EVROPSKÁ UNIE Evropské strukturální a investiční fondy Operační program Výzkum, vývoj a vzdělávání MINISTEHSTVO ŠKOLSTVÍ, MLÁDEŽE A TĚLOVÝCHOVY Hydrogeochemistry 3. Processes in natural water Spring 2022 Tento učební materiál vznikl v rámci projektu Rozvoj doktorského studia chemie č. CZ.02.2.69/0.0/0.0/16_018/0002593 ACID-BASE REACTIONS Br0nsted's definition of acids and bases Acid A substance capable of transferring a proton to another substance. Proton donor. HC1 + H90 -> FUO+ + C1 Base • A substance capable of accepting a proton from another substance. • Proton acceptor. NH3 + H20 -> NH^ + OH Arrhenius acids and bases are also acids and bases according to Br0nsted's definition, but this is not the case retrospectively. Conjugated pairs • Acid - H+ -> base _ HCl + H20 -> H30+ + CI" • Base + H+-> acid ,_, NHo + H90 -> NHÍ +0H" • The same substance can be an acid in one reaction and a base in another. • The designation acid/base is relative and indicates the ability to bind/release a proton. Dissociation of acids HC1 + H20 -> H30+ + cr • It does not run up to 100% and runs simultaneously forward and backwards (like all reactions). • Dissociation leads to the establishment of protolytic equilibrium. • Characterized by an equilibrium constant. = [H30+][C1-] eq [HC1][H20] Dissociation constant In an aqueous environment, we take water a constant and include it directly in the constant, and then we talk about the dissociation constant o/acid _ [H30+][C1-] Ka ~ [HCl] Dissociation of bases Quite analogous to acid dissociation, it is characterized by the equilibrium constant Kpn or the dissociation constant of base K NH3 + H20 -> NHj + OH = [NH+][OH-] [NH+][OH-] e« [NH3][H20] Kb [NH3] Strength of acids and bases • Dissociation constants determine the strength of acids and bases - the intensity of proton release/capture. • For a strong acid, most molecules dissociate. Strong KA/B>10"2 Medium K A/B = 10-2 -10-4 Weak KA/B<10~4 The strength of acids • The degree of dissociation depends on the structure of the acid molecules. • Polybasic acids have multiple dissociation constants. • The polarity of the molecule and its bonds determines the strength. • Few acids dissociate completely and are very strong - HN03, H2S04, HCI04, HCI, HBr, HI • Weak acids dissociate to smaller extent - H2C03, organic acids Water dissociation Water can behave as both acid and base. H20 + H2O^H30+ + OH- [H30+][OH-] Ke« [H2OP The water concentration is considered to be constant and we derive the ionic product of water K^, Kw = [H30+][OH-] pH • For clean water: -At25°C Kw=10-14pHis7(iel00ppbH+) - At 0 °C Kw= 10 "14 9 pH = 7.45 -Atl00°C Kw= 10"13pH = 6.50 • The concentration of H30+ and OH" ranges from 10 to 10"15 mol L-1 • For practical reasons, a logarithmic pH scale has been introduced pH = -log[H30+] f Neutral V [H3Q + J 1.0 10"' 10"1 10 1 10g lQn 1QT ioR 10* ID10 10 11 10 13 10JJ 10"" p\l 0 12 3 4 5 e [7] 6 9 10 11 12 13 14 ____ * A " Neutral — " ^> ^Increasing acidic character Increasing basic character; CARBONATE SYSTEM Carbonates • The pH of most natural waters is controlled by reactions with carbonates • C02(g) dissolves in water to equilibrium in proportion to the partial pressure in air (PC02) C02(#) <-+ C02(aq) • C02 reacts with water to form an acid C02(aJ (a) CI- NO, Ca Mg 2+ 2+ K+ Na+ AZKib) - AZ Aia) = A[AIk] = -A[Acy] (b) Titration curves ! 