - Lecture Review Problem Set
Experiment 1 - Density Determination
Density Determination
Purpose
The purpose of this experiment is to determine the density of water and an unknown liquid.
Introduction
Chemistry is an empirical science meaning that it is based on evidence collected through experiment and observation. Experimental data is used to perform calculations to yield results. The accuracy of the results depends on a number of factors: the accuracy of the measuring device, the precision of the measurements, and how careful you are in taking the measurements. Accuracy is how close a measurement is to the known value, whereas precision is the reproducibility of a measurement. The best description of accuracy and precision is the bullseye example.
Figure 1: Accuracy versus precision
A set of six data points grouped together but far away from the bullseye would be considered precise but not accurate. A set of six data points randomly spaced near the bullseye would show low precision with some accuracy. A set of data is precise and accurate when all six values are on the bullseye.
To eliminate random errors always use the same measuring device (that is, the same balance for every mass measurement); this will help increase the precision of your data. Systematic errors can occur when a measuring device is not calibrated correctly and will affect accuracy. For example, a balance might weigh everything 5 mg too low. The error will always be in the same direction, thus preventing an accurate data set but maintaining precision.
Density is the ratio of substance’s mass to volume at a given temperature as shown by the equation below. Typical units of density are g/mL or g/cm3.
Density is a physical property because it can be measured without changing the composition of the substance. Density is also an intensive property—the density of a substance is the same no matter how much of it there is. Other intensive properties include color, smell, or boiling point. Mass and volume are extensive properties—properties that are affected by the amount of sample.
The density of a substance may change depending on the temperature. This is because the volume of the substance may change; volume increases with temperature. Density and volume are inversely related, meaning that if volume increases, the density will decrease. Literature values of density are usually reported at 20°C. Typically, we estimate the density of water to be 1.00 g/mL to simplify calculations; however, the true density of water at 20°C is 0.998203 g/mL.
You will calculate the density of water and an unknown liquid. You will identify the liquid based on your density results. You will compare your results to the standard in terms of accuracy and precision.
Procedure
Part 1 – Density of Water
- Determine the mass of a clean, dry 50-mL beaker. Record the mass to 3 decimal places.
- Fill a 100-mL beaker halfway with room temperature, deionized water.
- Using a clean, dry 10-mL volumetric pipet dispense 10.00 milliliters of DI water into the pre-weighed beaker. Instructor will demonstrate proper use of the pipet.
- Insert the pipet tip into the liquid, then suction liquid just above the calibration mark.
- While at eye level, adjust the level of the liquid in the pipet so that the bottom of the meniscus lies on the calibration line.
- Transfer the pipet to beaker and allow liquid to drain. Best practice is to have glass to glass contact when draining, meaning touch the pipet tip to the side of the beaker. Finally, the last drop is calibrated to stay in the pipet tip. Do not attempt to blow out this drop.
- Record the mass of the beaker plus 10.00 mL of water.
- On your data sheet, calculate the mass of the water. Report your answer to the correct number of significant figures.
- Using the density equation, calculate the density of water and record this value on your data sheet to the correct number of significant figures.
- Repeat steps 1 through 6 for Trials 2 and 3.
- Calculate the average density of water from the three trials.
- Record the temperature of water to the nearest 0.1 °C.
- Record the literature density of water at the recorded temperature using Table 2, located at the end of the lab procedure. Compare your average to the literature density. Was the calculated density accurate? Were the calculated density values precise?
Part 2 – Density of Unknown Liquid
- Obtain a vial containing unknown liquid from instructor.
- Determine the mass of a clean, dry 50-mL beaker. Record the mass to 3 decimal places.
- Using a clean, dry 10-mL volumetric pipet dispense 10.00 milliliters of unknown liquid into the pre-weighed beaker.
- Record the mass of the beaker plus the 10.00 mL of unknown. This should be done quickly, as some unknown might evaporate quickly.
- On your data sheet, calculate the mass of the unknown. Report your answer to the correct number of significant figures.
- Using the density equation, calculate the density of unknown liquid and record this value on your data sheet to the correct number of significant figures.
- Dispose of the unknown liquid as directed by your instructor.
- Repeat steps 2 through 7 for Trials 2 and 3.
- Calculate the average density of unknown liquid from the three trials.
- Identify the unknown liquid using Table 1.
Table 1: Density of liquids at 20 °C
| Substance | Density (g/mL) |
| Hexane | 0.659 |
| Ethanol | 0.791 |
| Seawater | 1.025 |
| Acetic Acid | 1.05 |
| Ethylene glycol | 1.109 |
Data Sheet
Name:
Unknown Number:
Density of Water
| Density Data | Trial 1 | Trial 2 | Trial 3 |
| Mass of beaker and 10.00 mL water | |||
| Mass of beaker | |||
| Mass of water | |||
| Volume of water | 10.00 mL | 10.00 mL | 10.00 mL |
| Density of water (density = mass/volume) | |||
| Average density of water |
Temperature of water:
Literature Density of Water at above Temperature:
Density of Unknown Liquid
| Density Data | Trial 1 | Trial 2 | Trial 3 |
| Mass of beaker and 10.00 mL unknown | |||
| Mass of beaker | |||
| Mass of unknown | |||
| Volume of unknown | 10.00 | 10.00 | 10.00 |
| Density of unknown (density = mass/volume) | |||
| Average density of unknown |
Density of Unknown Liquid:
Identity of Unknown:
Table 2: Standard Density of water (g/mL) at different Temperatures (°C)
| 0.0 | 0.1 | 0.2 | 0.3 | 0.4 | 0.5 | 0.6 | 0.7 | 0.8 | 0.9 | |
| 15 | 0.999099 | 0.999084 | 0.999069 | 0.999054 | 0.999038 | 0.999023 | 0.999007 | 0.998991 | 0.998975 | 0.998959 |
| 16 | 0.998943 | 0.998926 | 0.998910 | 0.998893 | 0.998877 | 0.998860 | 0.998843 | 0.998826 | 0.998809 | 0.998792 |
| 17 | 0.998774 | 0.998757 | 0.998739 | 0.998722 | 0.998704 | 0.998686 | 0.998668 | 0.998650 | 0.998632 | 0.998613 |
| 18 | 0.998595 | 0.998576 | 0.998558 | 0.998539 | 0.998520 | 0.998501 | 0.998482 | 0.998463 | 0.998444 | 0.998424 |
| 19 | 0.998405 | 0.998385 | 0.998365 | 0.998345 | 0.998325 | 0.998305 | 0.998285 | 0.998265 | 0.998244 | 0.998224 |
| 20 | 0.998203 | 0.998183 | 0.998162 | 0.998141 | 0.998120 | 0.998099 | 0.998078 | 0.998056 | 0.998035 | 0.998013 |
| 21 | 0.997992 | 0.997970 | 0.997948 | 0.997926 | 0.997904 | 0.997882 | 0.997860 | 0.997837 | 0.997815 | 0.997792 |
| 22 | 0.997770 | 0.997747 | 0.997724 | 0.997701 | 0.997678 | 0.997655 | 0.997632 | 0.997608 | 0.997585 | 0.997561 |
| 23 | 0.997538 | 0.997514 | 0.997490 | 0.997466 | 0.997442 | 0.997418 | 0.997394 | 0.997369 | 0.997345 | 0.997320 |
| 24 | 0.997296 | 0.997271 | 0.997246 | 0.997221 | 0.997196 | 0.997171 | 0.997146 | 0.997120 | 0.997095 | 0.997069 |
| 25 | 0.997044 | 0.997018 | 0.996992 | 0.996967 | 0.996941 | 0.996914 | 0.996888 | 0.996862 | 0.996836 | 0.996809 |
| 26 | 0.996783 | 0.996756 | 0.996729 | 0.996703 | 0.996676 | 0.996649 | 0.996621 | 0.996594 | 0.996567 | 0.996540 |
| 27 | 0.996512 | 0.996485 | 0.996457 | 0.996429 | 0.996401 | 0.996373 | 0.996345 | 0.996317 | 0.996289 | 0.996261 |
| 28 | 0.996232 | 0.996204 | 0.996175 | 0.996147 | 0.996118 | 0.996089 | 0.996060 | 0.996031 | 0.996002 | 0.995973 |
| 29 | 0.995944 | 0.995914 | 0.995885 | 0.995855 | 0.995826 | 0.995796 | 0.995766 | 0.995736 | 0.995706 | 0.995676 |
| 30 | 0.995646 | 0.995616 | 0.995586 | 0.995555 | 0.995525 | 0.995494 | 0.995464 | 0.995433 | 0.995402 | 0.995371 |
Handbook of Chemistry and Physics, 53rd Edition, p. F4
Experiment 2 - Identification of a Substance
Identification of a Substance Lab Prep Sheet
Unknowns: Key
Acetone, Isopropyl Alcohol, Hexane, Ethyl Alcohol
Fill numbered vials to neck with unknown. Make all Acetone unknowns, then ethanol etc. CAP TIGHTLY.
Solubility setup
- Dropper bottles
- Paraffin oil- dropper bottle
- Naphthalene with scoop
- DI Water-Dropper bottle on each table, and 1 for DEMO
- Hexane- Dropper bottle on each table, and 1 for DEMO
- Ethanol- dropper bottle on each table, and 1 for DEMO
- Test tube rack with labeled test tubes for DEMO solutions:
- Naphthalene/ water
- Naphthalene/hexane
- Naphthalene/ethanol
- Oil/water
- Oil/hexane
- Oil/ethanol
- Solubility guide (laminated card)
Waste Hood:
- ORGANIC WASTE CONTAINER- for unknowns (Waste hood)
- Return vial contain (Waste hood)
Density determination
- Pipette fillers
- 12- 10mL volumetric pipettes
- Rulers
Boiling point determination
- 12 Side arm test tubes inverted on test tube racks
- 12 #6 (2 y) cut-hole rubber stoppers in labeled beaker
- 12 thermometers
- 12 rubber hoses
BP DEMO- Instructor’s table
- Ring stand, ring, wire square
- 400 mL Beaker (with 300 mL water)
- Side arm test tube (including Boiling stone and 10mL water) with arm pointing to rod of stand clamped at neck to ring stand
- Rubber tubing (on one notch of side arm test tube) leading directly to wide mouth 250mL Erlenmeyer flask
- Thermometer in cut-hole rubber stopper with rubber stopper resting on top of test tube; bulb of thermometer should be about 2-2.5” from bottom of solution. Stabilize thermometer/rubber stopper with clamp
- Plastic container labeled “RETURNED VIALS”- waste hood (dispose in glass waste)
Cleaning
- Rinse pipettes and side arm test tubes with acetone between lab sections
At conclusion of experiments, clean rinsed (H2O) vials, pipettes, and side arm test tubes with cleaning solution. Rinse with tap H2O, then demineralized H2O. Let drain. Return vials, vial caps Cleaned with Sparkleen 5 min., rinsed with tap and DI H2O) to box. Return side arm test tubes to box in cabinet. Return pipettes to labeled drawer.
Identification of a Substance
Purpose
The purpose of this experiment is to identify an unknown liquid using three physical properties: solubility, density, and boiling point.
Introduction
Solubility
A solution is a homogeneous mixture of two or more components: solute and solvent. The solute is the substance being dissolved and the solvent is the substance doing the dissolving. The solubility of a solute in a solvent can be expressed in quantitative terms that denote the concentration of the solution. For example, a common way to express solubility of a substance is the mass of solute in grams per 100 grams of solvent at a definite temperature. Relative solubilities can be expressed qualitatively by terms such as soluble (s), slightly soluble (ss), and insoluble (i). A general rule regarding solubility is "likes dissolves likes", that is, substances of similar structure are soluble in one another.
Density
Another specific property of a substance is density, that is, how much mass is contained in a unit volume of the substance. In the case of solids, the unit volume is 1 cubic centimeter (cm3), for liquids: 1 milliliter (mL), and for gases: 1 liter. Pure substances can often be identified by measuring their densities precisely, since it is rare that any two substances have identical densities.
