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Mathew Mahindaratne
Mathew Mahindaratne
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If not adding it (nitric acid), does it affect our concentrate solution later?

Yes it does because slow hydrolysis with water, some $\ce{Fe^3+}$ would separate out from the solution as solid particles (Ref.1). This review focuses on studies on the hydrolysis and precipitation of $\ce{Fe^3+}$ from aqueous solutions of its inorganic salts, chiefly the nitrate, perchlorate, chloride, and sulfate. Note that the $\ce{Fe^3+}$ ion has been shown to exist as the octahedral $\ce{[Fe(OH2)6]^3+}$ complex in aqueous solutions of $\ce{Fe(NO3)3.9H2O}$ (Ref.2).

The hydrolysis reactions of these inorganic salt solutions, most of which occur slowly, consists of several steps: (1) formation of low-molecular-weight species; (2) formation of a red cationic polymer; (3) aging of the polymer, with eventual conversion to oxide phases; and (4) precipitation of oxide phases directly from low-molecular-weight precursors.

The characterized crystalline iron(III) oxides and hydrous oxides are $\ce{Fe2O3}$ and $\ce{FeO(OH)}$, each of which is polymorphic and very insoluable in water. For example, the $\ce{\alpha-Fe2O3}$ (hematite) has a $\mathrm{p}K_\mathrm{sp}$ of $41.7$ while that of $\alpha-, \beta-,$ and $\gamma-$ phases of $\ce{FeO(OH)}$ ((goethite, akaganeite, and lepidocrocite) are $41.7, 36,$ and $38.7$, respectively. Thus, one can expect after preparing in aqueous solutions, $\ce{Fe^3+}$ can lost significant amount to hydrolysis.

The first step of the hydrolysis of inorganic Fe(III) solutions is the formation of low-molecular-weight species. Two such reactions are given in Ref.1 as: $$\ce{Fe(OH2)6]^3+ + H2O <=> Fe(OH2)5(OH)]^2+ + H3O+} \qquad K_\mathrm{eq} = 3.5 \tag2$$$$\ce{Fe(OH2)6]^3+ + H2O <=> Fe(OH2)5(OH)]^2+ + H3O+} \qquad K_\mathrm{eq} = 2.2 \tag1$$ $$\ce{Fe(OH2)5(OH)]^2+ + H2O <=> Fe(OH2)4(OH)2]^+ + H3O+} \qquad K_\mathrm{eq} = 2.2 \tag1$$$$\ce{Fe(OH2)5(OH)]^2+ + H2O <=> Fe(OH2)4(OH)2]^+ + H3O+} \qquad K_\mathrm{eq} = 3.5 \tag2$$

These formation reactions of low-molecular-weight species can be revised by adding appropriate strong acid (here it is nitric acid since we used $\ce{Fe(NO3)3}$) before they react further to make polymeric iron salts. Note that Hydrolyzed Fe(III) solutions containing the red polymer are not at equilibrium (Ref.1).

The polymer separated from hydrolysis has been studied to find the rate of their degradation by $\ce{HNO3}$, which results in immediate precipitation of an undetermined fraction of the iron. Thus, it should be advised that addition of the nitric acid should be immediate after the preparation of iron solution.

References:

  1. Charles M. Flynn Jr., "Hydrolysis of inorganic iron (III) salts," Chem. Rev. 1984, 84(1), 31-41 (DOI: https://doi.org/10.1021/cr00059a003).
  2. Neil J. Hair and James K. Beattie, "Structure of hexaaquairon(III) nitrate trihydrate. Comparison of iron(II) and iron(III) bond lengths in high-spin octahedral environments," Inorg. Chem. 1977, 16(2), 245-250 (DOI: https://doi.org/10.1021/ic50168a006).
  3. Barbara A. Sommer, Dale W. Margerum, John Renner, Paul Saltman, and Thomas G. Spiro, "Reactivity and aging in hydroxy-iron(III) polymers, analogs of ferritin cores," Bioinorganic Chemistry 1973, 2(4), 295-309 (DOI: https://doi.org/10.1016/S0006-3061(00)80202-1).

If not adding it (nitric acid), does it affect our concentrate solution later?

Yes it does because slow hydrolysis with water, some $\ce{Fe^3+}$ would separate out from the solution as solid particles (Ref.1). This review focuses on studies on the hydrolysis and precipitation of $\ce{Fe^3+}$ from aqueous solutions of its inorganic salts, chiefly the nitrate, perchlorate, chloride, and sulfate. Note that the $\ce{Fe^3+}$ ion has been shown to exist as the octahedral $\ce{[Fe(OH2)6]^3+}$ complex in aqueous solutions of $\ce{Fe(NO3)3.9H2O}$ (Ref.2).

The hydrolysis reactions of these inorganic salt solutions, most of which occur slowly, consists of several steps: (1) formation of low-molecular-weight species; (2) formation of a red cationic polymer; (3) aging of the polymer, with eventual conversion to oxide phases; and (4) precipitation of oxide phases directly from low-molecular-weight precursors.

