Periodic Properties of The Elements

Introduction

The purpose of this experiment was to determine the similarities and differences in the chemical and physical properties of elements and elemental groups. This was achieved by the observation and comparison of chemical reactions of different elements. 

Results

Group IA

Reactions in Air:

A piece of lithium, sodium, and potassium we cut with a spatula and introduced with oxygen.  The elements appeared to be metallic at first, but over time a white color formed where the element was cut. The balanced chemical reactions for these metals in air are as follows:

4Na(s) + O₂(g) 2Na₂O(s)

4K(s) + O₂(g) 2K₂O(s)

4Li(s) + O₂(g) 2Li₂O(s)

Potassium was the fastest element to have this reaction  Sodium was in the middle and Lithium was the slowest.

Reactions in Water:

Pieces of the alkali metals were placed into separate beakers of distilled water.  Lithium and sodium quickly dissolved in the water creating gas emissions.  Potassium sparked and lit on fire and created gas when it touched the water.  All three alkali metals left white precipitant at the bottom of the beaker.  The pH balances were then tested to find if the solution was basic or acidic.

2Na(s) + 2H₂O(l)  2Na(OH)(aq) + H₂(g)

2K(s) + H₂O(l) 2K(OH)(aq) + H₂(g)

2Li(s) + 2H₂O(l) 2Li(OH)(aq) + H₂(g)

The pH paper showed a basic pH for the alkaline earth metals. This is because alkaline earth metal’s valence electrons only have one electron so the atom easily gives that electron away.

If cesium was added to water it would react the same as potassium; that as soon as it touched water it would ignite and explode, but the reaction may be stronger.

The reactivity of these metals increases as you go down the group.  As the atom grows larger it becomes more reactive since the electrons are farther away from the nucleus.

Group IIA

Reactions in Water:

Pieces of solid magnesium and calcium were put in test tubes of water.  Magnesium had no reaction with room temperature water or with boiling water.  The calcium bubbled and formed gas and white precipitate.  The pH balances were then tested.

Mg(s) + H₂O(l) NR

Ca(s) + 2H₂O(l) Ca(OH)₂ + H₂(g)

Reactions in Hydrochloric Acid:

Magnesium and calcium were placed in test tubes with hydrochloric acid.  Both metals bubbled.  The magnesium completely dissolved, while the calcium formed a white precipitant at the bottom of the test tube.  The pH was tested.

Mg(s) + 2HCl(aq) MgCl₂(aq) + H₂(g)

Ca(s) + 2HCl(aq) CaCl₂(s) + H₂(g)

Reactions in heat/oxygen:

Two more pieces of the alkaline earth metals were subjected to heat and oxygen.  Each piece of metal ignited and burned.  The calcium took longer to ignite then the magnesium.  The magnesium was then placed into water and the pH was then tested. 

2Mg(s) + O₂(g) 2MgO(s)

MgO(s) + H₂O(l) Mg(OH)₂(aq)

2Ca(s) + O₂(g) 2CaO(s)

CaO(s) + H₂O(l) Ca(OH)₂(aq)

Calcium is more reactive than the magnesium, but they had similar reactions and pH.  

They vary only slightly. Most of the reactions were the same but were slightly stronger with the calcium and the calcium oxide reacted stronger than the magnesium. The elements are more reactive as one descends the column.

The reactions of this group were different, but just as strong as those done with the IA group. They were less explosive, but had impressive reactions.

Group IIIA

Reactions in Water:

Boron and aluminum were placed into test tubes of water.  The aluminum created effervescence while the boron had no reaction. 

2Al(s) + 6H₂O(l) 2Al(OH)₃(aq) + 3 H₂(g)

B(s) + H₂O(l) NR

Reactions in Hydrochloric Acid:

The same was done only instead of water the elements were introduced to hydrochloric acid.  The aluminum produced effervescence and the boron had no reaction.

2Al(s) + 6HCl(aq) 2AlCl₃(aq) + 3H₂(g)

Br(s) + HCl(aq) NR

Reactions in Sodium Hydroxide:

A piece of aluminum and of boron were placed into a test tube of sodium hydroxide.  The aluminum created effervescence and the boron had no reaction.

