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Microchimica Acta

, Volume 75, Issue 1–2, pp 63–72 | Cite as

Semimicro determination of solubility constants: Copper(II) carbonate and iron(II) carbonate

  • F. Reiterer
  • W. Johannes
  • H. Gamsjäger
Article

Summary

A semimicro technique for the determination of solubility constants of metal oxides, hydroxides, carbonates and sulphides has been developed and successfully applied to copper(II) and iron(II) carbonate. All experimental data can be interpreted according to the heterogeneous equilibria
$$\begin{gathered} CuCO_{3(s)} + 2H_{(l)}^ + \rightleftharpoons Cu_{(l)}^{2 + } + CO_{2(g)} + H_2 O_{(l)} ; \hfill \\ ^* K_{ps0} (CuCO_3 ) = [Cu^{2 + } ] \cdot p_{CO_2 } \cdot [H^ + ]^{ - 2} , \hfill \\ FeCO_{3(s)} + 2H_{(l)}^ + \rightleftharpoons Fe_{(l)}^{2 + } + CO_{2(g)} + H_2 O_{(l)} ; \hfill \\ ^* K_{ps0} (FeCO_3 ) = [Fe^{2 + } ] \cdot p_{CO_2 } \cdot [H^ + ]^{ - 2} . \hfill \\ \end{gathered} $$
The individual solubility constants obtained are (errors 1δ):
$$\begin{gathered} \log ^* K_{ps 0} (CuCO_3 ) = 6.95 \pm 0.04, 25^0 C, I = 0.20M (NaClO_4 ), \hfill \\ \log ^* K_{ps 0} (FeCO_3 ) = 7.61 \pm 0.05, 50^0 C, I = 1.0 mole/kg (NaClO_4 ). \hfill \\ \end{gathered} $$
With these values the free enthalpies of formation of copper(II) and iron (II) carbonate have been calculated as
$$\begin{gathered} \Delta _f G_{298}^ \ominus (CuCO_3 ) = - 126.21 \pm 0.06 kcal/mole \hfill \\ \Delta _f G_{298}^ \ominus (FeCO_3 ) = - 159.90 \pm 0.11 kcal/mole. \hfill \\ \end{gathered} $$
Thermodynamic implications of these results are discussed.

Keywords

Oxide Iron Copper Experimental Data Sulphide 
These keywords were added by machine and not by the authors. This process is experimental and the keywords may be updated as the learning algorithm improves.

Halbmikrobestimmung der Löslichkeitskonstanten von Kupfer(II)carbonat und Eisen(II)carbonat

Zusammenfassung

Eine Semimikro-Technik zur Bestimmung der Löslichkeitskonstanten von Metalloxiden, -hydroxiden, -carbonaten und -sulfiden wurde entwickelt und mit Erfolg auf Kupfer (II)- und Eisen (II)-carbonat angewendet. Alle experimentellen Daten konnten mit Hilfe der beiden folgenden heterogenen Gleichgewichte erklärt werden
$$\begin{gathered} CuCO_{3(s)} + 2H_{(l)}^ + \rightleftharpoons Cu_{(l)}^{2 + } + CO_{2(g)} + H_2 O_{(l)} ; \hfill \\ ^* K_{ps0} (CuCO_3 ) = [Cu^{2 + } ] \cdot p_{CO_2 } \cdot [H^ + ]^{ - 2} \hfill \\ FeCO_{3(s)} + 2H_{(l)}^ + \rightleftharpoons Fe_{(l)}^{2 + } + CO_{2(g)} + H_2 O_{(l)} ; \hfill \\ ^* K_{ps0} (FeCO_3 ) = [Fe^{2 + } ] \cdot p_{CO_2 } \cdot [H^ + ]^{ - 2} . \hfill \\ \end{gathered} $$
Folgende Löslichkeitskonstanten wurden gefunden (Fehlergrenze 1δ):
$$\begin{gathered} \log ^* K_{ps 0} (CuCO_3 ) = 6,95 \pm 0,04; 25^0 C, I = 0.20 mol dm^{ - 3} (NaClO_4 ), \hfill \\ \log ^* K_{ps 0} (FeCO_3 ) = 7,61 \pm 0,05; 50^0 C, I = 1,0 mol kg^{ - 1} (NaClO_4 ). \hfill \\ \end{gathered} $$
Mit diesen Werten ergaben sich die freien Bildungsenthalpien von Kupfer (II)-und Eisen(II)-carbonat zu
$$\begin{gathered} \Delta _f G_{298}^ \ominus (CuCO_3 ) = - 126,21 \pm 0,06 kcal mol^{ - 1} , \hfill \\ \Delta _f G_{298}^ \ominus (FeCO_3 ) = - 159,90 \pm 0,11 kcal mol^{ - 1} . \hfill \\ \end{gathered} $$
Thermodynamische Konsequenzen dieser Resultate werden diskutiert.

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Copyright information

© Springer-Verlag 1981

Authors and Affiliations

  • F. Reiterer
    • 1
    • 2
    • 3
  • W. Johannes
    • 1
    • 2
    • 3
  • H. Gamsjäger
    • 1
    • 2
    • 3
  1. 1.Institut für Physikalische ChemieMontanuniversitätLeoben
  2. 2.Mineralogisches Institut der Technischen UniversitätHannover
  3. 3.Institut für Physikalische ChemieMontanuniversitätLeoben

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