Boiling Point
When bubbles of vapor form within a liquid and rise freely to the surface and burst, the liquid is said to boil. A liquid exposed to the air will boil when its equilibrium vapor pressure becomes equal to the pressure of the atmosphere. The normal boiling point of a liquid is that temperature at which the equilibrium vapor pressure becomes exactly equal to the standard atmospheric pressure at 760 torr or 1 atmosphere. A liquid will boil at temperatures higher than normal when under external pressures greater than 1 atmosphere; the boiling point of a liquid is lowered below normal by decreasing the external pressure below 1 atmosphere.
Procedure
Part 1 - Solubility Determination
- In separate labeled test tubes, place equal quantities, approximately 2-centimeter depth of the following solvent: water, hexane, and ethyl alcohol. To each test tube, add a few crystals of naphthalene and mix well by shaking and rolling the test tube between the palms of your hands. Body heat speeds up the dissolution process. Let stand for 30 minutes, shaking occasionally. Note results on the data sheet.
- In separate labeled test tubes, place equal quantities, approximately 2-centimeter depth of the following solvent: water, hexane, and ethanol. To each test tube, add 5 to 7 drops of paraffin oil and mix well by shaking and rolling the test tube between the palms of their hands. Note results on the data sheet.
- In separate labeled test tubes, place equal quantities, approximately 2-centimeter depth of the following solvent: water, hexane, and ethanol. To each test tube, add 20 drops of the unknown liquid and mix well by shaking and rolling the test tube between the palms of their hands. Let stand for 30 minutes, shaking occasionally. Note the results on the data sheet.
Part 2 - Density Determination
- Obtain a vial containing an unknown from the instructor.
- Determine the mass of a clean, dry 50-mL beaker. Record the mass to 3 decimal places.
- Using a clean, dry 10-mL volumetric pipet dispense 10.00 milliliters of unknown liquid into the pre-weighed beaker.
- Insert the pipet tip into the liquid, then suction liquid just above the calibration mark.
- While at eye level, adjust the level of the liquid in the pipet so that the bottom of the meniscus lies on the calibration line.
- Transfer the pipet to beaker and allow liquid to drain. Best practice is to have glass to glass contact when draining, meaning touch the pipet tip to the side of the beaker. Finally, the last drop is calibrated to stay in the pipet tip. Do not attempt to blow out this drop.
- Record the mass of the beaker plus the 10.00 milliliters of the unknown. This should be done quickly, as some unknown might evaporate quickly.
- Pour the unknown liquid from the beaker back into the vial and save for the boiling point determination in Part 3.
- On your data sheet, calculate the mass of the unknown. Report your answer to the correct number of significant figures.
- Using the density equation, calculate the density of your unknown and record this value on your data sheet to the correct number of significant figures.
Part 3 - Boiling Point Determination
- Assemble apparatus as shown in Figure 1. Place approximately 300 milliliters of tap water into the beaker.
Figure 1: Boiling point apparatus
- Attach rubber tubing to side arm test tube.
- Place the side arm test tube into the water bath so the bottom of the test tube is about 2 centimeters from the bottom of the beaker.
- Clamp the test tube at the upper edge and attach it to the ring stand. Adjust the tube so that the side arm points toward the rod of the stand and tubing leads directly to a beaker or flask used as a trough.
- Transfer approximately 10 milliliters of unknown liquid from the vial to side arm test tube, pouring the liquid down the side opposite the side arm.
- To promote even boiling, add 1 boiling stone to the unknown liquid in the side arm test tube.
- Spread the cut portion of a 2-hole rubber stopper and insert a thermometer.
- Clamp the rubber stopper loosely and attach it to the ring stand.
- Adjust the height so the bulb of the thermometer is centered 2 centimeters above the level of the liquid in the test tube and the rubber stopper rests on the test tube.
- Have the set-up checked by your instructor before starting to heat.
- Heat water bath with a moderate flame. The tip of the inner core flame should be about 2 to 3 centimeters below the ring. Heat the unknown until temperature remains constant for 30 seconds to 1 minute. This is the boiling point. Record the temperature on your data sheet to one decimal place.
- Remove flame. Avoid boiling to dryness as superheating may occur.
Part 4 - Identification of Unknown and Disposal of Materials
- Using Table 1 below, compare the data collected for the unknown substance with the properties of known substances. Report the identity of the unknown on the data sheet.
- Dispose of hexane solvent from solubility experiment in the Organic Waste Container. If the unknown liquid is soluble in water, other materials in test tubes may be poured in the sink. Otherwise, dispose of insoluble materials in the Organic Waste Container. Clean test tubes with brush and cleanser. If unknown is insoluble in water, dispose of the liquid from the side arm test tube and the vial in the Organic Waste Container. If unknown is soluble in water, dispose of the liquid from the side arm test tube and the vial in the sink.
- Place boiling stone in the garbage can.
- Return empty unrinsed vials to the container labeled Returned Vials.
Table 1: Physical properties of miscellaneous substances
| Substance | Density (g/mL) | Melting Point (°C) | Boiling Point (°C) | solubility in water | solubility in hexane | solubility in ethanol |
| Acetamide | 1.16 | 81 | 222 | s | i | s |
| Acetone | 0.79 | -95 | 56 | s | s | s |
| Benzene | 0.88 | 5.5 | 80 | i | s | s |
| Carbon tetrachloride | 1.60 | -22.6 | 77 | i | s | s |
| Chloroform | 1.49 | -63.5 | 61 | i | s | s |
| Diethyl Ether | 0.71 | -120 | 35 | ss | s | s |
| Ethyl alcohol | 0.79 | -112 | 78 | s | s | s |
| Hexane | 0.66 | -94 | 69 | i | s | s |
| Isopropyl alcohol | 0.79 | -86 | 83 | s | s | s |
| Methyl alcohol | 0.79 | -98 | 65 | s | i | s |
| Stearic acid | 0.85 | 70 | 291 | i | s | ss |
Data Sheet
Name:
Unknown Number:
Solubility Data
| Substance | Solubility in water | Solubility in hexane | Solubility in ethanol |
| Naphthalene | |||
| Paraffin Oil | |||
| Unknown |
Density Data
| Mass of beaker and unknown | |
| Mass of beaker | |
| Mass of unknown | |
| Volume of unknown | |
| Density of unknown (density = mass/volume) |
Boiling Point of Unknown: ____________________________________
Identity of Unknown: ________________________________________
Experiment 3 - Heat of Solution
Heat of Solution Lab Prep Sheet
DEMO AREA
- Calorimeter with stirrer
- Thermometer
- Cork stopper with thermometer stem extending 10.5 cm through “black cap”
- 50 mL pipette (for demo only)
- Pipette filler
- 250 mL beaker for DI H2O
Hood Area:
- 12- 10mL volumetric pipettes
- Pipette fillers
- Calorimeters with stirrers
- Thermometer
Balance Area:
- 3 bottles of Na2SO4· 10 H2O (Sodium Sulfate Decahydrate)store in refrigerator
Heat of Solution
Purpose
The purpose of this experiment is to determine the heat of solution for sodium sulfate decahydrate. The chemical formula for sodium sulfate decahydrate is ·
Introduction
When a solid dissolves in a liquid, energy is absorbed in overcoming the forces which hold the molecules, atoms, or ions in position in the crystal. Since this requires an input of energy, it is an endothermic change that accompanies the dissolution of all solids in all liquids. The physical process of dissolving is often accompanied by a chemical reaction between the solute and the solvent – usually exothermic in character. If the heat evolved in the chemical change is greater than that absorbed in the physical change, the net process is exothermic; that is, heat of solution, ΔH, is negative. If the net process is endothermic, ΔH is positive. An exothermic heat of solution involves an increase in temperature and an endothermic heat of solution involves a decrease in temperature.
The heat of solution of a solid compound can be easily measured in a calorimetric experiment. The temperature change of the solvent is measured, and the quantity of heat evolved or absorbed is calculated from the known specific heats and the masses of solvent and solute. Heat is measured in units of Joules, specific heat has units Joules per gram per °C, and the temperature change is measured in °C.
Heat energy = mass x specific heat x temperature change
The molar heat of solution can be calculated by multiplying the heat of solution, in units of Joules per gram, by molar mass, in units of grams per mole.
Procedure
- Obtain a calorimeter and a thermometer.
- Insert the thermometer into the calorimeter so that the bulb is not touching the sides or bottom of the calorimeter as shown in Figure 1.
Figure 1: Calorimeter setup
- Use a 10-mL volumetric pipet to transfer 50 milliliters of deionized water into the inner vessel of the calorimeter.
- Measure the temperature of the water.
- Weigh out approximately 5 grams of sodium sulfate decahydrate. Record the exact mass to the nearest 0.001 grams.
- Transfer the sodium sulfate decahydrate into the water. Quickly seal the calorimeter. While stirring continuously, watch the thermometer closely and record the maximum or minimum temperature reached as the solute dissolves.
- Dispose of solution in sink. Rinse the inner vessel of the calorimeter, the thermometer, and the stirrer with deionized water.
Data Sheet
| Measured quantity | Numerical value and units |
| Volume of water | |
| Initial temperature | |
| Density of water at initial temperature | |
| Mass of water | |
| Mass of weighing paper | |
| Mass of weighing paper and solute | |
| Mass of solute | |
| Final temperature |
| Calculations | Numerical value and units |
Heat change for water
| |
Heat change for solute | |
Total heat change | |
| ΔH per gram of solute | |
| Molar mass of solute | 322.1 grams / mole |
| ΔH per mole of solute (J/mole) | |
| ΔH per mole of solute (kJ/mole) |
Experiment 4 - Determining the Size of Zinc and Iron Atoms
Determining the Size of Zinc and Iron Atoms Lab Prep Sheet
Instructor's Table:
- 12 Galvanized zinc squares, about 1.5x1.5” (purchased as sheet of ductwork metal from Lowe’s). Place on Instructor table.
- *Use metal press at Blount Co to cut ductwork.
- Blue rulers
- Beaker labeled "Waste Zn Metal"
Hoods 2-5:
- 6.0M HCl solution w/ 250mL beaker, 25mL graduated cylinder (300 mL per Lab)
- Metal tongs - 2 in each hood with 6M HCl
Waste:
Waste HCl collected in plastic waste bucket. Metal squares can be thrown away.
Solution Prep:
6.0M HCl - hydrochloric acid
- 48.4mL conc. HCl + DI H2O → 100mL
- 242mL conc. HCl + DI H2O → 500mL
- 484mL conc. HCl +DI H2O → 1L
Determining the Size of Zinc and Iron Atoms
Purpose
The purpose of this experiment is to use the mole concept to determine the volume of zinc and iron in a thin square of galvanized metal and to experimentally determine the diameter of individual zinc and iron atoms.
Introduction
Using mole relationships, it is possible to calculate things that cannot possibly be measured. Metals that are susceptible to corrosion are often coated in a thin protective layer. The most common is to add a layer of more reactive metal in a process known as cathodic protection. Steel and iron are often protected from corrosion by a thin layer of zinc, creating galvanized metal. The zinc layer corrodes before the iron, thereby preventing the iron from rusting.
In this experiment, a sample of galvanized iron is placed into hydrochloric acid. The zinc reacts and dissolves; the iron, being a less reactive metal, does not. The reaction is:
Avogadro’s number and molar mass can be used to calculate the number of atoms of each element present. Combined with density, the atomic size of each metal can be determined.
Procedure
- Obtain a sample piece of galvanized iron. Measure the original, exact mass of the metal sample.
- Place the metal sample in a 250-mL beaker and add 25.0 milliliters of 6 M hydrochloric acid. When the reaction stops, pour the acid away as waste.
- Rinse the metal sample thoroughly with tap water, then wash thoroughly with deionized water. Dry the metal sample well.
- Measure the exact mass of the metal sample after reaction.
- Clean up, including washing your hands and returning all equipment.
- Complete calculations to determine the size of zinc atoms, paying careful attention to units.
- Since the zinc was removed by the hydrochloric acid, calculate the mass of zinc by determining the mass lost after the reaction.
- Use the molar mass of zinc to calculate the moles of zinc removed from the metal sample.
- Use Avogadro’s number to calculate the number of zinc atoms removed from the metal sample.