The characterized crystalline iron(III) oxides and hydrous oxides are $\ce{Fe2O3}$ and $\ce{FeO(OH)}$, each of which is polymorphic and very insoluable in water. For example, the $\ce{\alpha-Fe2O3}$ (hematite) has a $\mathrm{p}K_\mathrm{sp}$ of $41.7$ while that of $\alpha-, \beta-,$ and $\gamma-$ phases of $\ce{FeO(OH)}$ ((goethite, akaganeite, and lepidocrocite) are $41.7, 36,$ and $38.7$, respectively. Thus, one can expect after preparing in aqueous solutions, $\ce{Fe^3+}$ can lost significant amount to hydrolysis.

The first step of the hydrolysis of inorganic Fe(III) solutions is the formation of low-molecular-weight species. Two such reactions are given in Ref.1 as: $$\ce{Fe(OH2)6]^3+ + H2O <=> Fe(OH2)5(OH)]^2+ + H3O+} \qquad K_\mathrm{eq} = 3.5 \tag2$$ $$\ce{Fe(OH2)5(OH)]^2+ + H2O <=> Fe(OH2)4(OH)2]^+ + H3O+} \qquad K_\mathrm{eq} = 2.2 \tag1$$

These formation reactions of low-molecular-weight species can be revised by adding appropriate strong acid (here it is nitric acid since we used $\ce{Fe(NO3)3}$) before they react further to make polymeric iron salts. Note that Hydrolyzed Fe(III) solutions containing the red polymer are not at equilibrium (Ref.1).

The polymer separated from hydrolysis has been studied to find the rate of their degradation by $\ce{HNO3}$, which results in immediate precipitation of an undetermined fraction of the iron. Thus, it should be advised that addition of the nitric acid should be immediate after the preparation of iron solution.

References:

  1. Charles M. Flynn Jr., "Hydrolysis of inorganic iron (III) salts," Chem. Rev. 1984, 84(1), 31-41 (DOI: https://doi.org/10.1021/cr00059a003).
  2. Neil J. Hair and James K. Beattie, "Structure of hexaaquairon(III) nitrate trihydrate. Comparison of iron(II) and iron(III) bond lengths in high-spin octahedral environments," Inorg. Chem. 1977, 16(2), 245-250 (DOI: https://doi.org/10.1021/ic50168a006).
  3. Barbara A. Sommer, Dale W. Margerum, John Renner, Paul Saltman, and Thomas G. Spiro, "Reactivity and aging in hydroxy-iron(III) polymers, analogs of ferritin cores," Bioinorganic Chemistry 1973, 2(4), 295-309 (DOI: https://doi.org/10.1016/S0006-3061(00)80202-1).

If not adding it (nitric acid), does it affect our concentrate solution later?

Yes it does because slow hydrolysis with water, some $\ce{Fe^3+}$ would separate out from the solution as solid particles (Ref.1). This review focuses on studies on the hydrolysis and precipitation of $\ce{Fe^3+}$ from aqueous solutions of its inorganic salts, chiefly the nitrate, perchlorate, chloride, and sulfate. Note that the $\ce{Fe^3+}$ ion has been shown to exist as the octahedral $\ce{[Fe(OH2)6]^3+}$ complex in aqueous solutions of $\ce{Fe(NO3)3.9H2O}$ (Ref.2).

The hydrolysis reactions of these inorganic salt solutions, most of which occur slowly, consists of several steps: (1) formation of low-molecular-weight species; (2) formation of a red cationic polymer; (3) aging of the polymer, with eventual conversion to oxide phases; and (4) precipitation of oxide phases directly from low-molecular-weight precursors.

The characterized crystalline iron(III) oxides and hydrous oxides are $\ce{Fe2O3}$ and $\ce{FeO(OH)}$, each of which is polymorphic and very insoluable in water. For example, the $\ce{\alpha-Fe2O3}$ (hematite) has a $\mathrm{p}K_\mathrm{sp}$ of $41.7$ while that of $\alpha-, \beta-,$ and $\gamma-$ phases of $\ce{FeO(OH)}$ ((goethite, akaganeite, and lepidocrocite) are $41.7, 36,$ and $38.7$, respectively. Thus, one can expect after preparing in aqueous solutions, $\ce{Fe^3+}$ can lost significant amount to hydrolysis.

The first step of the hydrolysis of inorganic Fe(III) solutions is the formation of low-molecular-weight species. Two such reactions are given in Ref.1 as: $$\ce{Fe(OH2)6]^3+ + H2O <=> Fe(OH2)5(OH)]^2+ + H3O+} \qquad K_\mathrm{eq} = 2.2 \tag1$$ $$\ce{Fe(OH2)5(OH)]^2+ + H2O <=> Fe(OH2)4(OH)2]^+ + H3O+} \qquad K_\mathrm{eq} = 3.5 \tag2$$

These formation reactions of low-molecular-weight species can be revised by adding appropriate strong acid (here it is nitric acid since we used $\ce{Fe(NO3)3}$) before they react further to make polymeric iron salts. Note that Hydrolyzed Fe(III) solutions containing the red polymer are not at equilibrium (Ref.1).

