Purpose:
The purpose of this experiment is to investigate some properties of ionic composite compounds containing H2O molecules, specifically experiments on blue hydrated copper (II) sulfate. The number of coordinated H2O molecules will be determined.
Introduction:
The calculation of the value of “x” (the number of coordinated H2O molecules) is based on the “relative molecular mass (Mr)” and “the mole.” Lister and Renshaw (2000) stated that Mr is the mass of a molecule compared to the mass of one H atom and is the sum of the relative atomic masses. The number of moles is equal to (mass in g) / Mr. The relative atomic masses of Cu, H, O, and S are 63.5, 1, 16, and 32, respectively, and CuSO4 is 160 (Lister and Renshaw, 2000).
Copper (Cu) is a d-block element in the periodic table and is a member of transition elements. Lister and Renshaw (2000) pointed out that transition elements, including Cu, have several specific properties, including being usually colored compared to s-block metals and the ability to form complexes with dative bonds. In Cu’s complex, other molecules such as NH3, which have lone pairs, can form dative bonds with Cu, and these molecules are called ligands. Such ligands include H2O, NH3, and Cl- (Lister and Renshaw, 2000).
According to Clark (2000), a dative covalent bond, also called a coordinate bond, is a covalent bond (sharing a pair of electrons) where both electrons are supplied by the same atom. Normally, a lone pair (unshared pair of electrons) is recognized as a cation such as Cu(II) to help obtain a full outer shell (Lister and Renshaw, 2000). A typical example is NH3, which has a lone pair to form a dative bond.
Hydrated salts are compounds containing water molecules (Farlex Inc., 2009). When a crystal of the substance forms, some waters are incorporated. They will be driven off when the crystal is heated, and the substance becomes an anhydrous salt. A typical example is Cu(II) sulfate. According to Bennett (1998), hydrated Cu(II) sulfate has four water molecules directly bound to the Cu (by dative bonds) and forms a Cu(H2O)42+ ion first. In that ion, the Cu is at the center of a square surrounded by the O’s of the water molecules. One sulfate ion and one water molecule are bonded to each other by an H bond and connect with Cu(H2O)42+ as a line. Hydrated Cu(II) sulfate becomes anhydrous with the following expression:
CuSO4•xH2O(s) → CuSO4(s) + xH2O(l)
The important nature of transition elements (T.E.), including Cu, is that they are colored. Lister and Renshaw (2000) stated that this characteristic is caused by the energy spread between two energy degrees in the vitamin D orbital, which all T.E. have, and the energy spread needs to absorb light energy to be filled. An equation that relates energy spread (E) to frequency (v) is E=hv, where H is a constant, and if the frequency of the substance falls in the visible part of the spectrum, the remainder of visible radiations (not absorbed) will appear as the substance’s color. It is the ligands that make the difference in energy degrees. According to Clark, J (2000), when ligands such as H2O, NH3, and Cl- approach the ions of T.E., there is a repulsive force between the electrons from ligands and the 500 orbital of T.E., resulting in the splitting of the energy into two groups, one of which has been promoted to a higher energy degree, creating a spread.
Method
The source of this method is Lane, R (2009).
The following chemicals were provided: Copper Sulphate (s), concentrated hydrochloric acid (L), and ammonia solution (L).
The following setups were provided: spatula, tongs, paper cartridge holder, electric balance, desiccator, crucible, burner, base, pipeclay triangle, and conical flask.
Part A:
To record information, a crucible was cleaned with a tissue. After that, a paper cartridge holder was placed in the crucible, and both were weighed using an electric balance and recorded to 0.01g. Then, 2-3g of Cu sulfate was added to the crucible with a paper cartridge holder and weighed.
After that, hydrated Cu (II) sulfate was heated over an alcohol burner placed under a base. The crucible was placed on the base and heated for 5 minutes. The crystal was stirred and broken up using the paper cartridge holder and observed. Next, using tongs, the crucible was placed inside a drying desiccator for 5 minutes to cool down (the paper cartridge holder remained in the dish). When it was cool enough to touch, it was reweighed. The process was repeated twice until a constant weight was achieved.
Part B:
Copper (II) sulfate was made into a solution and reacted with other solutions. Copper (II) sulfate and water were put into 3 conical flasks and shaken to dissolve. Then, using a pipette, concentrated hydrochloric acid was dropped into one flask and observed. The process was repeated in flask 2 with ammonia solution replacing hydrochloric acid. Two steps were performed, first to add a small amount of ammonium hydroxide and second to add additional solution.
