The aim of the experiment was to determine the empirical formula of hydrated copper (II) sulphate(CuSO4· xH2O)by experiment and to investigate the changes of copper (II) ions in solution.
Copper is a d-block Transition metals, which are elements in Group 3-12 of the Periodic Table. It has the electronic structure 1s22s22p63s23p63d104s1, and can form complex formation (Chemguide, 2003). Complex ions are compounds having a central atom surrounded by other molecules called ligands, and these ligands can form dative covalent bonds to the central particle (Lister and Renshaw 2000). Lister and Renshaw (2000) further state that good ligands can displace poorer ligands from complexes, and the log of stability constants (logKc) can be used to measure the stability of complexes. The larger logKc, the more stable the complex (Lister and Renshaw, 2000). Most complexes are coloured since the movement of electrons between d orbitals will absorb a quanta of electromagnetic energy and the resulting energy gap is corresponding to frequencies of electromagnetic radiation in the visible region of the spectrum (Lister and Renshaw, 2000).
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Copper (II) sulphate has a considerable number of compounds, which have different degrees of hydration. Fishing (2009) points out that pentahydrate (CuSO4· 5H2O) is bright blue due to the water of hydration and when heated, the blue copper sulphate can be dehydrated to a grey-white power called anhydrous salt. The structure can be seen in Figure 1. Two water molecules will be lost at 30? when heated, and two more are then lost at 110?, followed by the last one at 250? (Fishing, 2009).
As stated by Chemguide (2003), the typical blue hexaaquacopper (II) ion- [Cu(H2O)6]2+ is the simplest form in solution. Forming stronger bonds than water molecules with the center particle (Cu2+), chloride ions can displace water molecules and form a yellow-green [CuCl4]2- (aq) whose value of logKc is 5.6 (Lister and Renshaw, 2000). Figure 2 shows the structure of the [CuCl4]2-. The reaction of hexaaquacopper (II) ions with chloride ions can be shown as:
[Cu (H2O)6]2+(aq) + 4Cl-(aq) [CuCl4]2-(aq) + 6H2O(l) (Chemguide, 2003)
When copper (II) sulphate solution reacts with ammonia, it has two separate stages. Chemguide (2003) suggests that in the first stage, a small amount of ammonia can lead to hydrogen ions being removed from the hexaaqua ion. As a result, a neutral complex is produced, which is a precipitate with a pale blue colour (Chemguide, 2003). The equation for this reaction can be written as:
[Cu(H2O)6]2+(aq) + 2NH3(aq) ?[Cu(H2O)4(OH)2](aq) + 2NH4+(aq) (Chemguide, 2003)
Chemguide (2003) further reports that when adding excess ammonia solution, the ammonia will replace four of the six water molecules from [Cu(H2O)6]2+, forming a deep blue [Cu(NH3)4(H2O)2]2+(aq). Its value of logKc is 13.2 (Lister and Renshaw, 2000). The reaction can be shown as:
[Cu(H2O)6]2+(aq) + 4NH3(aq)?[Cu(NH3)4(H2O)2]2+(aq) + 4H2O(l) (Chemguide, 2003)
The apparatus consisted of crucible, spatula, burner and tongs, electronic balance, desiccator and stand, as well as three conical flasks, and the chemicals included copper sulphate, concentrated hydrochloric acid and ammonia solution. The method of this experiment was divided into two parts.
According to Lane (2009.a), the practical for determining the formula in part 1 was done as following steps, and all figures gained were corrected to two decimal places. To start with, the inside of the crucible was cleaned with a cloth. Subsequently, a paper clip was placed in the crucible. The crucible was then weighed and the weight was recorded. After that, using the electronic balance, 2.58g copper sulphate was added into the crucible. Before placing the crucible on the stand and heating, the burner was lit and placed under the stand. Using the paper clip, the crystals were stirred when heating. The change in colour was then noted. After that, using the tongs, the crucible was placed inside the desiccator to cool down. When cool down to room temperature, the crucible was reweighed and the procedures heating, cooling, and weighing were repeated until constant weight was recorded. Some water was added to the crucible at the end and the result was noted.
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In part 2, the steps making complex ions in solution can be shown below (Lane, 2009.b). At the beginning, some copper sulphate and water were put into three conical flasks and these flasks were shaken and observed. After that, using a pipette, concentrated hydrochloric acid was dropped into one flask. Any phenomena were noted. Before observing, a small amount of ammonia solution was then added into the second flask. At the end, excessive ammonia solution was put into the third flask and the result was recorded.