- Use the mass of zinc and the density of zinc, 7.14 grams per cubic centimeter, to calculate the volume of zinc removed from the metal sample.
- Calculate the volume of one zinc atom by dividing the total volume of zinc removed by the number of atoms removed.
- Assuming a simple cubic packing of zinc atoms, calculate the diameter of a zinc atom by taking the cube root of the volume.
- The actual atomic diameter of zinc is 0.268 nm. Calculate the percent error of your calculations compared to the true value.
- Complete calculations to determine the size of iron atoms, paying careful attention to units.
- Since the iron was unaffected by the hydrochloric acid, the mass of iron is the same as the mass left after the reaction.
- Use the molar mass of iron to calculate the moles of iron remaining in the metal sample.
- Use Avogadro’s number to calculate the number of iron atoms remaining in the metal sample.
- Use the mass of iron and the density of iron, 7.874 grams per cubic centimeter, to calculate the volume of iron in the metal sample.
- Calculate the volume of one iron atom by dividing the total volume of iron by the number of atoms in the remaining metal sample.
- Assuming a simple cubic packing of iron atoms, calculate the diameter of an iron atom by taking the cube root of the volume.
- The actual atomic diameter of iron is 0.280 nm. Calculate the percent error of your calculations compared to the true value.
Data Sheet
Name:
Experimental Data
| Original mass of metal sample: | |
| Mass of metal sample after reaction: |
Calculations
Size of Zinc
| Mass of zinc (mass lost): | |
| Moles of zinc removed: | |
| Atoms of zinc removed: | |
| Volume of zinc removed: | |
| Volume of one zinc atom: | |
| Diameter of zinc atom: | |
| Percent error: |
Size of Iron
| Mass of iron (mass remaining): | |
| Moles of iron: | |
| Atoms of iron: | |
| Volume of iron: | |
| Volume of one iron atom: | |
| Diameter of iron atom: | |
| Percent error: |
Experiment 5 - Determination of a Chemical Formula
Determination of a Chemical Formula Lab Prep Sheet
DEMO AREA:
- Magnesium - Mg strips- 25 cm long with forceps (one per group)
- Steel wool wads in 600ml beaker labeled “Return when finished”
- Vacuum grease tubes
INSTRUCTOR’S DEMO:
- Ring stand
- Burner
- Ring 3” above burner
- Clay triangle on ring
- Crucible and lid on clay triangle
- Gauge square
- Small desiccators
- 50 mL beaker
- Crucible tongs
- Stirring rod
- Eye dropper
HOOD:
- 24 Stained crucibles with lids in hood
- Desiccators under hood, -open doors
Determination of a Chemical Formula Experiment
Purpose
The purpose of this experiment is to determine the simplest formula of a compound by synthesizing the compound.
Introduction
A compound is composed of two or more different elements. The Law of Definite Composition states that in a pure compound, the elements are always combined in a definite ratio by mass. If the masses (or %) of each element in the compound are known, the simplest or empirical formula can be determined using the following method.
- Divide the mass (or percentage) of each element by its molar mass. This gives the relative number of moles of atoms of each element in the compound.
- Determine the simple ratio by dividing each value for the relative number of moles by the smallest value so that a ratio of value(s) to one will be obtained.
- Since atoms are indivisible under ordinary chemical conditions, determine the integral, or whole number, ratio by multiplying values in the simple ratio by the smallest integer that will produce a whole number. A number within 0.05 of a whole number may be rounded to a whole number.
- The integers thus obtained are the subscripts corresponding to the number of atoms of each element in the formula.
In this experiment, the formula for the oxide of magnesium is determined by reacting a weighed quantity of magnesium with oxygen in the air forming magnesium oxide. During the reaction, however, magnesium also reacts with nitrogen in the air forming magnesium nitride. To avoid a serious error in formula determination, magnesium nitride is converted to magnesium oxide by the following reactions:
Magnesium nitride + water → magnesium hydroxide + ammonia gas
Magnesium hydroxide → magnesium oxide + water vapor
Thus, a product containing only magnesium and oxygen in chemical combination is obtained. The mass of the oxygen that reacted can be determined by subtracting the mass of the magnesium from the mass of the magnesium oxide. From these data, the formula for magnesium oxide can be determined using the method outlined above.
Procedure
- Assemble the apparatus as shown in Figure 1.
Figure 1: Heating apparatus
- Obtain a crucible and lid. Check to ensure the crucible has no cracks. Support crucible and lid on clay triangle with the lid slightly displaced.
- Heat the crucible strongly for 5 minutes. Cool in air for 2 minutes. Transfer crucible and lid to a desiccator and cool to room temperature.
- Determine the mass of the crucible and lid, and record this on the data sheet.
- Repeat steps 3 and 4 until a constant mass is obtained, that is, the masses must be within 0.005 grams of each other.
- Average the two masses that agree within 0.005 grams and use this value for the constant mass on the data sheet.
- Clean a magnesium ribbon strip with steel wool until the oxide coating is removed, that is, there are no dull spots on the surface.
- Wipe strip with a paper towel.
- Using a paper towel to avoid fingerprints, crumble the magnesium strip into a loose coil. The size of the coil should be such that it fits into the bottom half of the crucible.
- Place the magnesium in the crucible. Obtain the mass of the crucible, lid, and magnesium. Record this mass on the data sheet.
- Determine the mass of the magnesium by subtraction.
- Heat the magnesium in the crucible with the lid slightly displaced using a moderate flame until the magnesium begins to ignite. Adjust the lid to cover the crucible completely and continue to heat with a hot flame for 15 minutes. Control heating by adjusting the flame so that no white smoke escapes from beneath the lid.
- After heating for 15 minutes, adjust the lid so that a slight opening exists to let more air enter the crucible. Continue heating for an additional 15 minutes, adjusting the heat so that no product escapes.
- Remove flame and cool for 2 minutes before removing the lid. Gently tap the lid with the tongs to collect any loose particles into the crucible.
- Place crucible on ring stand base. Using a stirring rod, pulverize the contents. If any rigid form of metal remains, continue heating with a moderate flame and with the lid slightly displaced. Place the stirring rod at the base of the ring stand so that the tip is not touching anything, and any adhered particles may be recovered later.
- Place the stirring rod close to the product in the cooled crucible and drop approximately 10 to 15 drops of deionized water onto the stirring rod to retrieve adhered particles. Move the stirring rod during the water addition so the entire product is just moist. Do not stir.
- Dry the product by heating the covered crucible gently for 2 minutes, and then heat with a stronger flame with lid slightly displaced for an additional 5 minutes.
- Cool the crucible in air for 2 minutes. Transfer to the desiccator and cool to room temperature.
- Obtain the mass of the crucible, lid and product and record this value on the data sheet.
- Pulverize the product with the stirring rod and repeat steps 16 through 19 until a constant mass is obtained, that is, the masses must be within 0.005 grams of each other. No odor of ammonia should be detected if the magnesium nitride has been completely converted to the oxide during the final heating.
- Average the two masses that agree within 0.005 grams and record on the data sheet.
- Subtract the mass of the crucible, lid, and the magnesium from the mass of the crucible, lid, and product to determine the mass of the oxygen.
- Calculate the relative number of moles of each element by dividing the mass of the element by its molar mass. Record these values on the data sheet to the proper number of significant figures.
- Determine the simple ratio of moles of magnesium to moles of oxygen by dividing the relative number of moles of each by the smaller value. Record these values on the data sheet to the proper number of significant figures.
- Determine the integral ratio by multiplying the simple ratio by the smallest numeral that will yield a value withing 0.05 of a whole number. Round to whole numbers and record these values on the data sheet.
- Write the chemical formula for the product using the integers obtained in step 25 as subscripts.
- Place the contents of the crucible in the trash can. Clean crucible and lid with a brush and cleanser.
Data Sheet
Name: ____________________________
| Mass of crucible, lid, and magnesium | |
| Mass of crucible and lid from first heating | |
| Mass of crucible and lid from second heating | |
| Constant mass of crucible and lid | |
| Mass of magnesium used | |
| Mass of crucible, lid, and product from first heating | |
| Mass of crucible, lid, and product from second heating | |
| Constant mass of crucible, lid, and product | |
| Mass of oxygen in the compound | |
| Molar mass of magnesium | |
| Moles of magnesium in the compound | |
| Molar mass of oxygen | |
| Moles of oxygen in the compound | |
| Simple ratio of moles of magnesium to oxygen | |
| Integral ratio of moles of magnesium to oxygen | |
| Formula |
Experiment 6 - Absorbance Spectrum of a Substance
Absorption Spectrum of a Substance Lab Prep Sheet
Solutions: Need about 50-75 mL per semester**
0.1M Co(NO3)2 | (29.104g Co(NO3)2·6 H2O → 1L) Colbalt(II) Nitrate |
0.1M K2CrO4 | (19.420g K2CrO4 → 1L) Potassium Chromate |
0.1M CuSO4 | (24.968g CuSO4·5 H2O → 1L) Copper(II) Sulfate |
Student Benches:
- 1 Spectrophotometers
- 4x Kimwipe lined 50 mL beakers labeled DI H2O, Co(NO3)2, K2CrO4, CuSO4 at each instrument with glass cuvettes
- DI water bottles
Waste Container in waste hood
**NOTE: Only 1st lab section prepares and fills cuvettes. Leave filled cuvettes in beakers for next lab until all labs have done the experiment.
When all lab sections have performed the experiment, cuvettes should be emptied in waste container and cleaned with cleaning solution, rinsed with tap water, rinsed inside and outside with DI water, and left to dry inverted in fresh tissue lined beakers.
NOTE: To avoid scratching, only 4-5 cuvettes should be cleaned at a time
Absorption Spectrum of a Substance
Purpose
The purpose of this experiment is to determine the absorption spectra of solutions using a spectrophotometer and to relate the color of the solution to the corresponding wavelengths in the spectrum where the percent transmittance is high and where absorbance is high.
Introduction
Much of the knowledge of chemical properties of substances has come from the study of how light interacts with matter. When a beam of white light is dispersed into a spectrum, the components of the light are spread according to their wavelength. In the case of white light, there is a continuous gradation of color, that is, it is difficult to decide where the blue part of the spectrum ends, and the green part begins. When these dispersed wavelengths are passed through cells containing samples of either atomic or molecular substances, it is found that the emergent light is no longer continuous. Some of the light waves have interacted with the sample in the cell. The light absorbed is characteristic of the substance just as in the case of emission spectra resulting when an excited atom emits light energy of specific wavelengths. In general, an object transmits (reflects) those wavelengths of light corresponding to its color and absorbs wavelengths that are complementary to its observed color.
An instrument which can separate light into its component colors and measure the intensity of each of these small bands of color is called a spectrophotometer. This instrument is calibrated to read both the percent transmittance and absorbance. The percent transmittance (%T) at a specific wavelength is determined electronically within the instrument by comparing the intensity of the light after passing through the sample cell with the solvent (blank) to the intensity of the light before passing through the cell with the substance of interest (solution).
The percent transmittance is related to the measurements of the instrument (intensity of light measured at the detector), but absorbance is a more useful measurement because of its relationship to the substance. The mathematical relationship is between percent transmittance and absorbance is
.
Absorbance is directly proportional to the concentration of the substance through the Beer-Lambert Law:
Where A is the absorbance, ε is the molar absorptivity, b is the pathlength, and c is the molar concentration. The molar absorptivity is an experimental value that depends on each substance and each wavelength. As long as the wavelength and substance are fixed, then the molar absorptivity is a constant. The path length is the distance light travels through the substance. In most solution spectroscopy, the path length is determined by the size of the cuvette, which is the same during the experiment. Cuvettes are optically sensitive and need to be handled with care to avoid scratches or other defects. Since both molar absorptivity and path length are constant during an experiment, the molar concentration of a substance can be calculated directly from the absorbance. This simple, linear relationship is why spectroscopy is one of the most useful methods in chemistry.
Procedure
- Turn on the spectrophotometer and allow the instrument to warm up for 15 minutes. This gives time for the electronics and light source to reach an equilibrium and have reproducible output/input signals.
