The Alcohol Content of Whiskey

Introduction

The purpose of this experiment was to determine the alcohol content of a sample of whiskey and a sample of vodka. To achieve the purpose of the experiment, the entire class worked to obtain the densities of a wide range of ethanol/water mixtures. These densities were put into a plot that related them to their percent composition, and were compared to the densities of the whiskey and vodka samples.

Procedure

Different alcohol/water solutions were assigned to each student. Volumetric pipets were used to prepare three 10.0 mL samples that composed of 4.5 mL ethanol and 5.5 mL water, creating the assigned 45% ethanol solution. The samples were combined in a beaker and covered and then 5 mL were withdrawn using the volumetric pipet and put into pre-weighed vials. The vials were weighed again and their masses recorded. For each of the 5 mL solution samples, the mass without the vial and the density were calculated, along with the average density and the standard deviation using a TI-83 Plus calculator. The data of each of the different assigned ethanol/water solutions, the whiskey, and the vodka, were combined and put onto an Excel document, and made available to the class. The average densities and standard deviation of the whiskey and vodka samples were calculated on Excel. A scatter plot was produced of the density vs the percent composition on the ethanol/water solutions. It was determined that the data was linear so an equation was calculated and plotted using Excel that best fitted the to the data. The calculated densities of the whiskey and vodka where plugged into the equation to calculate the alcohol content of the whiskey and vodka.

Detailed procedures may be found in reference 1.

Results

Figure 1 shows a plot of the densities calculated of the samples of different ethanol/water compositions. Figure 1 is important because it allowed me to easily see if there was a pattern in the data, to see how close that data was, and use the equation of the best-fit to calculate alcohol content of the whiskey and vodka samples.

Figure 1. Scatter plot of the calculated densities of the ethanol/water samples.

From the calculated average densities of the whiskey and vodka and the equation of best-fit line in Figure 1, the alcohol concentration of the whiskey was calculated to be 36.2% (v/v) and the vodka’s  to be 35.7% (v/v). The calculated PRE of the whiskey is 9.63% and the calculated PRE of the vodka is 10.9%.

Discussion

The the plot (figure 1) had a respectable R2 value of 0.9632 and there were no outliers, but the majority of the points were not directly on the best-fit line. The points were linear pattern, the average density decreased as the ethanol concentration increased. The concentration of alcohol in the sample of whiskey (36.2% (v/v)) and the sample of vodka (35.7% (v/v)) as calculated from the best-fit line in figure 1, were less than the reported concentration of alcohol (40% (v/v)) given on the label of the bottles. The PRE of the whiskey was 9.63% and the PRE of the vodka was 10.9%. The differences in the calculated concentrations and the reported may have been caused by not making precise measurements, and the assumption that the whiskey and vodka only contained ethanol and water. Because the calculated and the reported concentrations of the whiskey and vodka were different, the fundamental assumption that they only contain alcohol and water is not valid. The data points in figure 1 appeared to be fairly precise, which made the rejection of that assumption more valid.

References

1. General Chemistry Experiments: A Manual for Chemistry 204, 205, and 206, Department of Chemistry, Southern Oregon University: Ashland, OR, 2009

Determination of Water Density

Introduction

The purpose of this experiment was to compare the precision of some of the laboratory instruments and determine the density of purified water. The density of water was determined using three different volume-measuring devises. Five milliliters of water was measured using the three different volume-measuring devises three times each, and was then weighed on the analytical balance.

Procedure

Three vials were cleaned and then weighed on the analytical balance. Using a 100 mL graduated cylinder, five milliliters of water were measured and poured into each of the vials. The vials were then weighed on the analytical balance again and their masses recorded. The vials were dried, and the steps repeated using a 100 mL graduated pipet and again using a volumetric pipet. The density of water was calculated for masses of the three different measuring devises. The densities were then compared to each other and to the actual density of water.

Results

Tables one, two, and three show the mass of the vials and the mass of the vials containing water that was measured using the three different volume-measuring devises.

Table 1: Data obtained using a 100 mL graduated cylinder to measure a 5.0 mL sample of water. 