2Al(s) + 2NaOH(aq) + 6H₂O 2NaAl(OH)₄ (aq) + 3H₂(g)

Ba(s) + NaOH(aq) NR

Boron is not active compared to aluminum.  Aluminum is beneath boron, which means it’s larger in size becoming more reactive.  

The elements became less reactive in water the farther towards the right of the periodic table they are, and they become more acidic then basic.

Group IVA

Reaction in Water:

A piece of carbon, tin, and lead were placed in test tubes of distilled water.  Five minutes were allowed for a reaction to occur, and none of them reacted in that time.  

C(s) + H₂O(l) NR

Sn(s) + H₂O(l) → NR

Pb(s) + H₂O(l) → NR

Reaction with Hydrochloric Acid:

Another piece of carbon, tin, and lead were placed in separate test tubes containing hydrochloric acid.  Then the test tubes were placed in a hot water bath.  Tin and lead slowly created effervescence, while carbon had no reaction

C(s) + HCl(aq) →  CCl₂(aq) + H₂(g)

Sn(s) + HCl(aq) → SnCl₂(aq) + H₂(g)

Pb(s) + HCl(aq) → PbCl₂(aq) + H₂(g)

All three elements are non-reactive in water, but as they mix with acids they become reactive.  Lead is the most reactive since it is the largest of these elements in this elemental group.  

As the element closer to the right side of the table are less reactive since the valence shell progressively fills up, making it harder for other atoms to pull electrons from the valence shell.

Group VA 

Reactions in oxygen/heat:

A small portion of phosphorus in a deflagrating spoon was ignited by a gas burner.  The emitted smoke was then put in a flask that had distilled water at the bottom.  The captured smoke was mixed with the water, then the pH tested.

P₄(s) + O₂(g) → P₄O₁₀(g)

P₄O₁₀(g) + 6H₂O(l) → 4H₃PO₄(aq)

The pH paper resulted in being acidic.  It reacted this way because phosphorus is three electrons away from having a full valance electron so it doesn’t react with water easily. 

Reaction in Hydrochloric Acid and Water:

Pieces of bismuth were introduced into two test tubes; containing water and the other containing hydrochloric acid.  After a few minutes no reaction occurred.

Bi(s) + H₂O(l) → NR

Bi(s) + HCl(aq) → NR

Group VIA

Reaction in Oxygen/heat:

A small quantity of sulfur was introduced to burner by a deflagrating spoon.  Once the sulfur was ignited the spoon was put into a flask containing a small amount of distilled water in the bottom.  The smoke contained was then mixed with the water and the pH balance was tested.

S₈(s) + O₂(g) → 8SO₂(g)

SO₂(g) + H₂O(l) → H₂SO₄(aq)

S₈(s) + O₂(g) → SO₃(aq)

SO₃ + H₂O(l) → H₂SO₃(aq)

The pH paper resulted in being acidic.  As with phosphorus, sulfur is getting closer to having a full valance shell becoming less reactive and more acidic.

Group VIIA

Reactions in Chlorine Water:

Chlorine water was added to a test tube containing potassium bromide and another test tube containing potassium iodide.  After observing the color changes, hexane was added, to test to see if an actual chemical reaction occurred.  In the potassium bromide: a yellow color was the starting color and ended with a pale yellow color.  Potassium iodide however underwent a reaction, starting with the color brown and changing to the color purple.

KBr(aq) + Cl₂(aq) → 2KCl(aq) + Br₂(aq)

2KI(aq) + Cl₂(aq) → 2KCl(aq) + I₂(aq)

The hexane changed color because it was indicating as to whether a chemical reaction occurred or not.  When potassium iodide is mixed with chlorine water a chemical reaction will occur, but not when potassium bromide is mixed with chlorine water.

The same experiment was completed only using bromine water that was added to a test tube containing potassium iodide and potassium chloride.  The potassium iodide started out as a brown color and ended with a purple color.  The potassium chloride however started out yellow and ended in a yellow color, but still underwent a chemical reaction.

KI(aq) + Br₂(aq) → KBr(aq) + I₂(aq)

KCl(aq) + Br₂(aq) → KBr(aq) + Cl₂(aq)

The hexane changed color again, because of the result in a chemical reaction occurring.  

Based on this experiment it can be concluded that iodine is more reactive than bromine which is more reactive than chlorine.  In the experiments iodine separated it-self from the potassium, meaning that it would rather be by itself, showing that it is more stable than chlorine and bromine.