- Cuvettes should be cleaned by rinsing with deionized water several times. The outside of the cuvette should be cleaned of fingerprints using tissues (not paper towels). Paper towels could scratch the cuvette.
- Four cuvettes should be filled to the mark using each solution: deionized water (blank), 0.1 M , 0.1 M , and 0.1 M . After filling, be sure to wipe the outside of each cuvette with a tissue (not a paper towel) to wipe away excess solution and fingerprints.
- Measure the %T of light using the following steps starting at 400 nm.
- Set the wavelength to the correct setting.
- Insert the blank solution, deionized water, and set the %T to 100.0%.
- Switch out each solution and record the %T at the selected wavelength for each solution.
- Repeat step #4, increasing the wavelength measured by 25 nm each time. The last iteration will be for a wavelength of 700 nm.
- Review the collected data and look for large gaps in the %T results for each solution. For any gap of 20% between two wavelengths, repeat steps #4 and #5 at 5 nm increments. For example, if %T = 84.5% at 500 nm and 61.2% at 525 nm of a solution, then %T data needs to be obtained at 505, 510, 515, and 520 nm.
- Once all the data has been obtained and there are not any large gaps in %T, use the formula to convert %T to absorbance. Significant Figures: since absorbance is a logarithmic function, digits to the left of the decimal are never significant, but digits to the right are always significant. For example, if %T = 94.2%, then A = 0.026; or if %T = 8.37%, then A = 1.077. Absorbance is typically reported with three significant figures, that is, three decimal places.
Data Sheet
Name:
Experimental Data
Complete the table using the data for the 400-700 nm range for each solution.
| Wavelength (nm) | %T (%) | Abs. (a.u.) | %T (%) | Abs. (a.u.) | %T (%) | Abs. (a.u.) |
| 400 | ||||||
| 425 | ||||||
| 450 | ||||||
| 475 | ||||||
| 500 | ||||||
| 525 | ||||||
| 550 | ||||||
| 575 | ||||||
| 600 | ||||||
| 625 | ||||||
| 650 | ||||||
| 675 | ||||||
| 700 |
Experimental Data to Smooth Gaps
Use this table as needed to fill in gaps in %T that were identified in the original data.
| Wavelength (nm) | %T (%) | Abs. (a.u.) | %T (%) | Abs. (a.u.) | %T (%) | Abs. (a.u.) |
Lab Report
To practice writing lab reports, data analysis, and graphical presentations, complete a partial lab report with the following components.
- Title page with the experiment title, your name, date of experiment, and lab section.
- One page that has a well-organized data table that clearly and neatly lists all the data and results for each solution. Table headings, units, substance ID, percent transmittance, and absorbance should all be present.
- One page that has a graph that plots percent transmittance versus wavelength. The y-axis should be set to range from 0% to 100%, and the x-axis should be set with a domain of 400 nm to 700 nm. The scale for each axis should be set appropriately. The %T data for each solution should be displayed on this one graph, which means that a legend will be needed to clearly distinguish between the data for each solution. The graph for each solution should be set to use a smooth curve with data markers visible. Include all the necessary components of a graph: title, and axes labels with units.
- One page that has a graph that plots absorbance versus wavelength. The y-axis should range from 0 (a.u.) to 2 (a.u.), and the x-axis should be set to have a domain of 400 nm to 700 nm. The absorbance results for each solution should be displayed on this one graph, which means that a legend will be needed to clearly distinguish between the data for each solution. The graph for each solution should be set to use a smooth curve with data markers visible. Include all the necessary components of a graph: title, and axes labels with units.
- One page that includes a purpose and conclusion. Restate the purpose of this experiment in your own words. The conclusion needs to address the following prompts.
- Account for the color of each solution by relating the maximum %T to the color observed. You may need to consult the electromagnetic spectrum.
- Account for the color of each solution by relating the maximum absorbance to the color observed. You may need to consult a color wheel for complementary colors.
- Describe which of the explanations above makes the most sense.
Experiment 7 - Effect of Temperature on Solubility of Potassium Nitrate
Effect of Temperature on Solubility Lab Prep Sheet
DEMO SETUP:
- Ring stand w/ iron ring, wire square, 400 mL beaker
- 2 clamps (same type and extension)
- Wide mouth test tube- clamped to ring vertically
- Cut-hole rubber stopper (#6 2-hole with thermometer clamped to ring stand so stopper rests on mouth of test tube with thermometer about 1 cm above bottom of test tube.
- Cu stirrer
- 5 mL pipette on paper towel
- Pipette filler
- 250 mL beaker for stock DI H2O
1st HOOD:
- Pipette holder (container)
- 24- 5 mL volumetric pipettes
- Pipette fillers and rollers-12
- Thermometers -12
- Rulers-12
MIDDLE HOOD:
- Beaker of #6 (two hole) cut-hole rubber stoppers- 12
- Beaker of Cu stirrers- 12
BALANCE AREA:
- KNO3 (3 Bottles) on separate paper towels with scoop on paper towel for each bottle
Effect of Temperature on Solubility of Potassium Nitrate
Purpose
The purpose of this experiment is to study the effect of temperature on the solubility of potassium nitrate, , in water.
Discussion
The maximum amount of a solute that can be dissolved in a given quantity of solvent at a specified temperature is called that solute’s solubility. Solubility is commonly expressed in units of grams solute per 100 grams solvent. Alternatively phrased, solubility is the concentration of a saturated solution at a specified temperature.
In general, solids are more soluble at higher temperatures (that is, as temperature increases, solubility increases). In this experiment, various concentrations of potassium nitrate in water are prepared and the temperatures at which crystallization starts are determined. These crystallization temperatures correspond to the solubility of potassium nitrate in water at the measured temperature.
Procedure
- Assemble the apparatus shown in Figure 1.
Figure 1. Heating Apparatus.
- In a weighed beaker, obtain 10.000 grams (± 0.001 gram).
- Crush the crystals of with a stirring rod to increase surface area and reduce the time required for dissolution.
- Using a paper funnel that extends to within two inches from the bottom of the test tube, transfer all the weighed into the test tube.
- Place about 50 milliliters of deionized water into a 250-mL beaker. Portions of this water will be pipetted into the test tube at the appropriate times.
- Pipet two 5.00 mL portions of deionized water into the test tube. Avoid splattering. Thus, 10.000 grams per 10.00 grams is the solubility of this first solution.
- Place the thermometer into the liquid about 1 centimeter above the bottom of the test tube.
- Warm the test tube gently and stir gently until all the solid dissolves. At this concentration, all should be in solution by 80°C. Do not allow the temperature to get much above 60°C.
- Allow the solution to cool while stirring gently.
- Watch the test tube closely, and when crystals first form (looks like “snowing” in the test tube), record the temperature in the Solubility Data Table.
- Add exactly 5.00 milliliters more deionized water. This makes a total of 15.00 milliliters of water present – thus 10.000 grams per 15.00 grams is the new solubility.
- When all the solid has re-dissolved (heating if necessary), allow the solution to cool again, while stirring gently. Determine the temperature at which crystallization begins.
- Add another 5.00 milliliter portion of deionized water. This makes a total of 20.00 milliliter water present – thus 10.000 grams per 20.00 grams is the new solubility.
- Determine the crystallization temperature.
- Complete the Solubility Data table.
- Draw a graph titled “Effect of Temperature on Solubility of ”. Plot Solubility (on the vertical axis) and Temperature (on the horizontal axis). Be sure to choose a convenient scale so that the graph nearly fills the page. Label both axes with appropriate units. Plot the four data points and draw a smooth curve.
- Without altering the slope of the curve, extend the curve slightly below 30°C.
- Read from the graph the solubility of at 30.0°C and 50.0°C to one decimal place. Record these predictions on the Solubility Predictions Table.
Data Sheet
Name:
Solubility Data Table
| Temperature (°C) | Mass of (grams) | Mass of (grams) | Solubility of (grams / 100 grams ) |
| 89.5 | 10.000 | 5.00 | 200. |
| 10.000 | 10.00 | ||
| 10.000 | 15.00 | ||
| 10.000 | 20.00 |
Solubility Predictions Table
Temperature (°C) | Predicted solubility of from graph (grams / 100 grams ) |
| 30.0 | |
| 50.0 |
Experiment 8 - Preparation and Properties of Oxygen
Preparation and Properties of Oxygen Lab Prep Sheet
DEMO SETUP: See lab setup pictures
- Pneumatic trough with shelf toward front
- 3 gas collecting French square bottles
- 3 glass plates
- Wide mouth test tube attached to ring stand
- Angle bend in #4 rubber stopper attached to wide mouth test tube and tube of bend attached to rubber hose attached at other end to trough.
- Red rubber tube for overflow spout of pneumatic trough
- Pinch clamp on rubber hose to trough
- Vial of neutral litmus
- Wood splints
- Wash bottle with DI H₂O
- Crucible and clay triangle on ring stand
- Mg 2” strips in 600mL beaker (forceps on a paper towel)
Last Hood:
- Bunsen burner and striker
- Deflagration spoons
- Red phosphorus- scoop on paper towel
- PbO2 with scoop
- Fe2O3 with scoop
- SiO2 with scoop
- CaO with scoop
- Box with expendable test tube for PbO2 test
- Test tube rack with test tubes labeled PbO2, Fe2O3, SiO2, CaO
- Test tube clamp
Middle hoods:
- 12 pneumatic troughs
- 12 red rubber hoses
- 12 white rubber hoses (or yellow/as available)
- 12 angle bends in #4 rubber stoppers
- 12 pinch clamps
- 12 crucibles (no lids)
- French square bottles
- Glass plates
Balance area:
- KClO3- Scoop on paper towel- Potassium Chlorate
- MnO2 w/ scoop-Manganese Dioxide
Preparation and Properties of Oxygen
Purpose
The purpose of this experiment is to prepare oxygen gas, to prepare oxides, and to become familiar with the properties of oxygen and oxygen-containing compounds.
Introduction
Preparation
Oxygen is often prepared in the laboratory through thermal decomposition of potassium chlorate, KClO3.
Manganese (IV) oxide (MnO2) acts as a catalyst for the thermal decomposition, allowing the reaction to occur more rapidly and at a lower temperature.
Basic Anhydrides
Oxygen is a very reactive element. It forms compounds known as oxides by reacting either directly or indirectly with nearly all the other elements. The oxides formed by reacting metals (M) with oxygen (O2) are called metallic oxides, basic oxides, or basic anhydrides (MO).
General: M + O2→ MO
Example: 2Ca + O2→2CaO
In general, oxides of metals, that is basic anhydrides, react with water to form bases. Bases are substances that produce hydroxide ions (OH-) in water, which is why they have a general formula of metal plus hydroxide (MOH). One key idea is that when a basic anhydride reacts with water, the oxidation state of the metal does not change.
General: MO + H2O→ MOH
Example: CaO + H2O→ CaOH2
Acidic Anhydrides
Oxides formed by reacting nonmetals (Nm) with oxygen (O2) are known as nonmetallic oxides, acidic oxides, or acidic anhydrides (NmO).
General: Nm+O2→ NmO
Examples: 2 S + 3O2→ 2SO3 or S + O2→ SO2
In general, acidic anhydrides react with water to form acids. Acids are compounds that release an acidic proton (H+) in water, which is why they have a general formula of HNmO.
General: NmO + H2O→ HNmO
Examples: SO3 + H2O→ H2SO4 or SO2 + H2O→ H2SO3
In the case of elements that react with oxygen to form multiple oxidation states, as the example with sulfur above, the oxides will produce different acids when reacted with water: sulfuric acid versus sulfurous acid in the example above. As with the metal in a basic anhydride, the oxidation state of the nonmetal in the acidic anhydride does not change when it is reacted with water.
Procedure
Part 1 – Preparation of Oxygen
- Weigh out approximately 0.3 grams of manganese (IV) oxide into a crucible.
- Strongly heat the uncovered crucible for 2-3 minutes with a Bunsen burner (Figure 1), until the bottom of the crucible turns orange. This ensures all the combustible impurities in the manganese (IV) oxide are removed.
