In this article we are going to investigate the Manganese solubility in natural water with the help of various Eh-pH diagrams. We will also investigate the effect of bicarbonate (alkalinity) and sulfate on Manganese solubility. At the end we will compare the solubility of manganese against iron solubility.

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**Background:**

In natural surface and groundwater, manganese is usually found in the rage of 0.01 ppm to 0.1 ppm. Stability fields on an **Eh-pH diagram **vary depending on the concentration of manganese. We will build three simple cases to study the solubility of manganese. Each case is subdivided in to two Types – one with 0.01 ppm of Mn and another with 0.1 ppm of Manganese.

**Eh-pH diagrams that we will generate will have:**

1 (a): Case 1 Type 1: Mn+2=0.01 ppm; HCO3-=0 ppm; SO4–=0

1 (b): Case 1 Type 2: Mn+2=0.1 ppm; HCO3-=0 ppm; SO4–=0

2 (a): Case 2 Type 1: Mn+2=0.01 ppm; HCO3-=150 ppm; SO4–=0

2 (b): Case 2 Type 2: Mn+2=0.1 ppm; HCO3-=0 ppm; SO4–=0

3 (a): Case 3 Type 1: Mn+2=0.01 ppm; HCO3-=150 ppm; SO4–=300 ppm

3 (b): Case 3 Type 2: Mn+2=0.1 ppm; HCO3-=0 ppm; SO4–= 300 ppm

In short, we will start with an initial solution with Manganese only. It will help us understand the solubility of manganese in pure water. In case 2, we will identify the effect of bicarbonate and in the third case we will study the effect of added sulfate.

Finally we will build a special case for iron with:

4 – Case (4) Special case: Fe+2=0.1 ppm, HCO3-=0 ppm; SO4–= 300 ppm

We will compare case 4 with case 3 to check solubility of Mn and Fe in natural water. Iron is the next member to manganese in the periodic table.

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**Assumption**: For our exercise, we will assume activity of species is equal to morality.

**Main factors that control the geochemistry of manganese in natural water are:**

- Mn species: Mn+2/Mn+3/Mn+4
- redox potential (Eh)
- pH
- Dissolved bicarbonate (HCO3-)
- Dissolved Sulfate (SO4–)

Eh-pH diagrams will be produced using **Geochemist’s Workbench**. Every system is different. So, those diagrams would not exactly represent your field conditions, but should be close. If you are outside the range of Mn, SO4 and HCO3- activities, you need to generate the stability diagrams your self.

As always, all of our stability field calculations and diagrams assumes that **equilibrium **exist in the geochemical system.

Remember that in natural water, there are more species that what we considered in this exercise.

**NOTE:** In natural ground and surface water, Mn concentration usually varies between 0.01 to .1 ppm.

**CASE 1**

**First, let’s assume there is no HCO3- and SO4-2 present in the system and the only cation present is Manganese.**

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**Main governing equations:**

**Mn++ + H2O + .5 O2(aq) = Pyrolusite (MnO2) + 2 H+ ———Equation 1**

- Polynomial fit: log K = 1.546 – .004138 × T – 1.133e-5 × T^2 – 4.69e-8 × T^3 + 2.142e-10 × T^4
- Equilibrium equation: log K = – log a[Mn++] + 2 × log a[H+] – log a[H2O] – .5 × log a[O2(aq)]

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**Mn++ + H2O + .25 O2(aq) = .5 Bixbyite (Mn2O3) + 2 H+ —————-Equation 2**

- Polynomial fit: log K = -4.079 + .02119 × T – 8.2e-5 × T^2 + 1.066e-7 × T^3 + 5.826e-11 × T^4
- Equilibrium equation: log K = – log a[Mn++] + 2 × log a[H+] – log a[H2O] – .25 × log a[O2(aq)]

**Mn++ + H2O + .1667 O2(aq) = .3333 Hausmannite (Mn3O4) + 2 H+….Equation 3**

- Polynomial fit: log K = -6.838 + .03192 × T – .0001124 × T^2 + 1.722e-7 × T^3 – 7.928e-12 × T^4
- Equilibrium equation: log K = – log a[Mn++] + 2 × log a[H+] – log a[H2O] – .1667 × log a[O2(aq)]

**Mn++ + 2 H2O = Mn(OH)2(am) + 2 H+ …………..Equation 4**

- Polynomial fit: log K = -16.87 + .06945 × T – .0002745 × T^2 + 6.336e-7 × T^3 – 6.29e-10 × T^4
- Equilibrium equation: log K = – log a[Mn++] + 2 × log a[H+] – 2 × log a[H2O]

**Mn++ + 1.5 H2O = .5 Mn2(OH)3+ + 1.5 H+……Equation 5**

- Polynomial fit: log K = -11.95 + 0 × T + 0 × T^2 + 0 × T^3 + 0 × T^4
- Equilibrium equation: log K = – log a[Mn++] + .5 × log a[Mn2(OH)3+] + 1.5 × log a[H+]- 1.5 × log a[H2O]

**Mn++ + 3 H2O = Mn(OH)3- + 3 H+ ……….Equation 6**

- Polynomial fit: log K = -34.21 + 0 × T + 0 × T^2 + 0 × T^3 + 0 × T^4
- Equilibrium equation: log K = – log a[Mn++] + log a[Mn(OH)3-] + 3 × log a[H+]- 3 × log a[H2O]

**Note:** Let’s take a close look at the equations. All of the equations are written in terms of O2 and H+ which can be converted into equations with Eh and pH. Equation 1, 2 and 3 have both O2 and H+ as part of the equation. So, in an Eh-pH diagram their contact would be diagonal (you can calculate the actual slope, but I will explain that in another article!). Now equation 4,5 and 6 have only H+ component – Eh would not have any effect on the stability fields. and they would appear as vertical contacts between each phases.