Discussion

Going down the periodic table the elements are more reactive.  This is due to the amount the atoms size is increase which causes the electrons be farther away from the nucleus making it easier for other atoms to pull the electrons away.  The elements bigger but less reactive on the right side of the table than the left side.   The atoms grow in size since the proton numbers are increase, but the valence electrons are becoming fuller, making it harder for other atoms to pull the electrons away.

As size increases the ionization energy increases, which makes an element less reactive.  This is due to the fact that since the valence sub-shell is getting fuller, there are more electrons to pull from the atom.

References

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

Periodic Properties of The Elements

Introduction

The purpose of this experiment was to determine the similarities and differences in the chemical and physical properties of elements and elemental groups. This was achieved by the observation and comparison of chemical reactions of different elements. 

Results

Group IA

Reactions in Air:

A piece of lithium, sodium, and potassium we cut with a spatula and introduced with oxygen.  The elements appeared to be metallic at first, but over time a white color formed where the element was cut. The balanced chemical reactions for these metals in air are as follows:

4Na(s) + O₂(g) 2Na₂O(s)

4K(s) + O₂(g) 2K₂O(s)

4Li(s) + O₂(g) 2Li₂O(s)

Potassium was the fastest element to have this reaction  Sodium was in the middle and Lithium was the slowest.

Reactions in Water:

Pieces of the alkali metals were placed into separate beakers of distilled water.  Lithium and sodium quickly dissolved in the water, which created gas emissions.  Potassium sparked and lit on fire and created gas when it touched the water.  All three alkali metals left white precipitant at the bottom of the beaker.  The pH balances were then tested to find if the solution was basic or acidic.

2Na(s) + 2H₂O(l)  2Na(OH)(aq) + H₂(g)

2K(s) + H₂O(l) 2K(OH)(aq) + H₂(g)

2Li(s) + 2H₂O(l) 2Li(OH)(aq) + H₂(g)

The pH paper showed a basic pH for the alkaline earth metals. This is because alkaline earth metal’s valence electrons only have one electron so the atom easily gives that electron away.

If cesium was added to water it would react the same as potassium; that as soon as it touched water it would ignite and explode, but the reaction may be stronger.

The reactivity of these metals increase as you go down the group.  As the atom grows larger it becomes more reactive since the electrons are farther away from the nucleus.

Group IIA

Reactions in Water:

Pieces of solid magnesium and calcium were put in test tubes of water.  Magnesium had no reaction with room temperature water or with boiling water.  The calcium bubbled and formed gas and white precipitate.  The pH balances were then tested.

Mg(s) + H₂O(l) NR

Ca(s) + 2H₂O(l) Ca(OH)₂ + H₂(g)

Reactions in Hydrochloric Acid:

Magnesium and calcium were placed in test tubes with hydrochloric acid.  Both metals bubbled.  The magnesium completely dissolved, while the calcium formed a white precipitant at the bottom of the test tube.  The pH was tested.

Mg(s) + 2HCl(aq) MgCl₂(aq) + H₂(g)

Ca(s) + 2HCl(aq) CaCl₂(s) + H₂(g)

Reactions in heat/oxygen:

Two more pieces of the alkaline earth metals were subjected to heat and oxygen.  Each piece of metal ignited and burned.  The calcium took longer to ignite then the magnesium.  The magnesium was then placed into water and the pH was then tested. 

2Mg(s) + O₂(g) 2MgO(s)

MgO(s) + H₂O(l) Mg(OH)₂(aq)

2Ca(s) + O₂(g) 2CaO(s)

CaO(s) + H₂O(l) Ca(OH)₂(aq)

Calcium is more reactive than the magnesium, but they had similar reactions and pH.  

They vary only slightly. Most of the reactions were the same but were slightly stronger with the calcium and the calcium oxide reacted stronger than the magnesium. The elements are more reactive as one desends the column.

The reactions of this group were different, but just as strong as those done with the IA group. They were less explosive, but had impressive reactions.

Group IIIA

Reactions in Water:

Boron and aluminum were placed into test tubes of water.  The aluminum created effervescence while the boron had no reaction. 

