Since the beginning of chemistry, coordination compounds have been known (mostly through accidental discoveries). An example of such compounds includes Prussian blue, an accidental discovery 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 had, however, emigrated to America before he could do much more. In the following year, an American scientist, Oliver Wolcott Gibbs started an investigation and teamed up with Genth in 1856 to publish a paper in which the described thirty five cobalt-ammine complexes.

The term coordinate compounds, or simply 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, 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 orbital. Coordinate compounds differ from ionic and covalent chemical bonds in that both the electrons in a bond come from the same atom. However, a coordinate bond is indistinguishable from covalent bonds after chemical bonding occurs.

The idea of 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. 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.
An example of coordinate bonding: The Lewis base and the Lewis acid react to form a compound. The nitrogen donates its lone pair to the Lewis acid, which is an example of a coordinate bond.
An example of coordinate bonding: The Lewis base and the Lewis acid react to form a compound. The nitrogen donates its lone pair to the Lewis acid, which is an example of a coordinate bond.


Copper II Sulfate Pentahydrate
Copper II Sulfate Pentahydrate

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 and has a very strong and sharp odour. According to the EU (European Union) classification, ammonia is hazardous, caustic, and corrosive. Thus, when the 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 is to be handled with caution.


When using the Bunsen burner, safety goggles must be worn. Gloves and aprons are highly recommended. Beakers, test tubes, and flask tongs are used when appropriate. When using the hot plate, place all containers on top of the wired gauze to avoid spilling the substance on the hot plate.


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.


1. Place 10 grams of copper(II) sulfate pentahydrate into a 250mL beaker
2. Add 20mL of water to the beaker
3. Stir the solution to dissolve the solid
4. If the solution is too turbid, use filter paper to filter the copper (II) sulphate
5. Add 20mL of 15 M 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. Use a scoopula, carefully scrape of the crystalline structure of the compound formed.
10.Put one piece of the crystalline structure under the microscope for observation.


Upon mixing the Copper (II) Sulfate Pentahydrate with the distilled water, the mixture turned to a clear, light blue liquid and solid Copper (II) Sulfate crystals were observed at the bottom of the solution. Stirring the mixture using a stirring rod encouraged the crystals to dissolve. After the crystals had completely dissolved, the mixture was expected to turn a clear blue liquid. In reality, though, it was observed to be a cloudly translucent mixture due to the turbidity of the distilled water. 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 formed. 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, solid substance began forming on the bottom of the beaker.

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.

The crystals were observed to be opaque, thus a flashlight was needed in order to view them under a microscope by shining it on the crystals from above. A stochioscope is used for viewing the 3D structure of the crystal, which were observed to be ANGELA FILL THIS IN

After examining the crystals under a microscope 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 the same angles it is proven that a compound has only one possible lattice structure as unit cells can form different molecules.



The two main chemical reactions that took place during the experiment could be summarized as follows:

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

Calculation of the Limiting Reagent in the first stage of reaction:

Moles of CuSO4 present is calculated as follows:
n of CuSO4 = M of CuSO4/Mm of CuSO4
Mass of CuSO4 = % composition of CuSO4 in CuSO4 in CuSO4.5H2O X Mass of CuSO4.5H2O
% composition of CuSO4 = Mm of CuSO4/Mass of CuSO4 X 100%
% composition of CuSO4 = 159.62/249.62 X 100%
% composition of CuSO4 = 63.95 %
Mass of CuSO4 = 63.95 % X 10g
Mass of CuSO4 = 6.4g
n of CuSO4 = 6.4g/159.62gmol-1
n of CuSO4 = 0.04 moles

Mole ratio: CuSO4 : NH4OH = 1:2
Moles of NH4OH needed = n of CuSO4 x 2
Moles of NH4OH needed = 0.08 moles

Moles of NH4OH is 14M of NH4OH present is calculated as follows:
n of NH4OH = C X v
n of NH4OH = 14M x 0.02 L
n of NH4OH = 0.28 moles/2
n of NH4OH = 0.14 moles

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

Mass of Cu(OH)2 = n of CuSO4 x Mm of Cu(OH)2
Mass of Cu(OH)2 = 0.08 mol x 97.57gmol-1
Mass of Cu(OH)2 = 7.8056 g

Mass of (NH4)2SO4 = n of CuSO4 x Mm of (NH4)2SO4
Mass of (NH4)2SO4 = 0.08 mol x 131.13 gmol-1
Mass of (NH4)2SO4 = 10.4904 g

