Final+Final+Draft

toc =Introduction= Coordination compounds have existed since the beginning of chemistry (mostly through accidental discoveries). An example of such compounds includes Prussian blue, an accidental discovery made by Diesbach in 1704, an artists' pigments manufacturer. Frederick Augustus Genth, an assistant to Bunsen as the University of Marburg initiated the first serious study of cobalt-ammine complexes (the result of an inadvertent discovery). Instead of adding potassium hydroxide after precipitating the metals in his analytical group, Genth added ammonia to the mixture. He left the experiment, and upon his return, discovered that the mixture had crystallized. Genth, however, emigrated to America before he could do much more on the subject. In the following year, Oliver Wolcott Gibbs, an American scientist, started an investigation and worked with Genth in 1856 to publish a paper in which the described thirty five cobalt-ammine complexes. The term "coordinate compounds" (otherwise known as complex compounds) is used to describe molecules formed by the combination of ligands and metal ions. Ligands are atoms, ions, or molecules that donate an electron pair to the central metal in which it is bonding with, with the central metal including all metal compounds, except for metal vapours, plasmas, and alloys. The ligands are bound to the metal through a coordinate covalent bond, which is formed when a lone electron pair is donated to an empty metal's orbital. Coordinate compounds differ from ionic and covalent chemical bonds in that both the electrons in the bond come from the same atom. A coordinate bond is, however, indistinguishable from covalent bonds after the chemical bonding occurs. The process of forming a coordinate bond is depicted in Fig. 1 below:



The process of forming coordinate compounds can be linked to the Lewis theory. The Lewis theory was extended from the Lowry-Bronsted Theory (the theory that an acid is a proton donor and a base is a proton acceptor). A Lewis acid is a compound or atom that can accept a lone electron pair. A Lewis base is a compound that can donate an electron pair.

There are different types of unit cells which make up a lattice structure, one of them being FCC, which stands for Face Centered Cubic. FCC unit cell structures are formed when any atoms in the molecule belongs to group 9 to 11. When a compound has a coordination number of four, it usually forms a tetrahedral or square planar molecular structure. FCC unit cell structures is a characteristic of a orthorhombic crystal structure.

Coordination compounds are present in a variety of substances in everyday life, from vitamin B12, an essential element for the proper functioning of the brain and nervous system, to haemoglobin, chlorophyll, dyes, and pigments. In addition to its use in organic chemistry, coordination compounds also play an integral part in many industry processes as catalysts, and in extracting valuable metals from ores. It is therefore beneficial to gain an understanding of this important and very versatile compound in chemistry.

The purpose of this experiment is to synthesize and crystallize (form crystals from a liquid or gas) a coordinate compound and study its crystal structure.

Hypothesis
It was predicted that the unit cell structure of the synthesized complex compound would be FCC as it contains a group 11 metal, Copper. FCC crystal structures are an indication that the molecular structure of the complex compound be square planar or tetrahedral. Due to the unit cell structure of the complex compound being FCC the crystals formed should be orthorhombic in structure.

=Materials (Listed in alphabetical order by name)=

20mL of 14M Ammonia (NH3 + H2O) 1X Balance 2X 250mL beaker 1X Compound Microscope 10.0g of Copper II Sulfate Pentahydrate (CuSO4•5H2O) 1X Digital Camera 1X Dissection Kit 20mL of distilled water (H2O) 10mL of 95% Ethanol 2X Filter Paper 1X Flashlight/Lamp 1X Fume hood 1X Funnel 1X Hair ties (if needed) 1X Microscope slide 30X paper towel 1X Petri Dish 1X Retort Stand 1X Safety goggles 1X Scoopula 1X Stereo Microscope 1X Stirring rod 1X Timer

=Safety Considerations=

Ammonia (NH3(aq))
Ammonia is a colorless liquid that has a very strong and sharp odour. According to the EU (European Union) classification, ammonia is hazardous, caustic, and corrosive. Thus, when an ammonia solution is used with a beaker in a reaction, it must be accompanied with a fume hood. Gloves and safety goggles must be worn when performing an experiment with ammonia.