Trial	Mass of vial (g)	Mass of vial and water (g)
1	19.8155	24.745
2	20.6247	23.7373
3	21.2078	24.848

Table 2: Data obtained using a 100 mL graduated pipet to measure a 5.0 mL sample of water.

Trial	Mass of vial (g)	Mass of vial and water (g)
1	19.8155	24.9443
2	20.6247	25.4928
3	21.2078	26.1954

Table 3: Data obtained using a 100 mL volumetric cylinder to measure a 5.0 mL sample of water.

Trial	Mass of vial (g)	Mass of vial and water (g)
1	19.8155	24.7281
2	20.6247	25.2298
3	21.2078	25.6645

After the masses were determined, the densities of the data were calculated with the Microsoft Excel program. The Excel worksheet is stapled to this report.

Table 4: Average density water obtained from each devise. 

Measuring devise	Average density (g/mg)
100 mL Graduated Cylinder	0.9859 + or - 0.1869
100 mL Graduated Pipet	0.99897 + or - 0.02610
5.00 mL Volumetric Pipet	0.98252 + or - 0.04651

Discussion

The lab was fairly successful in accurately determining the density of water and the precision of the instruments. The second set of data, where the graduated pipet was used to measure the water volume, is the most accurate with a percent relative error of 0.097 when compared with the given density of water. The data where the graduated cylinder was used is slightly less accurate with a percent relative error of 1.21. This is most likely because it is more difficult to have precise measurements using a graduated cylinder. The data collected from the volumetric pipet is the least accurate with a percent relative error of 1.55. This is most likely because the instrument takes more practice and skill to use correctly than the other instruments, or it is not as precise as the other measuring instruments. 

References

1. General Chemistry Experiments: A Manual for Chemistry 204, 205, and 206, Department of Chemistry, Southern Oregon University: Ashland, OR, 2009

Molecular Modeling with Spartan: Polyatomic Molecules, VSEPR, Localized and Delocalized Bonding

Introduction
Spartan Student Edition was used to construct and examine molecular diagrams of H2O, NH3, CH4, and SF4. The properties examined include electron density, dipole moments, electrostatic potential maps, and equilibrium geometries. The geometries calculated with spartan were compared to the VSPER calculations. Spartan was used to examine the electron density, electrostatic potentials, and electrostatic charges of NO2 , NO2+, and NO2-, and compare them to the predictions made by using Lewis structures.
Procedure
The H2O molecule was created and the electrostatic charges were displayed for the atoms. All equilibrium geometry calculations were done at the B3LYP level in a 6-31G* basis set using Spartan Student Edition. The 0.002 electron/Å3 isodensity surface was calculated and displayed for H2O. Then a 0.08 electron/Å3 isodensity surface was calculated, displayed, and used to observe the electron density. An electrostatic potential map was created, made transparent, and legend displayed. To visualize the volume taken up by the lone pairs, a potential isodensity surface was calculated at the value of -83.68 kJ/mol. The angles and distances between the atoms were calculated. This procedure was repeated on the NH3 and CH4 molecules.
An electrostatic potential map was created and displayed for SF4. The bond angles and lengths were found and created. The molecules energy (au) was found and recorded. Another SF4 molecule was created which had three fluorine atoms in equatorial positions and one in an axial position, and the same procedure was done with it.
An electrostatic potential map was created for the NO2 molecule and the electrostatic charges displayed. An 0.08 electron/Å3 isodensity map was displayed. The bond lengths and angles were measured and recorded. The same procedure was followed on NO2+ and NO2- as was NO2.
Detailed procedures can be found in reference 1.

Results
The electron density of water was not distributed evenly throughout the molecule, as can be seen in Figure 1. The highest electron density was near the oxygen and the bonds where the red and green areas were.

Figure 1. Electrostatic Potential Diagram of Water Molecule.





The electron density of ammonia was not distributed evenly throughout the molecule, as can be seen in Figure 2.

Figure 2. Electrostatic Potential Diagram of Ammonia.
The electron density of Methane was not distributed evenly throughout the molecule but was symmetrical, as can be seen in Figure 3.

