Reaction+Mechanisms

Introduction: The major problem for offshore oil recovery companies is the formation of barium sulphate (barite) in the tubes. This occurs when the SO4²- rich sea water comes in contact with the Ba2+ contained in the brine to form the insoluble barium sulphate. This section will discuss the reaction mechanism behind the formation of barium sulphate and the different types of chemical reagents needed to remove barium sulphate.

Formation:
The formation of barium sulphate is a simple exothermic addition reaction that’s given by the following equation. Note: Other precipitates such as calcium carbonate and calcium sulphate are also formed in the offshore oil recovery process; however, the removal techniques for these scales is a simple acid dissolution. So, we'll focus our attention here on the mechanisms of how to remove barium sulpahte scale using the chelating technique.

Ba2+(aq) + SO­42-(aq) > BaSO4(s) ΔrHº= - 858.56 kJ/mol



Characteristics of Barium Sulphate:

 * 1) The barium sulphate is a rhomboildal shape structure that’s held together very closely and strongly
 * 2) The pH varies from 6.5-8.0 (close to neutral)
 * 3) The solubility in H2O(l) ≤ 0.2% ,and 0.6% in acetic acid
 * 4) It has an extremely high melting point of 1580 ⁰C
 * 5) The porosity of the barite is very low, which is shown in the above space filling diagram

Due to these characteristics of barium sulphate an acid dissolution is a very ineffective way to remove this type of scale. A more effective approach is to use a chelating agent when you have a very insoluble compound such as barite. Researchers have found that the best chelating agents for removing barite scale are dethylene triamine pentaacetic acid (DTPA) and ethylene diamine tetraacetic acid (ETDA).

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Introduction to EDTA and DTPA:
The purpose of using chelating agents such as **EDTA** and **DTPA** is to turn the insoluble barium sulphate into either barium carbonate or other ionic compounds that are easier to dissolve using an acid. You can control the stability of the metal complex by controlling the hydrogen ion concentration of the enviroment (i.e. pH). This is most commonly done by dissovling **sodium hydroxide** in solution to decrease the pH of the enviroment. A decrease in pH will promote the deprotonation of the **EDTA** and **DTPA** molecules (i.e the removing the hydrogen atoms) to form **(EDTA)4-** or **(DTPA)6-** ions. These ion complexes can now adhere to the barium sulphate and remove the Ba2+ ions; as a result, leaving the SO42- ions in solution. Another unique characteristic about **EDTA** and **DTPA** is that each nitrogen atom can donate its lone pairs to the Ba atom to form an additional single bond. Since **DTPA** has 5[| carboxylic acid]functional groups and 3 [|amine] functional groups it can form 8 bonds to the Ba, whereas **EDTA** can only form 6. The full octet (i.e. 8 bonds) stabilizes the the Ba better then the 6 bonds, making **DTPA** a better chelating agent then **EDTA** for barium sulphate. An example of the **(EDTA)4-** ion complex stabilizing a metal ion is shown in (Figure 4). Also, the stereochemistry of the **Ba(EDTA)** and **Ba(DTPA)** molecules make it very difficult to remove the Ba atom, because of the large oxygen atoms staggered around the Ba. This prevents any sort of nucleophilic substitution or nucleophilic elimination reaction to occur.



media type="file" key="EDTA complex.swf" width="360" height="360" align="center"

<span style="font-family: Arial,Helvetica,sans-serif;">When using chelating agents such as EDTA and DTPA along with sodium hydroxide in solution to promote the deprotonation of the EDTA or DTPA you end up with the following chemical reaction. Where the Na+(aq) are merely [|spectator ions] for the following reaction. <span style="font-family: Arial,Helvetica,sans-serif;">This reaction can be looked at through a thermodynamic point of view to see what type of pressures and temperaures will drive the reaction forward to precipitate out barium carbonate. Also, Le Chatelier's Principles can be applied to the equilibrium reaction to favour the precipitation of barium carbonate.