Copper (II) Sulphate Pentahydrate (CuSO4 · 5H2O)
Copper (II) Sulphate Pentahydrate is blue in colour, has no odour, and is non-flammable. According to the EU classification, Copper (II) Sulphate is both harmful and dangerous to the environment. It is considered a very acidic chemical and should therefore be handled with caution.

Ethanol
Ethanol is a highly flammable substance. Therefore, when handling ethanol, safety goggles are needed. Also, it is recommended to not place it near a flame source.

General safety
Wear safety goggles when performing the lab experiment. Know the location of the fire extinguisher, eye washer, and the chemical shower. Tie back long hair. Do not turn away from the Bunsen burner while it is in use. Turn off the gas when the Bunsen burner is not on operation. 

=Procedure=

1. Place 10 grams of copper(II) sulfate pentahydrate into a 250mL beaker **** 4. If the solution is too turbid, use filter paper to filter the copper (II) sulphate before proceeding (Fig. 2) **** 5. Add 20mL of 15M NH3(aq) to the solution under a fume hood **** 6. Concentrate the solution if needed by using a hot plate **** 7. Add 10mL of 95% ethanol to the solution slowly over a period of one minute ** 8. Allow 3-4 days for the liquid to evaporate 9. Using a scoopula, carefully scrape off the crystals that have formed in the beaker 10. Place one piece of the crystalline structure under a microscope to observe
 * 2. Add 20mL of distilled water to the beaker **
 * 3. Stir the solution to dissolve the solid

=Observations=

Upon mixing the Copper (II) Sulfate Pentahydrate with the distilled water, the mixture turned to a clear, light-blue liquid (Fig. 2) and solid Copper (II) Sulfate crystals were observed at the bottom of the solution. After the crystals had completely dissolved, the mixture was observed to be a cloudly translucent mixture. This issue was solved by filtering the solution to remove the excess crystals, impurities and turbidity, thus leaving a clear, light-blue solution.

Upon mixing the Copper (II) Sulfate with the Ammonia, a dark, opaque, navy-blue liquid was observed (Fig. 3). As the substance was left in the open to allow all the water to evaporate, small dark-blue crystals began accumulating on the side of the beaker while a thick, viscus substance began forming on the bottom of the beaker which eventually became solid over time.

When the water had completely evaporated, the crystals observed at the bottom of the beaker were of different sizes, although none were larger than a few millimeters in size.

 After examining the crystals under a microscope by using a lamp and a flashlight, it was observed that crystals did indeed form. The crystals were a strong blue colour and have flat edges similar to that of polished gemstones under low and medium magnification under a microscope and a stochioscope. Since the each crystal of the tetraaminecuprate (II) complex have similar angles, it was thus proven that a compound has only one possible lattice structure as unit cells can form different molecules. Under low magnification small crystals were visible. Under medium magnification the crystal structure was very easily seen but under high magnification it was impossible to focus the microscope so that the crystals were visible.

=Calculations=

__Solubility according to the solubility chart found on page 635 of Chemistry 11:__

Solubility of Reactants:
Copper (II) Sulphate Pentahydrate is soluble, as sulphates and hydrates are usually soluble.

Solubility of Products: A precipitate is formed because of the hydroxide present, however the precipitate soon dissolves to become the Tetraammine. Tetraamminecuprate (II) complex is soluble. .......................... __The two chemical reactions that took place in the experiment can be summarized using the following equations:__

CuSO4(aq)+2NH4OH --> Cu(OH)2+(NH4)2SO4 Cu^2+ + 4NH3(aq) --> [Cu(NH3)4]^2+ ..........................