Excess NH4OH = (0.28 - 0.08) x Mm of NH4OH
Excess NH4OH = 0.2 x 35.06
Excess NH4OH = 7.012 g

Calculation of the Limiting Reagent in second stage of the reaction:
Mass of NH3 = % composition of NH3 x excess of NH4OH
% composition of NH3 = 17.04/35.06 x 100%
% composition of NH3 = 48.6%
Mass of NH3 = 3.408 g
n of NH3 = Mass of NH3/Mm of NH3
n of NH3 = 3.408g/17.04gmol-1
n of NH3 = 0.02 moles x 3
n of NH3 = 0.06 moles

Mass of Cu(OH)2 needed = n of NH3 x Mm of Cu(OH)2
Mass of Cu(OH)2 needed = 0.06 mol x 97.57gmol-1
Mass of Cu(OH)2 needed = 5.8542 g

The NH3 is the limiting reagent.

Mass of [Cu(NH3)4](OH)2 = n of NH3 x Mm of [Cu(NH3)4)(OH)2
Mass of [Cu(NH3)4](OH)2 = 0.06 mol x 165.73 gmol-1
Mass of [Cu(NH3)4](OH)2 = 9.9438g

The colour change and formation of precipitate is present due to the following statements:
Transition metals forms different colours and complexes because of their structure. These colours form when change in energy distribution of photons gets either absorbed or reflected after being in contact with another substance.

Solubility of reactants:
Copper (II) Sulfate Pentahydrate is soluble. As sulphate and hydrates are usually soluble.
Ammonia is soluble.

When Copper(II) Sulfate Pentahydrate reacted with NH3 with water, a double displacement occurs, the Copper displaces and reacts with Hydroxide to form a blue precipitate. The blue ppt of Cu(OH)2 will, then, dissolve to form a deep blue solution called tetraamminecopper (II) ions or just simply call it a soluble complex ion.

Solubility of Products
A precipitate is formed because Copper hydroxide is present. Hydroxides are insoluble unless reacted with Group 1 elements and some coordinate compounds.
Tetraamminecopper (II) Hydroxide is soluble as a co-ordinate compound is present (NH4+).

Structure of Tetraamminecuprate (II) complex
Structure of Tetraamminecuprate (II) complex

The Tetraamminecuprate(II) complex is of square planer structure.

After examining the structure of the complex crystals....

*calculate yield*


Crystals formed when
Copper (II) Sulfate was mixed with the Ammonia. After the solution was allowed to evaporate, crystals were formed, therefore proving that the Copper (II) Sulfate had formed coordinate bonds with the Ammonia. This is true because

*calculate error*

Helen: I found something interesting about the colours that a transition metal forms with another substance, I don't know whether or not you guys would like it but it is very well related with coordination compounds/complexes. It is probably good to include it under the discussion so here's the following.
"Coordination by ligands can play a part in determining color in a transition compound, due to changes in energy of the d orbitals. orbitals remove degeneracy of the orbitals and split them in to higher and lower energy groups. The energy gap between the lower and higher energy orbitals will determine the color of light that is absorbed, as electromagnetic radiation is only absorbed if it has energy corresponding to that gap. When a ligated ion absorbs light, some of the electrons are promoted to a higher energy orbital. Since different frequency light is absorbed, different colors are observed.
The colour of a complex depends on:
  • the nature of the metal ion, specifically the number of electrons in the d orbitals
  • 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 then the greater the energy difference between the split high and low 3d groups.
The complex ion formed by the d block element zinc (though not strictly a transition element) is colorless, because the 3d orbitals are full - no electrons are able to move up to the higher group."

Ammonia can act as a ligand in transition metal complexes. It is a pure σ-donor. A complex needs ligands to be coloured. Without ligands, all five d orbitals are equal in energy or degenerate. When ligands are present, some of the d orbitals become higher in energy than before, and some become lower. This happens because some of the d orbitals are nearer the ligands than others so they experience more repulsion from the ligand electrons and thus have a higher energy.
Ligands have lone pairs of electrons that they use to bond to transition metal atoms or ions. The presence of these ligand electrons repels the transition metal's electrons in its d orbitals. The metal d electrons can't go anywhere to avoid the repulsion, so the only effect is that their energy increases. When their energy increase they produce different colours.

As for the crystal structure sigh... well we might have to examine the crystal structure but i don`t think we`ll find anything.

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.