1 X L X pH j \ OH-Atk 1 Acy U----- C02-Acy ^*""--*<^ I p-Aik 1 Alk (a) 1 *0 «1 "2 Volume of strong acid, CX Figure 4.12. Sketch of an acidimetric titration curve (a). In (b) the results of (a) are plotted in terms of Gran functions: F< is multiplied by scale factors nt. The F, values are defined by Equations 48 and 51-54. jc0t vXf and v2 are the volumes of strong acid corresponding to the equivalence points/ = 2,/ - 1, and / = 0, respectively. INTERACTION WITH CARBONATE MINERALS Calcium carbonates CaC03(s) = C0|" + Ca2+ • The product of solubility is Kc = aCa2+ aC02- = 10"8'48 calcite (25°C) Ka = aCa2+ aC02- = 10-8'34 aragonite (25°C) • We can solve the systems by adding the appropriate equations to the carbonate system - KC02, Kl7K2/ Kw, Kc, PEN and one other condition • Fixed PC02 for open system or sum of carbonates for closed system Example What will be the pH of the water in equilibrium with atmospheric C02and calcite at 25°C, assuming ideal behavior and no other solutes? • 6 unknowns (PC02, [H2C03*], [HC03-], [C032-], [Ca2+] and [H+]) • 6 equations (PC02, Kco2' Kcand PEN) PC02 = io-3-4 PEN: [H+] +2[Ca2+] = [HC031 +2[C0321 + [OH"] — At neutral pH we simplify to 2 [Ca2+] = [HCCV] - Substituting equations by expressing [H+] Example • What will be the pH of the water in equilibrium with atmospheric C02and calcite at 25°C, assuming ideal behavior and no other solutes? • 6 unknowns (PC02, [H2C03*], [HC03-], [C032"], [Ca2+] and [H+]) • 6 equations (PC02, Kco2' Kcand PEN) PC02 = io-3-4 PEN: [H+] +2[Ca2+] = [HC031 +2[C0321 + [OH"] - At neutral pH we simplify to 2 [Ca2+] = [HCCV] - Substituting equations by expressing [H+] - [H+] = 10"8-5 ie pH = 8.5 Dependence of solubility on PC02 ar HCO 3 Ca2+ "HCQ3 _ Kcal K± _ ^-4,30 fQr 25°C aH2CO^ K' 2 aCa2+ aHCQ3 _ Kcai K± KCp2 _ - q-5,97 PC02 K2 decrease in C02 leads to an increase in saturation and vice versa Respiration, decay, photosynthesis The dependence of [Ca2+] on PC02is nonlinear 0.0345 HjCO, CONCENTRATION (mmol/lilerJ 0.345 T 3.0*10 en CD T3 C C0S PftESSJRE (alm.1 Fig. 2.—Changes in composition of carbonated water during equilibration with calcite at 25° C. in the presence and in the absence of a vapor phase. Curves I and II describe the behavior of a solution which was originally in equilibrium with a vapor phase with a COa pressure of 0.10 atm., curves III and IV describe the behavior of a solution that was originally in equilibrium with a vapor phase with a C03 pressure of 0,01 atm. Water mixing 0 0.02 0.04 0.06 0.08 0.10 CD > co2 Mixing usually produces unsaturated solutions Dolomite CaMg(C03)2(s) = COi" + Ca2+ + Mg2+ • The product of solubility is KD = aCa2+aMg2+aC02- = 10"17'2(25°C) • The value of KD is quite uncertain and can vary greatly for dolomites of various origins • At normal temperatures it dissolves only slowly and hardly grows (very unsaturated/supersaturated solutions) Formation of dolomite Dolomite is often formed by alteration of calcite CaMg(C03)2 + Ca2+ = 2CaC03 + Mg2+ _ aMg2+ RCD - ~~- aCa2+ In solutions where the Mg/Ca ratio is greater than KCD, dolomite is more stable than calcite and vice versa However, the process is kinetically very slow Calcite changes to dolomite only under conditions of a very large excess of magnesium Dissolution of dolomite • Analogous to calcite • Significantly slower • The need for extreme retention times to achieve "balance" (thousands of years) — equilibrium • At low temperatures, it