Calculation of the limiting reagent in the first stage of reaction
The following equation is used: //**n = M/Mm**, where n is the number of moles, Mm is the molar mass, and M is the mass in grams.//

The number of moles of CuSO4 present is calculated as follows:

Mm(CuSO4) = 159.62gmol-1 Mm(CuSO4·5H2O) = 249.62gmol-1 M(CuSO4·5H2O) = 10g

% composition(CuSO4) = [M(CuSO4) / M(CuSO4·5H2O)] * 100% % composition(CuSO4) = [159.62gmol-1 / 249.62gmol-1] * 100% % composition(CuSO4) = 63.95 %

M(CuSO4) = % composition(CuSO4) * M(CuSO4·5H2O) M(CuSO4) = 0.6395 * 10g M(CuSO4) = 6.39g

n(CuSO4) = M(CuSO4) / Mm(CuSO4) n(CuSO4) = 6.39g / 159.62gmol-1 n(CuSO4) = 0.04mol //Therefore, 0.04 moles of CuSO4 are present.//

Mole ratio CuSO4 : 2NH4OH = 1:2 Moles of NH4OH needed = n(CuSO4) * 2 Moles of NH4OH needed = 0.04mol * 2 Moles of NH4OH needed = 0.08 mol //Therefore, 0.08 moles of NH4OH are needed//

14M of NH4OH is present (see materials above) In order to convert this quantity into moles, the follow calculations are completed using the following equation:


 * //C = n/V//**//, where C is the concentration, n is the number of moles, and v is the volume.//

C(NH4OH) = 14M = 14mol/L V(NH4OH) = 20mL = 0.02L

n(NH4OH) = C(NH4OH) * V(NH4OH) n(NH4OH )= 14mol/L * 0.02L n(NH4OH) = 0.28mol //Since Mole ratio CuSO4 : 2NH4OH = 1:2,// n(NH4OH) needed is: n(NH4OH) = 0.28mol / 2 n(NH4OH) = 0.14 moles //Therefore, 0.14 moles of NH4OH is present.//

0.14mol > 0.08mol. //There is more NH4OH present than is needed for the chemical reaction. Therefore, CuSO4 is the limiting reagent and NH4OH is the excess reagent.//

CuSO4(aq) is the desired limiting reagent as it will later be dissolved into the NH3 in the second stage of the reaction.

n(CuSO4) = 0.04 mol (from above) Mm(Cu(OH)2) = 97.57 gmol-1 M(Cu(OH)2) = n(CuSO4) x Mm(Cu(OH)2) M(Cu(OH)2) = 0.04 mol x 97.57gmol-1 M(Cu(OH)2) = 3.9028 g

M(Cu^2+) = 0.04 mol x 63.55gmol-1 M(Cu^2+) = 2.542 g

Mm((NH4)2SO4) = 131.13 gmol-1 M((NH4)2SO4) = n (CuSO4) x Mm of (NH4)2SO4 M((NH4)2SO4) = 0.04 mol x 131.13 gmol-1 M((NH4)2SO4) = 5.2452 g

Excess NH4OH = n(NH4OH) present - n(NH4OH) needed) x Mm(NH4OH) Excess NH4OH = (0.28 mol - 0.08 mol) x Mm(NH4OH) Excess NH4OH = 0.2 mol x 35.06 gmol-1 Excess NH4OH = 7.012 g //There are 7.012 g of excess NH4OH that will used in the second stage of the reaction.//

Calculation of the Limiting Reagent in second stage of the reaction
M(NH3) = % composition of NH3 x excess of NH4OH % composition of NH3 = 17.04/35.06 x 100% % composition of NH3 = 48.6% M(NH3) = 3.408 g

n(NH3) = M(NH3)/Mm(NH3) n(NH3)= 3.408g/17.04gmol-1 n(NH3) = 0.02 moles

M(Cu^2+) needed = n(NH3) x Mm(Cu(OH)2) M(Cu^2+) needed = 0.02 mol x 63.55gmol-1 M(Cu^2+) needed = 1.271g

2.542g> 1.271g //There is more Cu^2+ present has less than needed for the chemical reaction. Therefore, Cu^2+ is the limiting reagent and NH3 is the excess reagent.//

M[Cu(NH3)4]^2+ = n(NH3) x Mm[Cu(NH3)4]^2+ M[Cu(NH3)4]^2+ = 0.02 mol x 131.71gmol-1 M[Cu(NH3)4]^2+ = 2.6342g

//There are 2.6342g of the Tetramminecuprate (II) Complex in the final solution.//

=Conclusion= In this experiment, a coordination compound was successfully synthesized and crystallized. Ammonia, a ligand, was added to Copper II Sulphate Pentahydrate to form a Tetrammine Cuprate complex, a complex compound. Ethanol was added during the experiment to ensure the rapid evaporation of the solution so that the crystal formation could be studied (Fig. 4). As the molecular structure of Tetrammine Cuprate II Hydroxide is square planar and that the physical crystal structure of the is an elongated shape, it is related to the orthorhombic crystal formation, the crystals observed confirmed the hypothesis stated in the introduction in that they were prisms.