2Al(s) + 6H₂O(l) 2Al(OH)₃(aq) + 3 H₂(g)

B(s) + H₂O(l) NR

Reactions in Hydrochloric Acid:

The same was done only instead of water the elements were introduced to hydrochloric acid.  The aluminum produced effervescence and the boron had no reaction.

2Al(s) + 6HCl(aq) 2AlCl₃(aq) + 3H₂(g)

Br(s) + HCl(aq) NR

Reactions in Sodium Hydroxide:

A piece of aluminum and of boron were placed into a test tube of sodium hydroxide.  The aluminum created effervescence and the boron had no reaction.

2Al(s) + 2NaOH(aq) + 6H₂O 2NaAl(OH)₄ (aq) + 3H₂(g)

Ba(s) + NaOH(aq) NR

Boron is not active compared to aluminum.  Aluminum is beneath boron, which means it’s larger in size and more reactive.  

The elements became less reactive in water and are become more acidic then basic, the farther towards the right side of the periodic table they are.

Group IVA

Reaction in Water:

A piece of carbon, tin, and lead were placed in test tubes of distilled water.  Five minutes were allowed for a reaction to occur, and none of them reacted in that time.  

C(s) + H₂O(l) NR

Sn(s) + H₂O(l) NR

Pb(s) + H₂O(l) NR

Reaction with Hydrochloric Acid:

Another piece of carbon, tin, and lead were placed in separate test tubes which contained hydrochloric acid.  Then the test tubes were placed in a hot water bath.  Tin and lead slowly created effervescence, while carbon had no reaction

C(s) + HCl(aq)   CCl₂(aq) + H₂(g)

Sn(s) + HCl(aq) SnCl₂(aq) + H₂(g)

Pb(s) + HCl(aq) PbCl₂(aq) + H₂(g)

All three elements are non-reactive in water, but as they mix with acids they become reactive.  Lead is the most reactive since it is the largest of these elements in this elemental group.  

As the element closer to the right side of the table are less reactive since the valence shell progressively fills up, which makes it harder for other atoms to pull electrons from the valence shell.

Group VA 

Reactions in oxygen/heat:

A small portion of phosphorus in a deflagrating spoon was ignited by a gas burner.  The immited smoke was then put in a flask that had distilled water at the bottom.  The captured smoke was mixed with the water, then the pH tested.

P₄(s) + O₂(g) P₄O₁₀(g)

P₄O₁₀(g) + 6H₂O(l) 4H₃PO₄(aq)

The pH paper showed that the solution was acidic.  It reacted this way because phosphorus is three electrons away from having a full valance electron so it doesn’t react with water easily. 

Reaction in Hydrochloric Acid and Water:

Pieces of bismuth were introduced into two test tubes; one contained water and the other contained hydrochloric acid. No reaction occurred.

Bi(s) + H₂O(l) NR

Bi(s) + HCl(aq) NR

Group VIA

Reaction in Oxygen/heat:

A small quantity of sulfur was introduced to burner by a deflagrating spoon.  Once the sulfur was ignited the spoon was put into a flask which contained a small amount of distilled water in the bottom.  The smoke contained was then mixed with the water and the pH balance was tested.

S₈(s) + O₂(g) 8SO₂(g)

SO₂(g) + H₂O(l) H₂SO₄(aq)

S₈(s) + O₂(g) SO₃(aq)

SO₃ + H₂O(l) H₂SO₃(aq)

The pH paper resulted in being acidic.  As with phosphorus, sulfur is closer to having a full valance shell so it’s less reactive and more acidic.

The metallic oxides reacted more explosively with water and the non metallic oxides tended to react by creating gas.

Group VIIA

Reactions in Chlorine Water:

Chlorine water was added to a test tube which contained potassium bromide and another test tube which contained potassium iodide.  After observation of the color changes, hexane was added to test to see if an actual chemical reaction occurred.  In the potassium bromide: It started yellow color and ended with a pale yellow color.  Potassium iodide changed to the color pink.

KBr(aq) + Cl₂(aq) 2KCl(aq) + Br₂(aq)

2KI(aq) + Cl₂(aq) 2KCl(aq) + I₂(aq)

The hexane changed the colors to indicate whether a chemical reaction occurred or not.  When potassium iodide is mixed with chlorine water a chemical reaction will occur, but not when potassium bromide is mixed with chlorine water.