In part 1, using the data from Table 1 and the formula n = m/Mr, where n = moles, m = Mass of sample and Mr = relative Molecular Mass (Lane, 2009.b), the steps in the calculation of the formula (CuSO4· x H2O) can be shown as follows.
Mr (CuSO4) =159.6
m (CuSO4) =1.64g
n (CuSO4) = m/Mr =1.64/159.6 = 0.010289 moles
Mr (H2O) = 18
m (H2O) = 0.94g
n (x H2O) = m/Mr = 0.94/18 = 0.052222 moles
The ratio can be determined: CuSO4: x H2O = n (CuSO4): n (x H2O)
Therefore: 1 : x = 0.010289 : 0.052222 x = 5.0755
The value x = 5.0755 can be approximate to x = 5; therefore, the empirical formula is CuSO4· 5H2O.
As stated by theory, when heating the crystals, the water will evaporate depending on the temperature. Repeating heating and stirring the crystals can ensure the hydrated CuSO4 is dehydrated completely. The positive result of x = 5 indicates that the heater can reach 250? at least in the experiment, which provides the power to drive off all water molecules from the crystal.
Copper (II) oxide (CuO) and sulfur trioxide (SO3) will be produced when heating the crystal at around 600? (Fishing, 2009). It can be deduced that when over heating, the black CuO (s) and the pungent smell SO3 (g) would be observed.
The colour changing grey-white to blue when adding water into anhydrous copper sulphate can explain why the crucible needs to cool down inside the desiccator. It is to avoid the anhydrous copper sulphate absorbing H2O from the air and reforming hydrated CuSO4. This is also why anhydrous CuSO4 can be widely used for testing the presence of water in other chemical experiments.
The green colour in part 2 could be explained in terms of a mixture of colours from [Cu(H2O)6]2+ (blue) and [CuCl4]2- (yellow-green). It can be deduced that if adding enough water into the green solution, the green solution would turn back to blue, because a high concentration of H2O would lead to the reversible reaction tending to produce more [Cu(H2O)6]2+. The results about copper (II) ions reacting with NH3 (aq) mean that the blue precipitate (Cu(OH)2) can dissolve when adding excessive of ammonia.
All reactions tend to high stability, low energy. As mentioned by theory, the logKc value of [Cu(NH3)4 (H2O)2]2+ (aq) (13.2) is larger than [CuCl4]2-(aq) (5.6), which means that the complex [Cu(NH3)4 (H2O)2]2+ is more stable than [CuCl4]2-, and NH3 is a better ligant than Cl-. Therefore, the reaction NH3 replacing Cl- from [CuCl4]2-(aq) can be deduced.
Compared with the empirical formula CuSO4· 5H2O, the calculated answer is slightly high, although it can be estimated to x = 5. There are three main reasons can explain the result. The crucible may not be completely dry, and extra water evaporated will give a higher value. In addition, when over heating, the mass of gases would be regarded as the loss mass of water. Finally, some crystals would splash out when stirring, which can lead to the calculated value higher.
In order to produce more accurate result, several areas could be improved. Controlling the flame intensity of burner, putting an asbestos net under the crucible or granulating the crystal can reduce the possibility of decomposition and ensure the crystal is dehydrated completely.
In conclusion, the empirical formula of hydrated copper sulphate can be determined as CuSO4· 5H2O by experiment. Water molecules can make the copper complex ion blue. Ammonia (NH3) causes deep blue colour, and chloride ions (Cl-) make the copper complex solution yellow-green. Therefore, the nature of ligands can affect the energies of the d orbitals and produce complexes with different colours.
- Chemguide (2003) [online] Copper Available at: http://www.chemguide.co.uk/inorganic/transition/copper.html [Access at: December 20, 2009]
- Fishing (2009) [online] Copper Sulfate, Equation for Decomposition Available at: http://www.finishing.com/116/07.shtml [Access at: December 22, 2009]
- Kecheng (nd) [online] Indentify the Formula of Hydrated Copper Sulphate Available at: http://kecheng.edu.people.com.cn/index/newscontent/snsy/czhx/syzl6_2_4_2.htm [Access at: December 28, 2009]
- Lane, R (2009.a) Chemistry Practical Handout
- Lane, R (2009.b) IFY Course Notes
- Lister, T. and Renshaw, J. (2000). Chemistry for Advanced Level (third edition). Cheltenham: Nelson Thornes Ltd.
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