The polymer separated from hydrolysis has been studied to find the rate of their degradation by $\ce{HNO3}$, which results in immediate precipitation of an undetermined fraction of the iron. Thus, it should be advised that addition of the nitric acid should be immediate after the preparation of iron solution.

References:

  1. Charles M. Flynn Jr., "Hydrolysis of inorganic iron (III) salts," Chem. Rev. 1984, 84(1), 31-41 (DOI: https://doi.org/10.1021/cr00059a003).
  2. Neil J. Hair and James K. Beattie, "Structure of hexaaquairon(III) nitrate trihydrate. Comparison of iron(II) and iron(III) bond lengths in high-spin octahedral environments," Inorg. Chem. 1977, 16(2), 245-250 (DOI: https://doi.org/10.1021/ic50168a006).
  3. Barbara A. Sommer, Dale W. Margerum, John Renner, Paul Saltman, and Thomas G. Spiro, "Reactivity and aging in hydroxy-iron(III) polymers, analogs of ferritin cores," Bioinorganic Chemistry 1973, 2(4), 295-309 (DOI: https://doi.org/10.1016/S0006-3061(00)80202-1).
Source Link
Mathew Mahindaratne
Mathew Mahindaratne
  • 41k
  • 28
  • 55
  • 109

If not adding it (nitric acid), does it affect our concentrate solution later?

Yes it does because slow hydrolysis with water, some $\ce{Fe^3+}$ would separate out from the solution as solid particles (Ref.1). This review focuses on studies on the hydrolysis and precipitation of $\ce{Fe^3+}$ from aqueous solutions of its inorganic salts, chiefly the nitrate, perchlorate, chloride, and sulfate. Note that the $\ce{Fe^3+}$ ion has been shown to exist as the octahedral $\ce{[Fe(OH2)6]^3+}$ complex in aqueous solutions of $\ce{Fe(NO3)3.9H2O}$ (Ref.2).

The hydrolysis reactions of these inorganic salt solutions, most of which occur slowly, consists of several steps: (1) formation of low-molecular-weight species; (2) formation of a red cationic polymer; (3) aging of the polymer, with eventual conversion to oxide phases; and (4) precipitation of oxide phases directly from low-molecular-weight precursors.

The characterized crystalline iron(III) oxides and hydrous oxides are $\ce{Fe2O3}$ and $\ce{FeO(OH)}$, each of which is polymorphic and very insoluable in water. For example, the $\ce{\alpha-Fe2O3}$ (hematite) has a $\mathrm{p}K_\mathrm{sp}$ of $41.7$ while that of $\alpha-, \beta-,$ and $\gamma-$ phases of $\ce{FeO(OH)}$ ((goethite, akaganeite, and lepidocrocite) are $41.7, 36,$ and $38.7$, respectively. Thus, one can expect after preparing in aqueous solutions, $\ce{Fe^3+}$ can lost significant amount to hydrolysis.

The first step of the hydrolysis of inorganic Fe(III) solutions is the formation of low-molecular-weight species. Two such reactions are given in Ref.1 as: $$\ce{Fe(OH2)6]^3+ + H2O <=> Fe(OH2)5(OH)]^2+ + H3O+} \qquad K_\mathrm{eq} = 3.5 \tag2$$ $$\ce{Fe(OH2)5(OH)]^2+ + H2O <=> Fe(OH2)4(OH)2]^+ + H3O+} \qquad K_\mathrm{eq} = 2.2 \tag1$$

These formation reactions of low-molecular-weight species can be revised by adding appropriate strong acid (here it is nitric acid since we used $\ce{Fe(NO3)3}$) before they react further to make polymeric iron salts. Note that Hydrolyzed Fe(III) solutions containing the red polymer are not at equilibrium (Ref.1).

The polymer separated from hydrolysis has been studied to find the rate of their degradation by $\ce{HNO3}$, which results in immediate precipitation of an undetermined fraction of the iron. Thus, it should be advised that addition of the nitric acid should be immediate after the preparation of iron solution.

References:

  1. Charles M. Flynn Jr., "Hydrolysis of inorganic iron (III) salts," Chem. Rev. 1984, 84(1), 31-41 (DOI: https://doi.org/10.1021/cr00059a003).
  2. Neil J. Hair and James K. Beattie, "Structure of hexaaquairon(III) nitrate trihydrate. Comparison of iron(II) and iron(III) bond lengths in high-spin octahedral environments," Inorg. Chem. 1977, 16(2), 245-250 (DOI: https://doi.org/10.1021/ic50168a006).
  3. Barbara A. Sommer, Dale W. Margerum, John Renner, Paul Saltman, and Thomas G. Spiro, "Reactivity and aging in hydroxy-iron(III) polymers, analogs of ferritin cores," Bioinorganic Chemistry 1973, 2(4), 295-309 (DOI: https://doi.org/10.1016/S0006-3061(00)80202-1).