Figure 1: Heating apparatus
3. Remove the flame and allow the crucible to cool.
4. While the crucible is still cooling, weigh out approximately 4.0 grams of potassium chlorate into a beaker. Break up any large crystals using a stirring rod.
5. Add the cooled manganese (IV) oxide to the potassium chlorate and mix with a stirring rod.
6. Transfer the mixture to a wide mouth Pyrex test tube using a paper funnel. Wipe any loose powder from the mouth of the test tube.
7. Setup the oxygen collection apparatus as shown in Figure 2.
a. Attach the test tube with potassium chlorate mixture to a ring stand.
b. Tightly seal the test tube with a rubber stopper that has a glass tube and a clamped rubber hose.
c. Attach the end of the rubber hose to the inlet of the pneumatic trough.
d. Attach a rubber hose from the overflow drain of the pneumatic trough into the sink.
e. Add water to the pneumatic trough so that it is about half full.
f. Fill three bottles with water and invert them in the pneumatic trough. There should not be any air bubbles in the bottles.
g. Have the instructor check your setup before proceeding.
Figure 2: Oxygen collection apparatus
8. Be certain that none of the potassium chlorate or manganese (IV) oxide comes into contact with the rubber stopper during heating or a severe explosion will occur.
9. With the instructor’s approval, begin heating the test tube by slowly passing the flame back and forth along the bottom of the test tube. Once the mixture begins heating, do not remove the flame or water will be forced into the test tube.
10. Control the strength of heating so that oxygen gas is evolved at a rate convenient for collection.
11. When the first bottle of oxygen gas has been collected, quickly slide the bottle aside and move the next bottle to over the inlet to collect more oxygen. Repeat the quick transition to collect a third bottle of oxygen.
12. While inverted under water, each gas filled bottle should be covered by a watch glass and removed from the pneumatic trough and turned right-side up.
13. Keep each bottle of oxygen covered unless actively performing an experiment.
Part 2 – Preparation of Oxides
Oxides of several elements are prepared by burning them in oxygen.
Preparation of carbon anhydride
- Use the bottle with the least amount of oxygen for this reaction.
- Add approximately ¼ inch of deionized water to the bottle.
- Light a wood splint on fire and then blow out the flame so that only a glowing ember remains.
- Carefully insert the glowing ember and splint into the bottle of oxygen.
- Record observations.
- Allow the splint to remain in the bottle until it stops burning.
- Drop the splint into the bottle.
- Cover the bottle and gently shake to dissolve the oxide in the water.
- Add a piece of litmus paper to the solution and record the observation.
Preparation of phosphorus anhydride
- Perform this experiment in a fume hood.
- Place a small quantity of red phosphorus () in a deflagrating spoon.
- Ignite the phosphorus by heating it in a flame.
- Carefully lower the spoon into the second bottle of oxygen. Avoid letting the spoon touch the sides or bottom of the bottle or the glass may break from thermal stress.
- Record observations.
- Allow the spoon to remain in the bottle until it stops burning.
- Rinse the spoon and sides of the bottle with deionized water.
- Remove the spoon.
- Cover the bottle and gently shake to dissolve the oxide in the water.
- Add a piece of litmus paper to the solution and record the observation.
Preparation of magnesium anhydride
- Add approximately ¼ inch of deionized water to the third bottle.
- Hold a 2-inch ribbon of magnesium with a pair of tongs.
- Ignite the magnesium ribbon using a Bunsen burner.
- Quickly place the burning strip into the bottle. Do not look directly at the burning magnesium.
- Drop the remnants of the magnesium ribbon into the bottle.
- Cover the bottle and gently shake to dissolve the oxide in the water.
- Add a piece of litmus paper to the solution and record the observation.
Preparation and Properties of Oxygen - Data Sheet
Name:
Preparation of Anhydrides
| Element | Observations | Oxide Formula | Classification (metal/ nonmetal) | Balanced Chemical Equation to produce anhydride |
Anhydride Reactions with Water
| Anhydride | Observation from Litmus | Classification (Acid/ Base) | Balanced Chemical Equation of anhydride with water |
Relationship of Anhydrides to Acid or Base
The formula of an acidic or basic anhydride is related to its acid or base by the reaction with water. In general, the oxidation state of the metal or nonmetal is the same in both the anhydride (oxide) and acid or base compound, as shown in the examples below.
Acid compound:
Oxidation state of boron: +3
Acidic Anhydride:
Chemical Reaction:
Base compound:
Oxidation state of vanadium: +4
Basic Anhydride:
Chemical Reaction:
Apply the information learned to deduce the chemical formula of the anhydrides for the following acids or bases.
| Compound | Anhydride | Compound | Anhydride |
Experiment 9 - Copper Cycle Reaction
Copper Reaction Cycle Lab Prep Sheet
SOLUTION PREP:
- 3.0 M NaOH (120 g NaOH in 1 L) 720 mL per lab
- 6.0 M H2SO4 (333.3 mL H2SO4 in 1 L) 360 mL per lab
- Conc HNO3 100 mL
- Conc HCl 360 mL
- Acetone About 60 mL
- Methanol About 60 mL
DEMO SETUP:
Ring stand with ring and wire square.
400 mL beaker on wire square.
On a paper towel place a 50 mL beaker and eyedropper.
DEMO AREA:
- Cu foil (~0.5g), forceps on paper towel
- Steel wool wads in 600mL labeled “Return when finished”
- Boiling stones, forceps on paper towel
- 3.0M NaOH - 250mL labeled beaker and labeled graduated cylinder marked at 30mL
- 12 evaporating dishes
Hood 1:
- Conc HNO3 and labeled 10 mL graduated cylinder marked at 4 mL and 250 mL beaker
- 6.0M H2SO4 and labeled 25 mL graduated cylinder marked at 15 mL and 250 mL beaker
- Conc. HCl with graduated cylinder marked at 15mL and 250mL beaker
Hood 2:
- 1 container labeled WASTE Zn-H2SO4
- 1 100 mL beaker labeled EXTRA Zn
Hood 3:
- Conc. HCl with graduated cylinder marked at 15mL and 250mL beaker
- 1 container labeled WASTE HCl
- 1 10mL graduated cylinder marked H₂O
- 1 hot plate
Organic waste hood:
- Methanol and 10mL graduated cylinder marked at 5 mL and 250 mL labeled beaker
- Acetone with same
- Organic Waste Container
Balance Area:
- granular Zn with scoop on paper towel/one at each balance
Hot Plates:
- 1 on each bench
CLEANUP:
Clean beakers and cylinders with Sparkleen. Rinse thoroughly and return to cabinet after draining.
Copper Reaction Cycle
Measuring Technique Efficiency
Purpose
The purpose of this experiment is to study reactions of copper and copper compounds and to compare the mass of copper recovered with the starting mass of copper, that is, the percentage recovery.
Introduction
A weighed amount of copper metal is reacted with nitric acid to form copper (II) nitrate.
Redox reaction:
Copper (II) nitrate is then reacted with sodium hydroxide to form copper (II) hydroxide, a blue precipitate.
Double Displacement reaction:
Copper (II) hydroxide is decomposed by heat, forming copper (II) oxide, a black precipitate.
Thermal Decomposition reaction:
Copper (II) oxide is then converted to copper (II) sulfate by the addition of sulfuric acid.
Neutralization (or Double Displacement reaction):
Zinc metal is then added to the solution containing copper (II) sulfate. Zinc tends to lose electrons more readily than copper, consequently a displacement reaction takes place in which zinc displaces copper (II) ions from a solution of copper (II) sulfate forming elemental copper and zinc sulfate.
Metal Displacement reaction:
The precipitate copper is separated, dried, and weighed. The percentage recovery is then calculated as indicated:
Procedure
- Half-fill a clean 400-mL beaker with deionized water. Heat it to boiling using a flame. Transfer to a hot plate on medium-low setting until ready to use in Step 8.
- Obtain a copper square (approximately 0.5 grams) and clean it thoroughly with steel wool. Wipe with a paper towel.
- Weigh a clean, dry 250-mL beaker.
- Place the copper square in the 250-mL beaker and weigh again.
- In the fume hood, add 4 milliliters concentrated nitric acid to the copper square. Brown fumes of nitrogen dioxide will be produced. When the reaction is complete (blue solution and no more brown fumes), remove the mixture from the hood and add deionized water until the beaker is about half full.
- Add 30 milliliters of 3.0 M sodium hydroxide to precipitate out copper (II) hydroxide
- Add one large boiling stone. Then with gentle stirring (leave stirring rod in beaker), heat just to boiling using a flame to convert the copper (II) hydroxide to insoluble black copper (II) oxide. The precipitate will settle rapidly if very hot. To assure rapid settling, decrease flame at first indication of “bumping” and continue heating with low flame until solution is very hot.
- Remove beaker to base of ring stand. Remove stirring rod and place on base of ring stand in such a way that adhered particles will not be contaminated and may be recovered later. Allow copper (II) oxide to settle undisturbed. Decant supernatant liquid. To wash the precipitate, add about 200 milliliters of very hot deionized water (from Step 1). Stir then remove the stirring rod. Allow to settle and decant once more. Use medicine dropper to remove as much liquid as possible without removing product.
- Prepare a steam bath for use in Step 18, by half-filling a 400-mL beaker with tap water and place on hot plate on medium-low setting.
- Place the tip of the stirring rod in the beaker containing the copper solution. In the fume hood, pour 15 milliliters of 6.0 M sulfuric acid down the stirring rod. Stir, and roll the solution around the sides of the beaker in order to ensure all black copper (II) oxide dissolves. Leave the stirring rod in the beaker until Step 14. With forceps, hold the boiling stone above the solution and rinse it and the forceps with deionized water. Dispose of the boiling stone in a trash can.
- In a 50-mL beaker, obtain 2.5 grams of granular zinc.
- Place the beaker containing the copper solution in the hood, and slowly add about 2/3 of the granular zinc. This will precipitate copper metal. Stir continuously and continue to slowly add zinc until the solution becomes colorless. At this point, decant the liquid into the zinc-sulfuric acid waste container.
- Still in the hood, add 5 milliliters deionized water to the solid copper, then add 15 milliliters concentrated hydrochloric acid. This will react any excess zinc and will be observed as generation of hydrogen gas bubbles. When bubbling has ceased, ask the instructor for permission to warm the solution on a hot plate (still in the hood). Do not boil. When fumes of HCl are observed, remove the solution from the hot plate and let it cool in the hood. When fumes cease, bring the mixture out of the hood.
- Remove the stirring rod and decant away the HCl solution to another beaker, then dispose of as HCl Waste. Wash the copper precipitate with deionized water. Allow to settle, then decant as water waste.
- Weigh a clean, dry evaporating dish.
- Use a stream of deionized water from a wash bottle to transfer the precipitate from the beaker to the weighed evaporating dish. Allow to settle, then decant.
- In a fume hood, wash the precipitate with approximately 5 milliliters of methanol. Allow to settle and decant the methanol into the Organic Waste container. Next, wash the precipitate with about 5 milliliters of acetone, allow to settle, and decant the acetone into the Organic Waste container.
- Transfer the hot tap water (on hot plate; from step 9) from the 400-mL beaker to a 250-mL beaker. Place the evaporating dish over the beaker of hot water to make a steam bath. Allow at least 30 minutes for the precipitate to dry.
- Remove the evaporating dish from the steam bath and allow it to cool to room temperature. Wipe it to remove moisture and fingerprints. Weigh it and record the mass. Have the instructor observe this weighing, verify the mass, and initial this weighing on the report sheet.
- Calculate the percent of copper recovered.
- Dispose of granular copper in the trash can.
Data Sheet
Record your observations for:
Reaction of copper and nitric acid:
Reaction of copper (II) nitrate and sodium hydroxide:
Reaction of copper (II) hydroxide and heat:
Reaction of copper (II) oxide and sulfuric acid:
Reaction of copper (II) sulfate and zinc:
| Mass of 250-mL beaker with copper strip | grams |
| Mass of 250-mL beaker (empty) | grams |
| Mass of copper strip | grams |
| Mass of evaporating dish and product | grams |
| Mass of evaporating dish | grams |
| Mass of product | grams |
| Percent recovery of copper | % |
Show calculations below.