=Discussion (Analysis)=

Before proceeding with the experiment, it was known that a coordinate bond is located in the Ammonia (aqueous) solution. This coordinate bond would lead to the possible and now proven complex compound. When the Ammonia was diluted with water to become Ammonium Hydroxide. The NH4+ has a coordinate bond, when a Hydrogen cation ( a Lewis Base) bonded with the NH3 (a Lewis acid). Ammonia has one lone pair of electrons and it donated its electron to the Hydrogen cation to form a coordinate bond. The coordinate bond is what made NH4+ a ligand, and thus producing a complex compound.

The experiment was recognized to have two reactions. The complex compound was made when Ammonia (aqueous) is added to the Copper II Sulphate, a double displacement reaction occurs with the Hydroxide from the Ammonia (aqueous) bonding with the Copper to form a blue precipitate. The products from the first reaction are Copper II Hydroxide and Ammonium Sulphate. From the calculations it was seen that in the first reaction, Copper II Sulphate was the limiting reagent. There was an excess of Ammonia (aqueous).The excess of Ammonia (aqueous) from the first reaction then reacts with the Copper II Hydroxide of the products of the first reaction to form a coordinate compound, Tetrammine Cuprate II Hydroxide as well as the Sulphate, which is a spectator ion.

When performing the experiment, there were colour changes during the reactions. The colour that is displayed in a coordinate compound is a result of the orbitals in which the electrons occupy, specifically the changes in energy of the d orbitals in the case of a transition metal complex compound because transition metals have partially filled d orbitals.

When a coordinate compound is formed, the degenerate orbits are split into higher and lower energy groups. Because the colour of a substance is a result of the photons emitted, which is in turn affected by the difference in the energy levels, the colour a coordinate compound displays is determined by the size of the energy gap between two levels. This distance varies with the nature of the transition metal ion, its oxidation number, and the nature of the ligands in the coordinate compound. When light is shined through the complex compound, part of the energy is used to promote an electron from the lower orbital to a higher one. If the energy absorbed corresponds to the specific amount of energy required to promote the electron to a higher orbital, the colour observed will be the complementary colour to that specific wavelength. For example, if the wavelength absorbed corresponds to the colour yellow, then the coordinate compound’s colour will be a dark blue.

As for complex compounds that are composed of non transition metals, some will turn out to be colourless. The reason for this is because they do not have partially filled d orbitals. When this is the case, the energy absorbed by the electrons has to be the right amount for the electron to be promoted to a higher level. In the case of zinc, for example, since the 3d orbital is completely filled, zink complex compounds will turn out to be colourless.

Since the above is a major simplification of the actual process, one should take note that many other, more complex processes are also at play. Specifically, the determination of a colour of a coordinate compound can be summarized as follows:

- the nature of the metal ion (the number of electrons in the occupying the d orbital) - the arrangement of the ligands around the metal ion (for example geometric isomers can display different colours) - the nature of the ligands surrounding the metal ion. (the stronger the ligands, the greater the energy difference between the split high and low 3d groups)

After leaving the product for three days, the substance was broken off into pieces and were examined. The crystals were then carefully examined under a compound microscope using 10X magnificatoin (Fig. 5) 40X magnification (Fig. 6). The experiment was done with varying results in crystal structure due to differences in procedure. The first two examples produced larger crystals than the third specimen. The only difference in procedure was that the compound was filtered before adding ethanol to remove impurities. Therefore the first two experiments had impurities in the solution that was formed. It can be hypothesized that the impurities in the third solution aided in the crystallization of the Tetraamminecupric (II) complex as it provided a starting point for the crystals to grow. From the observations, it was seen that the crystals formed were rectangular prisms with a square base and peaked at the top. Since the each crystal of the Tetraaminecuprate(II) complex have the same angles it is proven that a compound has only one possible lattice structure as unit cells can form different molecules.The molecular structure of the Tetrammine Cuprate II Hydroxide is square planar as there are four ligands connected to the central metal. The Tetrammine Cuprate II Hydroxide contains the transition metal Copper, which is a group 11B transition metal and therefore has a FCC crystal structure.