Bromine water was added to test tubes which contained potassium iodide and potassium chloride.  The potassium iodide started out as a brown color and ended with a pink color.  The potassium chloride however started out yellow and ended in a yellow color, but still underwent a chemical reaction.

KI(aq) + Br₂(aq) KBr(aq) + I₂(aq)

KCl(aq) + Br₂(aq) KBr(aq) + Cl₂(aq)

Based on the results of this experiment it can be concluded that iodine is more reactive than bromine which is more reactive then chlorine. In the experiments iodine separated its-self from the potassium, which means that it would rather be by its self and is more stable than chlorine and bromine.

Discussion

The elements are progressively more reactive farther down the periodic table they are.  This is due to the atoms increase in size, which causes the electrons be farther away from the nucleus which makes it easier for other atoms to pull the electrons away.  The elements bigger and less reactive on the right side of the table than the left side.   The atoms grow in size as the proton numbers increase, but the valence electrons are fuller, which makes it harder for other atoms to pull the electrons away.

As size increases the ionization energy increases, it makes an element less reactive.  This is due to the fact that as the valence sub-shell is gets fuller, there are more electrons to pull from the atom.

References

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

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.

Farmingville

The migration of thousands of Latino day workers to the Farmingville community has created great culture shock and debate. Education about each culture is one step towards helping the Latino workers assimilate, decrease segregation and help both sides with culture shock. Many members of the community have differing opinions and beliefs. Gaining a better understanding of these beliefs is the first step to finding a solution.
Louise says in the film that she wants to know who is around her and who lives in her neighborhood. The migrant workers have made her lose this feeling of comfort in her community. Margaret Bianculli-Dyber simply wants the migrant workers to disappear. She feels that U.S. laws have been disrespected and does not trust or appreciate that changes in Farmingville. In her quest, she does not understand nor ask, in a way that prompts answers, why these workers are in her community and how tensions can be decreased so that the community is not constantly at war. Margret does not consider herself to be racist, but is trying to make the clock turn back to a time where she did not have to worry about who is standing on her street corner. Paul Tonna, the county legislator understands that something needs to be done in Farmingville to decrease tension and supports the use of tax payer dollars to build a hiring site. If the community could find a way to educate each other about the cultures that are now in this town maybe tensions over cultural differences would decrease.
Both the day workers and the Farmingville residents are experiencing culture shock in this situation. They do not know how to interact in the other group’s culture and therefore they do not understand and are sometimes intolerant of the differences in their cultures. With education, both cultural groups would learn about where the other was from and how to communicate in a way that promotes a community feeling instead of dividing the community.
All three of the interviewees in Farmingville agree that there is a problem that started with the migrant workers moving into the town. The differences in the three are how they view the rift in the community and how they approach healing that rift. Margaret believes that the immediate deportation of all the migrant workers will instantly solve the community’s problems. Louise wants the community to be a community again, no matter who is in it. Paul Tonna wants to find a way that the migrant workers can work side by side with the legal residents of Farmingville. He does not propose allowing the situation to stay as it is but instead proposes that steps are taken to try and reach a common ground between the two cultures.
We feel that the key to better understanding this situation and finding a solution is in education and responsibility. First, all sides need to come together with an open mind and have rational conversations without loosing their tempers. If the legal Farmingville residents got to know the day laborers, why they choose to come to their community, what their intentions are, and general knowledge about their culture, that would be a start to better understanding. I think this aspect of the solution would help to put Louise at ease. A meeting like this would also help the day laborers to gain more information about the previous Farmingville culture and help them integrate it with their Latino culture.
Another aspect of the solution is to assign responsibility and hold people accountable. We feel that the Farmingville citizens, day laborers, and employers all have some responsibility in allowing this situation to escalate the way it has. The citizens have to understand that there is a need for the day laborers. They would not have come to a place where there was not a demand for manual labor. These citizens have no interest in mowing lawns or doing dishes, so to condemn the men who are willing to work hard is a serious mistake. The day migrant workers also have to realize that their presence has seriously changed the dynamic of the community. They must be held responsible for the same laws that all other community members are. Harassment, destruction of property, and other crimes are not to be tolerated from anyone. The employers need to find a balance with the work they are doing. They cannot skate around tax laws just to make a buck. These employers are changing the value and payment of manual labor. Middle ground is needed from what was done in the past to what is going on now.