Experiment 10 - Determination of an Unknown Concentration of Nickel (II) Nitrate
Determination of an Unknown Concentration Lab Prep Sheet
Demo Area:
- Watch glasses
- Burette cards
Hood Area:
- 0.400 M Nickel (II) nitrate hexahydrate Ni(NO₃)₂∙6H₂O
- 250mL labeled beaker
- Rulers
- 3 Unknowns labeled A, B, C with labeled beakers and cuvettes
Student Benches: 6 stations
- 1 Spectrophotometer
- 6 Kim-wipe lined 50mL beakers with glass cuvettes- labeled 1-6
- 6x 50mL beakers - labeled 1-6
- 1 burette stand
- 2 burets
Solution Prep:
- 0.400M Nickel (II) nitrate – 116.3 g Ni(NO₃)₂∙6H₂O per 1 L DI H₂O
Each lab will need ~300mL of solution.
Unknowns:
A - 4.163g /100mL H₂O | 20.815g per 500 mL H2O 0.143 M |
B - 3.098g/100mL H₂O | 15.490g per 500 mLH2O 0.107 M |
C - 6.050g/100mL H₂O | 30.250g per 500 mL H2 0.208 M |
Each group will need 10mL of one unknown. These can be placed
in storage bottles labeled A, B, and C. Place labeled 250mL beaker, labeled
50mL beaker and cuvette with each unknown.
Waste bottle in hood
Determination of an Unknown Concentration of Nickel (II) Nitrate: An Introduction to Beer-Lambert Law
Purpose
The purpose of this experiment is to determine the concentration of an unknown nickel (II) nitrate solution by using a spectrophotometer and constructing a calibration curve from a series of nickel (II) nitrate solutions of known concentration.
Introduction
The nickel (II) nitrate solution used in this experiment is deep green in color and shows the greatest absorbance at a wavelength of 635 nm, referred to as . By setting a spectrophotometer to a wavelength of 635 nm, a red light from the light source will pass through the cuvette containing the solution. A fraction of the light is absorbed by the solution, the rest is transmitted through to the detector. Darker solutions, that is, higher concentration solutions, will absorb more light and transmit less light.
The Beer-Lambert Law states that the absorbance of a solution is proportional to the concentration of the solution.
Where A is the absorbance, ε is the molar absorptivity, b is the pathlength, and c is the molar concentration. This quantitative relationship between absorbance, A, and the solutions concentration, c, forms a linear relationship, as seen in Figure 1.
Five “standard” solutions of known concentration are prepared. Each is transferred to a cuvette and placed in a spectrophotometer for absorbance readings. A graph of absorbance as a function of the concentration is plotted and a direct relationship should result, see Figure 1. This direct relationship is known as the Beer-Lambert Law.
An unknown concentration is then determined by measuring and plotting its absorbance on the vertical axis and locating the corresponding concentration on the horizontal axis.
Procedure
- Turn on the spectrophotometer and allow the instrument to warm up for 15 minutes. This gives time for the electronics and light source to reach an equilibrium and have reproducible output/input signals.
- Obtain approximately 30 milliliters of 0.400 M nickel (II) nitrate in a 50-mL beaker.
- Obtain approximately 30 milliliters of deionized water in another 50-mL beaker.
- Label 4 clean, dry 50-mL beakers #1 through 4. Rinse a 50-mL burette twice with about 4 milliliters of the stock 0.400 M nickel (II) nitrate solution. Fill the burette with the remaining nickel (II) nitrate stock solution. Drain enough of the solution to remove any air bubbles from the tip of the burette. Use this burette to deliver 2, 4, 6, and 8 milliliters of nickel (II) nitrate to respective beakers #1 through 4, recording the initial and final burette volumes in Table 1.
- Using a second burette that has been rinsed with deionized water, air bubbles removed and filled with deionized water, deliver 8, 6, 4, and 2 milliliters to beakers #1 through 4 respectively, recording the initial and final burette volumes in Table 2.
- Stir each solution with a stirring rod and dry the stirring rod between beakers. Use the remaining 0.400 M nickel (II) nitrate solution for the 5th solution.
- Calculate the concentrations of the solutions in beakers #1 through 4 using the dilution equation, . Calculate each concentration and record in Table 3.
- Pour the solutions from each beaker into separate clean, dry cuvettes labeled # 1 through 5. Additionally, fill a cuvette with deionized water for the blank. Note – the cuvettes should be cleaned by rinsing with deionized water several times. The outside of the cuvette should be cleaned of fingerprints using tissues (not paper towels). Paper towels could scratch the cuvette.
- Using the spectrophotometer, set the wavelength to 635 nm. Place the cuvette with deionized water into the spectrophotometer and set the absorbance to zero. Next, read the absorbance for each of the cuvettes #1 through 5 and record these values in Table 3.
- Obtain a sample solution of nickel (II) nitrate of unknown concentration. Fill a clean cuvette with this solution and read the absorbance of the solution.
- Construct a graph of absorbance as a function of concentration, plot concentration on the x-axis and y-axis absorbance. Perform a linear fitting of the data points. Locate the absorbance of the unknown solution on the vertical axis and determine the corresponding concentration on the x-axis. Note this value on the Data Sheet.
Data Sheet
Name:
Table 1. Volume of nickel (II) nitrate used
| Beaker # | Initial burette reading (mL) | Final burette reading (mL) | Volume of nickel (II) nitrate transferred (mL) |
| 1 | |||
| 2 | |||
| 3 | |||
| 4 |
Table 2. Volume of deionized water used
| Beaker # | Initial burette reading (mL) | Final burette reading (mL) | Volume of deionized water transferred (mL) | Total volume of solution in beaker (mL) |
| 1 | ||||
| 2 | ||||
| 3 | ||||
| 4 |
Table 3. Concentration and Absorbance of solutions
| Beaker # | Concentration (M) | Absorbance (a.u.) |
| 1 | ||
| 2 | ||
| 3 | ||
| 4 | ||
| 5 | 0.400 |
Concentration of Unknown Solution: ___________________
Experiment 11 - Determination of the Gas Constant, R
Determination of the Gas Constant, R Lab Prep Sheet
DEMO SETUP:
- 500 mL Florence flask
- Wide mouth test tube
- 400 mL beaker
- Thermometer
- Utility clamp
- Clay Triangle
- Connecting tubes
- Pinch clamp
- Stirring rod on paper towel
- Ring stand -3
- Ring -2
- Crucible
- Wire square
- Evaporating dish
Hood:
- Pipette fillers
- 12 crucibles
- 12 pinch clamps
- 12 connecting tubes
- Box of thermometers (at least 12)
- Evaporating dishes
- Barometer
- Page #D-159 Vapor Pressure of H2O
Balance Area:
- 3-KClO3 on paper towel with scoop-Potassium chlorate
- 3-MnO2 on paper towel with scoop-Manganese @dioxide
Determination of the Gas Constant, R
Purpose
The purpose of this experiment is to determine a value for the gas constant, R, in the ideal gas law equation.
Introduction
All real gases deviate more or less from ideal behavior because their molecules do have some slight attraction for one another and do occupy some slight volume themselves. These deviations are more pronounced at high pressures and low temperatures; however, under ordinary conditions of temperature and pressure, most gases conform to the gas law equation:
or since , then
In this experiment, the values of P, V, n, and T are obtained for oxygen and R is calculated. Oxygen is prepared by the decomposition of potassium chlorate using manganese (IV) oxide as a catalyst.
If the potassium chlorate is accurately weighed before and after oxygen has been driven off, the mass of the oxygen can be obtained by difference. Oxygen is collected by water displacement and the volume can be determined by the volume of water displaced. Atmospheric pressure conditions are imposed on the oxygen gas. Thus, by reading the barometer and using the vapor pressure of water at the temperature the measurements were made, the pressure of oxygen can be calculated using Dalton’s law of partial pressure.
Procedure
- Weigh a crucible and place about 0.1 gram (±0.05 g) in the crucible. The particular value for this mass is not needed for calculation and need not be recorded on the report sheet because the serves as a catalyst.
- Heat the crucible and the vigorously for two or three minutes. Cool to room temperature.
- Weigh an evaporating dish and place approximately 1.5 grams (±0.2 g) in it. The particular value for this mass is not needed for calculation and need not be recorded on the report sheet because mass lost will be determined by weighing by difference.
- Carefully dry the sample by gently warming the evaporating dish over a small flame. Use a wire gauze mat as shown in Figure 1. DO NOT MELT. When the particles “dance” or “crackle,” the is dry. Cool to room temperature.
Figure 1: Evaporating dish used to dry potassium chlorate
5. While the is cooling, assemble the apparatus as shown in Figure 2, but do not attach the test tube.
Figure 2: Oxygen gas production and collection apparatus.
6. Nearly fill a 500-mL Florence flask with room temperature tap water. Fill it at least halfway up the neck. Attach the fitted connecting tube. The long tube should extend to within one-half inch of the bottom of the flask and the short tube should not extend more than one-half inch below the rubber stopper. Half-fill a 400-mL beaker with room temperature tap water.
7. Before attaching the test tube, use a pipet bulb to blow air into the Florence flask through the short tube side. This will displace some water from the Florence flask and fill the tube connecting the flask and the beaker. When air ceases to bubble into the beaker, place a clamp on the hose.
8. Crush any large crystals of in the evaporating dish with a stirring rod and add the cooled to it. Mix thoroughly.
9. Use a paper funnel to transfer the solid mixture into the dry wide-mouth test tube. Wipe the mouth of the test tube with a paper towel. BE CERTAIN THAT NONE OF THE SOLID MIXTURE COMES IN CONTACT WITH THE RUBBER STOPPER OR A SEVERE EXPLOSION MAY OCCUR.
10. Wipe the outside of the test tube thoroughly to remove any fingerprints.
11. Weigh the test tube with its contents ensuring to distribute mass evenly. Weigh accurately to ±0.001 g. Record the mass on the Data Sheet.
12. Spread the contents evenly over the lower third of the test tube and attach to apparatus.
13. Open the clamp and equalize air pressure inside the system with the atmospheric pressure outside the system by raising the beaker until the water levels are at the same height above the countertop.
14. Close the clamp and discard the water in the beaker, taking care not to invert the glass tip of the rubber tubing connector. Dry the beaker.
15. Place the empty beaker back under the glass tip and open the clamp. A little water will flow into the beaker, but if the system is air-tight and has no leaks, the flow will soon stop, and the glass tip will remain filled with water. If this is not the case, and water keeps flowing, check the apparatus for leaks and repeat steps 6, 7, 12, 13, and 14. The small amount of water that is expected to flow into the beaker is part of the sample, and should be kept in the beaker. Water levels will be adjusted again in a later step to correct for this.
16. OBTAIN INSTRUCTOR PERMISSION TO PROCEED ANY FURTHER.
17. Heat the lower part of the test tube gently (BE CERTAIN THAT THE CLAMP IS OPEN) so that a slow but steady stream of gas is produced, as evidenced by the flow of water into the beaker. KEEP THE TIP UNDER WATER at all times during the experiment, for if air is admitted through the tip at any time, the results will not be accurate.
18. Until the mixture has liquefied, tap the test tube gently with a stirring rod to contain “dancing” particles in the lower third of the test tube.
19. Carefully control the evolution of oxygen by removing the flame whenever the rate becomes too rapid. IF WHITE VAPORS SHOULD APPEAR, DISCONTINUE HEATING IMMEDIATELY, and allow the smoke to settle. All smoke particles should be contained in the test tube, for if the smoke enters the connecting tube, the loss of mass from the test tube will be a sizable source of experimental error; and, if it were to obstruct the flow of oxygen through the tube, it could create a major safety hazard.
20. When 300 milliliters of water have been collected in the beaker, remove the flame, and allow the apparatus to stand until it has cooled to room temperature with the glass tip beneath the surface of the water in the beaker and the clamp open.
21. After the apparatus has reached room temperature, equalize the gas pressure inside the flask with the atmospheric pressure by raising or lowering the beaker or flask until the water levels in the beaker and the flask are the same height above the countertop. DO NOT HANDLE THE FLASK in such a way that the gas contained therein will be warmed by the hands. CLOSE THE CLAMP while the water levels are the same height.