Modifications Made to the Experiment
The experiment procedure was completed for a total of three times, with each version of the experiment incorporating various modifications in order for the outcome to be as accurate as possible. As with all scientific experiments, in order for an outcome to be as accurate as possible, multiple attempts of an experiment must be made. This was applied to this experiment by conducting it a total of three times.

During the first experiment attempt, the Copper (II) Sulphate Pentahydrate solution was not filtered. Because of this, the mixture was observed to be turbid. This was modified in the second and third attempts by using filter paper to remove any impurities in the solution.

Modifications to the amount of water used to dissolve the Copper (II) Sulphate Pentahydrate were also made. This was in response to the observation that the Copper (II) Sulphate Pentahydrate was not dissolving completely when added to the water. There were fears that the mixture was too saturated so the particles could not dissolve. The amount of water used was therefore changed from 10mL to 20mL.

The amount of ethanol used was also changed. It was realized that ethanol was a catalyst (did not take part in the actual chemical reaction, but rather helps to speed up the evaporation process of the final solution to yield crystals). Because of this, the amount of ethanol used was decreased from 20mL to 10 mL.

In order to not disturb the formation of crystals, the ethanol that was added to the mixture during the second and third experiment attempts were added gradually and excessive turbulence to the beaker with the solution was avoided.

=Sources of Experimental Error=  As with all scientific experiments a certain degree of experimental error exists which would have yielded inaccurate results. Due to the length and number of steps required in completing this experiment, the amount of sources of experimental error was particularly numerous. This section will attempt to identify those sources of error and explain in detail what type of effect each one would have had on the outcome of the experiment. While identifying the sources of error in the experiment, it becomes necessary to distinguish them into two broad categories: unavoidable (systematic) errors, and avoidable (human) errors:

Unavoidable (systematic) Errors
The first source of systematic error was a systematic (unavoidable) one, and was attributed to the balance that weighed the Copper (II) Sulphate Pentahydrate. The electric balance used was a standard laborary balance, and was neither calibrated nor was it standardized, an error would have existed in the reading on the balance. This source of error was evident when the reading was observed to fluctuate as much as one decimal place even when nothing was placed on top. When the Copper (II) Sulphate Pentahydrate was placed on the balance, the reading would not show an accurate number, but would rather display a constantly changing number ranging from ** 9.97 grams to 10.02 grams. The amount of **Copper (II) Sulphate Pentahydrate used could not be calculated accurately. The difference in the amount would have yielded an inaccurate result in the end.

A second source of experimental error was in the amount of ammonia that was added to the Copper (II) Sulphate Pentahydrate. The amount of ammonia that was required for the experiment as explained in the procedure was 10mL. In order to obtain as close an amount to 10mL as possible, the ammonia was first put in a graduated cylinder. However, do to the presence of the meniscus, it was difficult to measure an accurate amount. The variation in the amount of ammonia added to the Copper (II) Sulphate Pentahydrate solution would have affected the outcome of the experiment.

Avoidable (Human) Errors
A third source of error, and one that was avoidable, was due to the fact that the substance in the beaker was left out on the laboratory counter to crystallize overnight uncovered. As a result, a variety of contaminants would have made their way into the beaker. These contaminants would have reacted with the substance in the beaker, thus making it no longer pure. The presence of an impure substance would have undoubtedly yielded inaccurate results. Another source of contamination came from the lab equipment used for the experiment: although the beaker was cleaned using water and paper towel prior to the experiment, residues of a previous chemical from another experiment would still have been present due to the lack of a thorough cleaning process. The scoopula that was used to transfer the Copper (II) Sulphate Pentahydrate was also placed on the laboratory counter prior to the experiment, which would have attracted more external contaminants.