**Figure 1 -Case 1: Type 1 – Mn=0.01 ppm (No HCO3- and SO4-2)**

This diagram is created assuming dissolved Manganese concentration= 0.01 ppm.

**Figure 2 -Case 1: Type 2 – Mn=.1 ppm (No HCO3- and SO4-2)**

This diagram is created assuming dissolved Manganese concentration= 0.1 ppm.

Figure 1 and 2 shows the solid and dissolved phases of Manganese in natural water. However, this system is never found in nature as we have not considered presence of any other cations or anions. But the diagrams show that Manganese is highly soluble over a wide range of pH and Eh. Between pH of 5-10, Manganese is widely soluble and the solubility of Manganese is enhanced by the formation of soluble complex ions such as Mn2(OH)3+.

**The Eh of neutral aerated water is about 0.5 V.****The pH range of natural water: 5-10****Mn is extremely soluble at pH 7, Eh 0.5**

Well, in nature, Mn is found with lower concentration levels that it appears from the presented Eh-pH diagrams.

**CASE 2**

**Le’s add 150 ppm HCO3 to the solution above. In nature, almost always we have some amount of alkalinity present. **

**Figure3 -Case 2: Type 1 – Mn=0.01 ppm, HCO3-=150 ppm**

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** ****Figure4 -Case 2: Type 2 – Mn=0.1 ppm, HCO3-=150 ppm
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**New Reaction**

**Mn++ + HCO3- = Rhodochrosite + H+ ……………….Equation 7**

- Polynomial fit: log K = -.05455 + .01181 × T + 9.001e-6 × T^2 – 1.601e-7 × T^3 + 4.535e-10 × T^4
- Equilibrium equation:log K = – log a[Mn++] + log a[H+] – log a[HCO3-]

Now, the two diagrams above decrease the soluble phases of Manganese by great extent. For example, when Mn concentration is 0.1 ppm and pH at 8 or more,** Rhodochrosite** becomes stable. So, raising pH from below neutral to +7 would precipitate the mineral. The stability field for **Rhodochrosite** is smaller when Mn concentration is 0.01 ppm.** In short, in presence of bicarbonate, Mn is no longer capable in staying complex ion forms. **

**The Eh of neutral aerated water is about 0.5 V.****The pH range of natural water: 5-10****Mn should precipitate at pH 8, Eh ~ 0.5**

**Case 3**

**Le’s add 300 ppm of SO4– in the system. We already had 150 ppm HCO3 in the initial solution. **

**Figure 5 -Case 3: Type 1 – Mn=0.01 ppm, HCO3-=150 ppm, SO4=300 ppm
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**Figure 6 -Case 3: Type 2 – Mn=0.01 ppm, HCO3-=150 ppm, SO4=300 ppm
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** **Figure 5 and 6 shows that in the presence of sulfate, the mineral alabandite could precipitate if enough oxygen is present in the system. Alabandite is however usually found at hydrothermal deposits. In reality, alabandite may not form at all. In short, if alabandite could not precipitate, then SO4 has no control on Mn solubility. ** **

**Remediation for Manganese? The Eh-pH diagrams clearly shows that raising pH of the solution would very much likely to precipitate Manganese out of the solution. So, we can conclude that in common lime applications if the pH is raised above 8, it could precipitate Manganese effectively. **

**Now, lets check the solubility diagram for iron at 0.1 ppm concentration. HCO3=150, SO4–=300.**

**Figure 7 -Special Case Fe=0.1 ppm, HCO3-=150 ppm, SO4=300 ppm
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** **Figure 7 shows that Fe is much less soluble than Manganese and if the pH is raised, Fe-minerals would precipitate earlier leaving Manganese in the solution. ** **

Some of the minerals:

- Hausmannite is a complex oxide of manganese containing both di- and tri-valent manganese. The formula can be represented as Mn2+Mn3+2O4 (Mn3O4)
- Bixbyite is a manganese iron oxide mineral with formula: (Mn,Fe)2O3. The iron:manganese ratio is quite variable and many specimens have almost
**no iron**leading to Mn2O3 **Pyrolusite**is a mineral consisting essentially of manganese dioxide (MnO_{2}) and is important as an ore of manganese.**Rhodochrosite**is a manganese carbonate mineral with chemical composition MnCO3. In its (rare) pure form, it is typically a rose-red color, but impure specimens can be shades of pink to pale brown.

If you want to learn the math behind the Eh-pH diagrams, please follow the exercise listed at http://www.umt.edu/geosciences/faculty/moore/G431/lectur7.htm\

- OSHA Chemical Sampling Information
- OSHA Methods
- Metal & Metalloid Particulates in Workplace Atmospheres (Atomic Absorption). ID-121, (Revised 2002).
- Metal and Metalloid Particulates in Workplace Atmospheres (ICP Analysis). ID-125G, (Revised 2002).

Manganese Toxicology report: http://www.atsdr.cdc.gov/toxprofiles/tp151.html