22. Take the temperature of the oxygen in the flask by placing a thermometer directly in the gas (one inch above the level of the remaining water). Do not let the bulb of the thermometer touch the wall of the flask. Record the temperature to one decimal place.
23. Weigh the test tube and contents, avoiding fingerprints, to ±0.001 gram and record the value.
24. With 100-mL and 10-mL graduated cylinders, determine the volume of the oxygen in the flask by measuring the volume of the water in the beaker. Record the volume to one decimal place.
25. Determine atmospheric pressure for your location.
26. Look up the vapor pressure of water at the recorded temperature.
27. Calculate the pressure oxygen is exerting, that is, the oxygen’s partial pressure.
28. Calculate the gas constant, R, from the data obtained. Show the calculation with units and proper significant figures in the calculation area of the Data Sheet.
Data Sheet
Name:
Quantity | numbers and units |
| Mass of test tube and contents before heating | |
| Mass of test tube and contents after heating | |
| Mass of oxygen gas | |
| Volume of oxygen sample in milliliters | |
| Volume of oxygen sample in liters | |
| Temperature in °C | |
| Temperature in Kelvin | |
| Atmospheric pressure in torr | |
| Vapor pressure of water in torr | |
| Pressure of oxygen gas in torr | |
| Pressure of oxygen gas in atmospheres | |
| Calculated value of gas constant, R |
Show the calculation the value of the gas constant, R as:
Problem Set 1 - Nomenclature
Compound Nomenclature – Naming Compounds and Writing Formulas Activity
Purpose
This activity introduces the naming conventions for common inorganic compounds. The naming conventions are classified by the compound's nature: ionic, molecular, and acid compounds.
Part 1 – Ionic Compounds
Ionic compounds are charge-based entities. They consist of positively charged particles (cations) in an extensive alternating array with negatively charged particles (anions). When naming an ionic compound, you do not really name “the compound.” You name the two parts (cation and anion) just as if they were independent.
Most cations are metals. These cations retain their own element names. For those metals which have a single dependable charge, that element name is enough. These metals are the metals of Group IA (the alkali metals), Group IIA (the alkaline earth metals), silver, zinc, cadmium, and aluminum. All other metals can have multiple charges, and a Roman numeral is used to specify the charge. For example, copper atoms will lose either one electron to make , copper (I) ions or will lose two electrons to make , copper (II) ions. When naming an ionic compound, the charge of the metal ion is determined from the chemical formula and the charge of the anion.
Binary Ionic Compounds
A binary ionic compound will consist of a metal cation and a monoatomic ion, that is, the anion consists of an atom of a single element. Monoatomic anions are typically nonmetals and are named by changing the ending of the nonmetal name to have the suffix “-ide.” Examples are fluorine atoms gain one electron to form , fluoride anions and oxygen atoms gain two electrons to form , oxide anions.
Ternary Ionic Compounds
A ternary ionic compound will consist of a metal cation and a polyion, that is, the anion consists of two or more atoms. The most commonly occurring polyions contain two elements with one being oxygen and are sometimes referred to as “oxyanions.” These polyions are named by dropping the ending of the of the non-oxygen element and adding the suffix “-ate” or “-ite,” depending on the number of oxygen atoms contained in the polyion. These polyions also sometimes contain the prefixes “hypo-“ and “per-“, again, to help indicate the number of oxygens atoms contained in the polyion. Table 1 lists some of the most common oxyanions. Look for patterns in the number of oxygens and charges within periodic table family groups.
Table 1 – Common Polyatomic Ions
Element | Root | hypo___ite | ___ite | ___ate | per___ate |
chlorine | chlor- | ||||
bromine | brom- | ||||
iodine | iod- | ||||
sulfur | sulf- |
|
| ||
nitrogen | nitr- |
|
| ||
phosphorus | phosph- |
| |||
arsenic | arsen- |
|
| ||
carbon | carbon- |
|
|
| |
boron | bor- |
|
|
|
It is important to note that there are a few exceptions to these naming rules. Hydroxide ion, , cyanide ion, , and peroxide ion, are polyions but they all end in with the “-ide” suffix, which is usually reserved for monoatomic ions. These ion names were assigned before the naming rules were established and were so engrained in the science community; their assigned names were left as an exception to the rule.
Another species that does not fit nicely into a category is ammonium ion, . This polyatomic cation will make ionic compounds with both monoatomic and polyatomic anions. Compounds containing ammonium cation appear quite often in general chemistry courses.
When determining the chemical formula from the name both binary and ternary ionic compounds, the resulting compound must be expressed so that the result is a neutral compound, that is, the sum of the charges must add up to zero. For a compound like sodium chloride, the name implies cations and anions. These cations and anions pair up in a 1:1 ratio, and the formula is NaCl, with the charges, +1 + (-1) equaling zero. The compound iron (III) chloride implies cations and anions. The compound would require 3 chloride anions for each iron (III) cation to maintain neutrality. The number of ions in the compound are indicated by adding a subscript after the anion, so the formula for iron (III) chloride is expressed as .
When the formula requires more than one polyatomic ion, parenthesis are used to group the polyatomic ion together. The compound iron (III) chlorate implies cations and anions. The compound would require 3 chlorate anions for each iron (III) cation, so the formula for iron (III) chlorate is expressed as .
Part 2 – Molecular Compounds
Molecular compounds typically are composed of nonmetals. Unlike ionic compounds where cations and anions are attracted through electrostatic interactions, molecular compounds form covalent bonds. In a covalent bond, two atoms share electrons. Nonmetals have high electron affinities and tend to attract electrons to themselves. In a compound comprised of two nonmetals, neither wants to give up electrons to the other, and the result is the sharing of electrons.
The naming convention for binary molecular compounds is similar to that of the binary ionic compounds in that the name of the element listed first in the compound is named without altering the name; the element listed second is named by changing the ending of its name to have the suffix “-ide.” In addition, Greek prefixes are used to designate how many atoms of each element are present in the compound. The prefixes are listed in Table 2. The prefix mono- is often omitted, with the implication that if no prefix is listed, it is assumed there is only one atom of that element type.
An example of this naming convention is dinitrogen tetrafluoride. The prefix di- indicates the compound has two nitrogen atoms, and the tetra- prefix indicates four fluorine atoms in the compound, thus the formula would be. In the case of oxygen being listed as the second element, often the “a” at the end of the prefix is omitted. In the case of , the compound would be named tetraphosphorus decoxide, rather than decaoxide.
Table 2 – Greek Prefixes
| 1 | mono- |
| 2 | di- |
| 3 | tri- |
| 4 | tetra-- |
| 5 | penta- |
| 6 | hexa- |
| 7 | hepta- |
| 8 | octa- |
| 9 | nona- |
| 10 | deca- |
Part 3 – Acids
Acids are compounds that release hydrogen ions, H+, when dissolved in water. Acids, like ionic compounds, can be broken up into two categories: binary acids and ternary acids, and have slightly different naming conventions.
Binary Acids
A binary acid will consist of hydrogen and one other nonmetal, and the hydrogen is listed first in the compound. In gas form, these substances follow the convention for naming molecular compounds. When they are dissolved in water, they release H+ ions and act like an acid; therefore, the name is changed. To name these compounds in aqueous solution, the word hydrogen is shortened to become the prefix “hydro-,” followed by the base part of the nonmetal, and the end of the nonmetal name is replaced with “-ic acid.” Table 3 lists a few examples of binary acid names in their gas and aqueous forms.
Table 3 – Names of Binary Acids
| hydrogen fluoride | hydrofluoric acid | ||
| hydrogen chloride | hydrochloric acid | ||
| hydrogen bromide | hydrobromic acid | ||
| 1 | hydrogen cyanide | hydrocyanic acid |
1 Cyanide ion, , is a polyion but as it is historically named as a monoatomic ion with an ending of “-ide,” it gets placed into the binary acid category as well.
Ternary Acids
A ternary acid will consist of a hydrogen and a polyatomic ion, typically the oxyanions, and like the binary acids, the hydrogen is typically listed first. To name these compounds, start with the name of the polyatomic ion, if the polyatomic ion ends in “-ate” change the ending to “-ic acid” and if the polyatomic ion ends in “-ite” change the ending to “-ous acid.” An example is . The polyatomic ion in this compound is the sulfate ion, Remove the “-ate” and replace with “ic acid” to get the name sulfuric acid.
See attached PDF for Naming problem sets.
Problem Set 2 - Oxidation Numbers
Oxidation Numbers
Purpose
This activity introduces the rules for assigning oxidation numbers to each element in a compound or ion.
Introduction
Oxidation numbers are useful for naming compounds and balancing oxidation-reduction reactions. The oxidation number is defined as the total number of electrons an atom has gained or lost in a chemical reaction to form a bond. Oxidation numbers can be calculated for covalent compounds even though electrons are not gained or lost, but rather shared. The oxidation number in covalent compounds represents the magnitude of relative electron density associated with each atom in a covalent compound. The oxidation number of an atom in a covalent compound indicates what the charges would be if that atom gained or lost the electrons involved in the bonding.
Oxidation Number Rules
The rules for assigning oxidation numbers are as follows, and should be considered in this order:
- The oxidation number of an element in its elemental form is zero. Examples: Fe, Al, H2, N2, O2, F2, Cl2, Br2, I2, P4, S8
- The oxidation numbers in any chemical species must sum to the overall charge on the species. That is, the oxidation numbers for a neutral compound must sum to zero and the oxidation numbers for a polyion must sum to the charge on the ion.
- The oxidation number of a monoatomic ion is equal to the charge on the ion.
- Fluorine has an oxidation number of -1 when in a compound.
- Metals of Group IA have an oxidation number of +1 when in a compound. Metals of Group IIA have an oxidation number of +2.
- When in a compound, silver, zinc, cadmium, and aluminum will have consistent oxidation numbers. Ag will be +1, Zn and Cd will be +2 and Al will be +3.
- Hydrogen has an oxidation number of +1 when it combines with nonmetals and polyions. Hydrogen has an oxidation number of –1 when it combines with metals. In these cases, hydrogen will be listed as the second element in the compound and will be named as “hydride.”
- Oxygen will usually have an oxidation number of -2. There are a few cases where oxygen combines with hydrogen or a Group IA Metal and it has an oxidation number of -1 (rules 5 and 6 take precedence), and forms the ion and is named as “peroxide.” In rare cases, oxygen will have an oxidation number of -½ and is named as “superoxide.”
- Nonmetals of Group VIIA, (halogens) have an oxidation number of –1 when they are listed second in a binary compound. When listed first in a binary molecule or in a polyion, their oxidation number can vary from +1 to +7.
- Nonmetals of Group VIA (chalcogens) have an oxidation number of –2 when they are listed second in a binary compound. When listed first in a binary molecular or in a polyion, their oxidation number can vary.
- Nonmetals of Group VA (pnictogens) have an oxidation number of –3 when they are listed second in a binary compound. When listed first in a binary molecular or in a polyion, their oxidation number can vary.
Examples
Example 1: magnesium sulfate, . Determine the oxidation numbers for P and S in the compound.
- According to rule 2, the oxidation number for all the atoms in the compound must sum up to zero. That is, Mg + S + 4O = 0. Likewise, the sum of the oxidation numbers of the atoms in the polyion (sulfate) must add up to the charge on the polyion. That is, S + 2O = -2.
- According to rule 3, magnesium will have an oxidation number of +2 since it is a Group IIA metal.
- According to rule 7, oxygen will have an oxidation number of -2 since it is oxygen in a polyion.
- According to rule 9, the oxidation for sulfur in a polyion can vary. Using either of the relationships listed above, the oxidation number of the sulfur can be calculated since the oxidation numbers of magnesium and oxygen are now known.
- By assigning Mg = +2 and O = -2, it is determined that S = +6.
Example 2: diphosphorus pentasulfide, . Determine the oxidation numbers for P and S in the compound.
- According to rule 2, the oxidation number for all the atoms in the compound must sum up to zero. That is, 2P + 5S = 0.
- According to rule 9, the oxidation number for sulfur will be -2 since it is listed second in the binary compound.