Discussion
The calculation of the value of x:
Number of moles of anhydrous Cu (II) sulfate: M1 (g)/M1 (g/mol) = 1.35g/(63.5+32+16•4) g·mol-1 = 8.46 x 10-3 mol
The number of moles of combined H2O: m2/M2= 0.78g/18g·mol-1 = 0.043 mol
Therefore, the value of x = 0.043 mol/8.46 x 10-3 mol = 5.12 ≈ 5 moles
The empirical formula of hydrated Cu sulfate is CuSO4·5H2O.
From the computation of the consequences, it appears that there are about five H2O molecules surrounding each Cu(II) sulfate. As the coordinated H2O was lost bit by bit, it is supposed that the bonds between H2O molecules and Cu ion broke one by one. The H bond may break first as its bond energy is low (Lister and Renshaw, 2000), and in turn, the dative bonds break. The fewer bonds combined with Cu may result in higher energy required to break them. The further expression can be written as follows:
CuSO4•5H2O (s) → CuSO4•H2O (s) → CuSO4 (s)
so CuSO4•5H2O (s) → CuSO4 (s) + 5H2O (l)
The way that Cu(II) exists in the water is [Cu(H2O)4]2+, and it is formed by dative bonding (Lister and Renshaw, 2000). Here, H2O molecules whose O has lone pairs are attracted to copper(II) to fill its empty orbital. When hydrochloric acid was added to copper(II) sulfate, in this case, Cl- is a ligand that forms stronger bonds than H2O molecules. In other words, H2O slightly splits the energy level, and Cl- can produce a large energy spread. As a result, the anion replaces H2O (Lister and Renshaw, 2000). The reversible equation can be written below:
[Cu(H2O)4]2+ (aq) + 4Cl- (aq) → [CuCl4]2- (aq) + 4H2O (l)
Here [CuCl4]2- is yellow for several reasons. According to Clark, J (2000), Cl- had split the Cu’s energy level somewhat related to H2O and procured a smaller energy spread that determines the wavelength of light being absorbed. Thus, the wavelength of the substance is higher, and the lower energy (dark color) light was absorbed. Consequently, in the spectrum, the lighter color, i.e., yellow, appears (from magenta to red as the wavelength increases). As the reaction happened gradually and is reversible, the color of the solution changed slowly from blue to green (the mixed color of blue and yellow).
When ammonia solution was added to copper sulfate solution, it is again a replacement. Cu(II) first reacts with OH- from ammonia water to form Cu(OH)2 so that a small amount of blue suspension was produced. Due to the small difference between OH- and H2O in splitting, the color change was small. Then the Cu(OH)2 reacted with ammonia solution. Similar to the reaction with hydrochloric acid, ammonium hydroxide molecules replaced H2O, and the compound produced is blue and soluble in water. This process proceeds to darker color because ammonium hydroxide makes a large energy spread. As a consequence, lower wavelengths of light were absorbed, resulting in a darker color, i.e., blue. The two reaction equations are:
Cu2+ (aq) + 2NH4 +OH- (aq) → Cu(OH)2 (s) + 2NH4+ (aq)
Then a reversible equation
Cu(OH)2 (s) + 2NH4 +OH- (aq) → [Cu(NH3)4]2+ (aq) + 4H2O (l)
There were some mistakes found during this experiment. While dehydrating the Cu(II) sulfate in the 3rd clip, the desiccator was not covered, resulting in the moisture mixing with anhydrous Cu(II) sulfate, so that the final record was greater than expected. To improve, the whole experiment should be conducted in a highly dry condition in order to get rid of water.
Decisions:
The value of x is five, which means five H2O molecules are combined with one Cu(II) sulfate. Copper(II) sulfate can react with hydrochloric acid, producing a green composite with dative bonds. Similarly, the reaction between Cu(II) sulfate and ammonia solution is relevant to form bonds and has two steps, forming an indigo complex.
References:
- Bennett, B. (1998) [online] What is Blue Vitriol? General Chemist Online! http://antoine.frostburg.edu/chem/senese/101/inorganic/faq/blue-vitriol.shtml (accessed 2010/1/3).
- Clark, J. (2000) [online] The Colors of Complex Metal Ions. Chemguide. http://www.chemguide.co.uk/inorganic/complexions/colour.html (accessed 2010/1/3).
- Clark, J. (2000) [online] Co-ordinate (Dative Covalent) Bonding. Chemguide. http://www.chemguide.co.uk/atoms/bonding/dative.html (accessed 2009/12/27).
- Farlex Inc. (2009) [online] Hydrate. The Free Dictionary. http://encyclopedia2.thefreedictionary.com/Hydrate+salt (accessed 2009/12/27).
- Lane, R. (2009) Chemistry Practical 2: Complex Ions of Copper (II). Handout.
- Lane, R. (2009) Chemistry Notes.
- Lister, T. and Renshaw, J. (2000) Chemistry for Advanced Level (3rd Edition). London: Stanley Thornes (Publishers) Ltd.
