Systematic Study of Iron Control Chemicals Used During Well Stimulation
- K.C. Taylor (Saudi Aramco) | H.A. Nasr-El-Din (Saudi Aramco) | M.J. Al-Alawi (Saudi Aramco)
- Document ID
- Society of Petroleum Engineers
- SPE Journal
- Publication Date
- March 1999
- Document Type
- Journal Paper
- 19 - 24
- 1999. Society of Petroleum Engineers
- 3 Production and Well Operations, 5.4.10 Microbial Methods, 5.3.1 Flow in Porous Media, 4.1.2 Separation and Treating, 3.2.4 Acidising, 4.2.3 Materials and Corrosion, 5.8.7 Carbonate Reservoir, 5.2 Reservoir Fluid Dynamics, 4.3.4 Scale, 4.1.5 Processing Equipment, 2.2.2 Perforating, 1.8 Formation Damage
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A new type of experimental procedure was developed to study iron control chemicals in acidizing treatments (citric acid, nitrilotriacetic acid, and acetic acid) at temperatures from 25 to 95°C. This procedure gives plots of iron in solution versus equilibrium pH, while previous work has generally studied iron concentration in completely spent acid solutions. Hydrochloric acid concentrations of 7.5 to 28 wt %, sodium chloride concentrations of 0 and 5 wt %, and iron (III) concentrations from 1000 to 10 000 mg/kg were examined with varying concentrations of commonly used iron control chemicals. A systematic study of this kind has not previously been reported in the literature. The following new results were found:
- Iron (III) hydroxide precipitated from spent acid at much lower values of pH than are generally believed in the literature. The general belief is that iron (III) hydroxide precipitation begins at a pH of about 2.2 and is complete at a pH of 3.3. From experimental results, precipitation begins at pH values of about 1 and is nearly complete by pH 2 at 25°C.
- As the temperature was increased, iron (III) became significantly less soluble in the spent acid. This means that iron (III) hydroxide precipitation is potentially more damaging to the formation than generally believed.
- Time-dependent iron precipitation can occur in the absence of iron control chemicals.
- Time-dependent precipitation of ferric/calcium complexes with nitrilotriacetic acid (NTA) and citric acid was found. In partially spent acid, precipitation of previously unreported complexes containing an iron/calcium/chelating agent in a 1/1/1 mole ratio was observed. This behavior is important because of potential formation damage during the acidizing treatment itself, or during shut-in of the well after stimulation.
- Synergism with citric acid/acetic acid mixtures was found. These two iron control chemicals were able to prevent iron precipitation better in combination than individually.
Iron compounds that precipitate during acidizing can reduce reservoir permeability in the critical near-wellbore area. This iron can originate from contaminated acid, dissolution of rust in the coiled tubing or well casing, from iron-containing minerals in the formation, from corrosion products present in the wellbore, or from surface equipment/tanks used during acid jobs.
Iron can be present in either the +2 or the +3 oxidation state. At equilibrium, all iron will be in the +3 state in an oxidizing environment. This will occur in the presence of air in surface facilities. In a reducing environment, the +2 state is favored (normal reservoir conditions or in the presence of hydrogen sulfide). It has been observed in field samples that the ratio of iron (II) to iron (III) in spent acidizing fluids is about 5 to 1. 1,2 This ratio can vary significantly, depending on the nature of the well and the formation. Both iron (II) and iron (III) can lead to precipitate formation under certain conditions. However, most damage will occur from precipitation of iron (III) hydroxide because of its lower solubility in spent acid.
Fig. 1 shows some common iron precipitation reactions. Iron (III) hydroxide precipitates as the pH increases above 1. Iron (II) hydroxide precipitates at pH values greater than about 6, making it less of a problem in acidizing. However, if hydrogen sulfide is present, then iron (II) can precipitate as iron (II) sulfide. It is also possible for iron (III) to be reduced by hydrogen sulfide, leading to precipitation of elemental sulfur. Sulfur in the formation is very difficult to remove, because it is not soluble in acids. Iron (II) can also form iron (II) carbonate as the acid is spent.
Fig. 2 shows the precipitation equilibria for iron (III) hydroxide. The dotted lines are not bonds, they are used to help outline the octahedral structure of the molecule. It is important to remember that in solution, iron (III) is generally six-coordinate, and that the water and hydroxide groups are not tightly bound to the iron. They can exchange with each other or with other groups in solution. As hydroxide exchanges with the water molecules, the charge of the iron complex decreases. The neutral ferric hydroxide can easily precipitate from solution.
Sources of Iron.
At each stage of an acidizing treatment, there is potential for contamination of the acid with iron compounds. Before injection, acid can dissolve rust in storage or mixing tanks. 1-3 Rust dissolution leads to a mixture of iron (II) and iron (III) in solution, but dissolved oxygen in the acid will rapidly oxidize iron (II) to iron (III).
During acid injection, millscale in new tubing,3 or iron-containing corrosion products in tubing 1,2,4,5 can be dissolved, resulting in large amounts of iron in solution. Corrosion products formed with oxygen in injection wells contain a mixture of Fe (II) and Fe (III). Corrosion products formed in production wells, even with trace amounts of H2S will be almost completely iron (II) compounds. Acid corrosion of steel tubing can occur to produce iron (II), and metallic iron can react with iron (III) to produce iron (II).1,2
Coulter and Gougler6 and Gougler et al.7 measured iron content in acid solutions at four different points in the acidizing process. They examined cleanup treatments of new wells only, before perforating. Acid was pumped down new tubing and up the annulus. They measured the iron content of the raw acid before transport, after transport, at the well head, and at the return line. Iron concentrations at the wellhead varied from 200 to 3500 mg/L, while returning acid showed iron concentrations of 9000 to 100 000 mg/L. Coulter and coworkers recommend pickling the tubing and annulus without the addition of iron control chemicals. The acid treatment can then be done with iron control additives sufficient to handle 1000 to 2000 mg/L of dissolved iron.
Walker et al.5 showed how thin deposits of iron sulfide scale on tubing can lead to large quantities of iron (II) in solution during acid treatment. Iron sulfide reprecipitation becomes a concern when sulfide scales are present, even if hydrogen sulfide concentrations are very low. Various types of iron-containing scale can be present in production tubing, and these have been summarized by Walker et al.5
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