Figure 3. Electrostatic Potential Diagram of Methane.
The bond angles, bond lengths, and dipole moments were calculated using spartan, and the bond lengths predicted by the VSEPR theory. This data can be found in Table 1.
Table 1. The bond angles, bond lengths, and dipole moments calculated using spartan, and the bond lengths predicted by the VSEPR theory.
Molecule Angles (degrees) VSEPR Angles (degrees) Bond Length (Å) Dipole Moment (D)
H2O 103.70° 109.5° 0.969 2.09
NH3 105.76° 109.5° 1.019 1.91
CH4 109.47° 109.5° 1.093 0.00
SF4 equilateral: 102.09°
axial: 87.29° 120°
90° 1.595
1.672 0.89
SF4 (w/ 3 F atoms in equilateral positions) equilateral: 119.14°
axial: 84.66° 120°
90° 1.679
1.067 0.85
NO2 133.82° 120° 1.203 -
NO2+ 179.99° 180° 1.129 -
NO2- 129.98° 120° 1.305 -
1.1 The bond angels in the molecules increase from water, to ammonia, to methane. This is because the bond angles decrease when more lone pairs are present. The lone pairs push the other bonds because they take up more space than bonded electrons.
1.2 The bond angles of the three molecules, found using the VSPER theory are all 109.5° because they are tetrahedrals. The bond angles found using Spartan are different because it into account the effects of lone pairs and double bonds.
1.3 The bond lengths are shorter in water than ammonia, and longer in methane than ammonia. This is because the electronegativity of the central atom, the greater the electronegativity the closer it pulls the bonding atom.
1.4 Electrostatic charges exist in water, ammonia and methane because each atom has a partial charge. Electronegativity causes the electrons in the molecule to be attracted to certain atoms, giving the atoms different partial charges and making the atom polar.
1.5 The dipole moment in water is larger than the one in ammonia, and methane doesn’t have a dipole moment. Each atom’s electronegativity along with the molecules lone pairs and the dimensions of the molecule effect the molecules electrostatic charge. The electrostatic charge can be represented by a dipole moment, which shows the size and direction of the force.
1.6 A molecular dipole moment of zero does not mean that there is not a separation of charge. There are often positively charged atoms and negatively charged atoms, but the charges are distributed evenly throughout the molecule.

The bond angles in SF4 found using Spartan were significantly different than than the angles predicted using the VSEPR theory. The bond angles in SF4 with three fluorine atoms in equilateral positions are close to the angles predicted using the VSEPR theory.

The electron density in the NO2/NO2+/NO2- molecules appeared different in each molecule and were consistent with the predictions made using the lewis structures. All of these predictions were close the calculations made using spartan. In the three molecules, NO2+ has the shortest bond length followed by NO2 and NO2- with the longest. NO2+ has the largest angle, followed by NO2, and NO2- has the smallest.

Discussion
The bond lengths and angels in the molecules increase from water, to ammonia, to methane. The bond lengths act in that manner because the bond angles get smaller when there are more lone pairs in the molecule, which repel the other atoms. The lone pairs push the other bonds because they are not localized and take up more space than a bond. The VSEPR theory defines all three of these molecules tetrahedrals and doesn’t take into account the lone pairs, which is why it’s angle predictions are not precise in many cases. The bond lengths are different because the electronegativity of the central atoms are not the same. The greater the electronegativity, the closer it pulls the bonding electron. Electronegativity also the causes these molecules to have electrostatic charges. In some molecules the charges were identical but went in opposite directions and canceled out. The difference in the equilateral and axial bond lengths were due to the atoms positions around the central atom, and the repulsion from the lone pairs. The differences in the bond angles and lengths from spartan and VSEPR demonstrate the importance of considering factors such as repulsion from lone pairs and atom sizes.

References
1. General Chemistry Experiments: A Manual for Chemistry 204, 205, and 206, Southern Oregon University: Ashland OR, Fall 2010

2. Brown, LeMay, Burnsten, Murphy. Chemistry: The Central Science. 11th Edition. Pearson Prentice Hall. Upper Saddle River, NJ. 2009.