- According to rule 10, The oxidation number for phosphorus can vary since it is listed first.
- By assigning S = -2, it is determined that P = +5.
Oxidation Number Exercises
Name:
Determine correct IUPAC name for each of the following formulas and then determine the oxidation number for requested element the compound or ion.
| Formula | Name | Oxidation Number |
| Zn = | ||
| H = | ||
| Br = | ||
| I = | ||
| Ti = | ||
| S = | ||
| P = | ||
| I = | ||
| Cr = | ||
| Ni = | ||
| Mn = | ||
| P = | ||
| Cr = | ||
| P = | ||
| H = | ||
| N= | ||
| I = | ||
| S = | ||
| P = | ||
| Mn = | ||
| N = | ||
| I = | ||
| Sn = | ||
| O = | ||
| I = |
Determine correct formula for each of the following IUPAC names and then determine the oxidation number for requested element the compound or ion.
| Name | Formula | Oxidation Number |
| Cr= | |
| Br = | |
| V = | |
| C = | |
| Mn = | |
| P = | |
| Cl = | |
| P = | |
| Cr = | |
| Hg = | |
| Cl = | |
| Te = | |
| Se = | |
| Cl = | |
| O = | |
| C = | |
| P = | |
| I = | |
| Br = | |
| S = | |
| P = | |
| H = | |
| N = | |
| Cd = | |
| I = |
Problem Set 3 - Molecular Geometry
Molecular Geometry
Purpose
The purpose of this exercise is to practice drawing Lewis structures of substances, determining electron domain and molecular geometry, determining bond polarity and molecular polarity, and determining bond angles.
Instructions
For each of the following substances draw the Lewis dot structure, then identify the electron domain geometry and molecular geometry (Table 1). Use the electronegativity chart (Table 2) to determine bond polarity. Determine if the molecule is polar or nonpolar using the molecular geometry and bond polarity. If the substance is charged, write “ion” in the molecular polarity square. Finally, record the approximate bond angles.
| Substance(Formula) | Lewis Dot Structure | Electron Domain Geometry | Molecular Geometry | Bond Polarity | Molecular Polarity | Bond Angles |
| Bromine | ||||||
| Hydrogen Chloride | ||||||
| Boron Trichloride | ||||||
| Nitrate Ion | ||||||
| Phosphorous Pentafluoride | ||||||
| Substance | Lewis Dot Structure | Electron Domain Geometry | Molecular Geometry | Bond Polarity | Molecular Polarity | Bond Angles |
| Ammonium Ion | ||||||
| Formaldehyde | ||||||
| Ammonia | ||||||
| Triiodide Ion | ||||||
| Chlorine Trifluoride | ||||||
| Substance | Lewis Dot Structure | Electron Domain Geometry | Molecular Geometry | Bond Polarity | Molecular Polarity | Bond Angles |
| Nitrite Ion | ||||||
| Carbon Tetrachloride | ||||||
| Sulfur Trioxide | ||||||
| Sulfite Ion | ||||||
| Ethylene |
Table 1: Electron Domain and Molecular Geometry
| Number of Bonding Domains | Number of Non-Bonding Domains | Electron Domain Geometry | Molecular Geometry | Bond Angles |
| 2 | 0 | Linear | Linear | 180° |
| 3 | 0 | Trigonal Planar | Trigonal Planar | 120° |
| 2 | 1 | Trigonal Planar | Bent | <120° |
| 4 | 0 | Tetrahedral | Tetrahedral | 109.5° |
| 3 | 1 | Tetrahedral | Trigonal Pyramidal | <109.5° |
| 2 | 2 | Tetrahedral | Bent | <109.5° |
| 5 | 0 | Trigonal bipyramidal | Trigonal Bipyramidal | 120°, 90° |
| 4 | 1 | Trigonal bipyramidal | Seesaw | <120°, <90° |
| 3 | 2 | Trigonal bipyramidal | T-shaped | <90° |
| 2 | 3 | Trigonal bipyramidal | Linear | 180° |
| 6 | 0 | Octahedral | Octahedral | 90° |
| 5 | 1 | Octahedral | Square Pyramidal | <90° |
| 4 | 2 | Octahedral | Square Planar | 90° |
Table 2: Selected Electronegativity Values
| Element | Electronegativity |
| Boron | 2.0 |
| Bromine | 2.8 |
| Carbon | 2.5 |
| Chlorine | 3.0 |
| Fluorine | 4.0 |
| Hydrogen | 2.1 |
| Iodine | 2.5 |
| Nitrogen | 3.0 |
| Oxygen | 3.5 |
| Phosphorous | 2.1 |
| Sulfur | 2.5 |
Problem Set 4 - Stoichiometry
Stoichiometry Problem Set
Purpose
This problem set practices calculation techniques for stoichiometry, theoretical yield, percent yield, percent purity, and limiting reactant determination.
- Use the balanced equation for the combustion of acetic acid in oxygen to produce carbon dioxide and water to calculate the answers for problems b through e.
a. Balance:
b. Calculate the number of moles of water produced from 1.50 moles of acetic acid reacting with sufficient oxygen.
c. Calculate the number of moles of acetic acid reacting with sufficient oxygen required to produce 6.2 moles of carbon dioxide.
d. Calculate the number of moles of oxygen required to react with 15.00 grams of acetic acid.
e. Calculate the number of molecules of carbon dioxide produced by reacting sufficient oxygen with 6.005 grams of acetic acid.
2. Iron metal reacts with oxygen to form iron (III) oxide, .
a. Write the balanced equation for this reaction.
b. If 25.00 grams of iron completely react, what mass of iron (III) oxide can form?
c. If 2.50 × 1023 molecules of oxygen completely react, what mass of iron (III) oxide can form?
d. If 25.00 grams of iron react with 2.50 × 1023 molecules of oxygen, what mass of iron (III) oxide can form?
e. If 30.45 grams of iron (III) oxide are produced in the reaction in part d, what is the percent yield?
3. Phosphoric acid can be prepared from elemental phosphorus, oxygen, and water by the following series of reactions. Be sure to balance the following reactions.
What volume of acid is produced from reacting 1.250 grams of phosphorus with sufficient oxygen and water? The specific gravity of phosphoric acid is 1.18.
4. Potassium chlorate decomposes on heating to form potassium chloride and oxygen.
a. Write the balanced equation for the reaction.
b. If 4.904 grams of potassium chlorate, with a 75.00% purity, is heated, what mass of oxygen may be produced?
c. 1.208 grams of oxygen are experimentally obtained. What is the percent yield?
5. What mass of magnesium oxide can be produced by reacting 0.750 grams of magnesium with 2.00 × 1022 molecules of oxygen? Remember to write the balanced equation for the reaction.
6. Ethyl alcohol, may be prepared by the fermentation of glucose, , as indicated by the unbalanced equation below.
If 3.856 mL of ethyl alcohol was collected from an impure 7.000-gram sample by this formation process, determine the percent purity for this glucose sample. The specific gravity of ethyl alcohol is 0.790.
Problem Set 5 - Gas Laws Problem Set
Gas Laws Problem Set
Purpose
This problem set practices the use of the several different gas laws: Boyle’s, Charles’, combined, ideal, Dalton’s, and Graham’s laws.
- Calculate the ratio of the rates of diffusion of neon to chlorine at room temperature. Which gas diffuses faster? How much faster?
- What is the density of fluorine gas at STP?
- A given mass of carbon dioxide has a volume of 250.0 milliliters at STP. To what temperature in degrees Celsius must the gas be heated in order to occupy a volume of 750.0 milliliters at 748.0 torr?
- Oxygen gas has the density of 1.43 g/L at STP. What will be the mass of 500.0 milliliters of oxygen at 21.8°C and 750.8 torr?
- How many liters of oxygen collected over water and measured at 27.0°C and 745.0 torr can be prepared by decomposing 4.421 grams gold (I) oxide (a noble metal oxide)? The vapor pressure of water at 27.0°C is 26.7 torr.
- Compute the volume in milliliters of 80.018 grams of carbon dioxide at STP.
- Analysis of a volatile liquid shows it contains 62.04% carbon, 10.41% hydrogen and some oxygen by mass. At 150.0°C and 760.0 torr, 500.0 milliliters of vapor have a mass of 1.673 grams. What is the molar mass of the compound and its chemical formula?
- Chlorine diffuses 2.19 times faster than gas Q. What is the molar mass of gas Q?
- A vessel contains 10.008 grams helium, 30.270 grams neon and 80.001 grams oxygen and exerts a pressure of 4.18 atm at a certain temperature. Calculate the partial pressure of each gas in torr.
- Given the gas phase reaction: nitrogen + oxygen yields dinitrogen pentoxide, if 106.4 grams of nitrogen are reacted with sufficient oxygen, calculate the following:
- Moles of nitrogen reacted
- Moles of oxygen reacted
- Moles of dinitrogen pentoxide produced
- Liters of nitrogen (STP) reacted
- Liters of oxygen (STP) reacted
- Liters of dinitrogen pentoxide (STP) produced
- Molecules of oxygen reacted
- Molecules of dinitrogen pentoxide produced
Problem Set 6 - Lecture Review Problem Set
Lecture Review Problem Set
Purpose
This problem set practices concepts covered throughout General Chemistry I and aid in review for the final exam.
- After an object with mass 32.118 grams is placed in a graduated cylinder containing water, the water level rises to 47.13 milliliters. The density of the object is 10.8 grams per cubic centimeter. Determine the initial level of the water in the graduated cylinder.
- If 28.7 grams of iron at 126.2ºC is placed in 725 milliliters of water at 23.6ºC in an insulated container, what will be the final temperature? The specific heat of iron is 0.444 J/g∙℃ and specific heat of water is 4.184 J/g∙℃.
- What is the percent composition by mass of oxygen in rubidium dichromate?
- The pretend isotope chemogen, Ch, has the following data:
| Isotope | Atom mass (a.m.u.) | Percent abundance |
| Ch-115 | 115.0003 | |
| Ch-114 | 113.9893 | 18.77 |
| Ch-113 | 112.9990 | 50.24 |
Determine the average atomic mass of chemogen to one decimal place.
5. “Antimonal saffron” is a pigment used in paint. A 1.117-gram sample of “antimonal saffron” was found to contain 60.30% of the element antimony. The remainder was sulfur. The true molar mass is known to be 404 grams per mole. Determine the empirical and molecular formulas of this compound.
6. Ethyl alcohol, may be prepared by the fermentation of glucose, , as indicated by the unbalanced equation below.
What volume of ethyl alcohol may be prepared from 0.200 kilograms of glucose that is 95.6% pure? The specific gravity of ethyl alcohol is 0.791.
7. What mass of aluminum oxide can be produced by reacting 6.248 grams of aluminum with molecules of oxygen? Hint – write the balanced equation first.
8. Use the thermochemical equations below to determine the standard heat of formation, , at 25ºC for
9. If 17.37 milliliters of 0.4000 M acid solution are added to 20.00 milliliters of water, what is the new concentration?
If 10.00 milliliters of this new solution are added to enough water to make 50.00 milliliters of acid solution, what is the new concentration
10. What volume of oxygen collected over water and measured at 26.1ºC and 754.8 torr can be prepared by decomposing 1.226 grams of potassium chlorate? The vapor pressure of water at 26.1ºC is 25.21 torr.
11. A 1.235-gram sample of a hydride of boron was found to contain 1.001 grams of boron. At 23.0ºC and 765.0 torr, 51.15 milliliters of this gas have a mass of 0.113 g. What is the molecular formula for the compound?
12. 8.08 grams of hydrogen, 48.0 gram of methane, and 41.90 grams of krypton in a closed vessel exert a pressure of 3.00 atmospheres at a certain temperature. Calculate the partial pressure of each gas in units of torr.
13. Using the values in the table below, calculate the amount of heat needed to convert 225.0 grams of ice at -5.0ºC to water at 98.0ºC.
Standard enthalpy of phase change (J/g) | Specific heat of substance |
| Fusion of water = 334 J/g | Ice = 2.092 |
| Vaporization of water = 2260 J/g | Water = 4.184 |