should dissolve uncongruently due to kinetics and thermodynamics Magnesium calcite Ca(1_x)MgxC03(s) = C0§" + (1 - x)Ca2++xMg2" • Low magnesium Mg-calcite (<5%) • High-magnesium Mg-calcite (> 10%) - recent deep-sea sediments • All high-magnesium are unstable (conversion to dolomite and low-magnesium), but under surface conditions it inhibits the growth kinetics of dolomite • Mg-calcite in seawater more stable than pure calcite Dissolution of Mg-calcite Equilibrium with solution definable as cation exchange or dissolution - As a result, both processes must be in balance The correctness of the approaches is not clarified, the relationships are very complex The determination of K is complicated by a combination of congruent and non-congruent dissolution 6.5E-03 CO co CO CO co co co o o p p p p p p UJ LU LU LU LU LU LU LU o O o m o LO O lO c\i CO CO CaTot [mol L"1] Properties of water in carbonate aquifers • Major ions Ca2+and HC03_ • Close to balance with calcite • The total mineralization depends mainly on the PC02 in the system - More C02 in the soil than in the air - Decisive role of C02 replenishement in to the solution - With the exception of water of deeper circulation mixed with other sources of minerals The ammount of dissolved calcite is critically dependent available C02gas phase during dissolution HaC03 CONCENTRATION (mmol/liler) C04 PRESSURE {aim.) Fig. 2.—Changes in composition of carbonated water during equilibration with calcite at 25° C. in the presence and in the absence of a vapor phase. Curves I and II describe the behavior of a solution which was originally in equilibrium with a vapor phase with a COn pressure of 0.10 atm., curves III and IV describe the behavior of a solution that was originally in equilibrium with a vapor phase with a C02 pressure of 0.01 atm. Example • Pure water in equilibrium with P C02= 10~2 a) How much calcite dissolves in a closed system? 6 unknowns (PC02, [H2C03*], [HCCy], [C032-], [Ca2+] and [H+]) 6 equations (KC02, K1; K2, Kc, PEN and ZC02) ZC02 = ZC02° + ZC02diss ZC02= [H2C03*]° + [Ca2+] We assume negligible [C032_] in ZC02and obtain [H2C03*]°= [HC031 + [H2C03*] - [Ca2+] Example • Pure water in equilibrium with P C02= 10~2 a) How much calcite dissolves in a closed system? 6 unknowns (PC02, [H2C03*], [HCCy], [C032-], [Ca2+] and [H+]) 6 equations (KC02, K1; K2, Kc, PEN and IC02) ZC02 = ZC02° + ZC02diss ZC02= [H2C03*]° + [Ca2+] We assume negligible [C032~] in ZC02and obtain [H2C03*]°= [HC031 + [H2C03*] - [Ca2+] Substituting the rest of the equations we get [Ca2+] = 3.34 x 10"4 mol/L Dissolve approximately 33.4 mg of calcite per liter of water Example • Pure water in equilibrium with PC02 - b) How much does it dissolve in an open system? - The system we have already solved 6 unknowns (PC02, [H2C03*], [HC03-], [C0321; [Ca2+] and [HI) 6 equations (PC02, Kco2' Ki> K2, Kcand PEN) Example • Pure water in equilibrium with PC02 = 10"2 b) How much does it dissolve in an open system? - The system we have already solved 6 unknowns (PC02, [H2C03*], [HC03"]# [C032-], [Ca2+] and [H+]) 6 equations (PC02/ KC02, K1# K2, Kcand PEN) P = in-2 rC02 ±u From the equations we express [Ca2+] and substitute the parameters [Ca 2+] = 1.39 x 10"3mol/L Dissolve approximately 139 mg of calcite per liter of water Punkva Cave Býčí skála TCI TC2 CP1 RD1 Angel BS1 BS2 ZD1 mS/m (25 conductivity °C) 28.1 54.8 62.5 56.1 61.1 25.7 33.3 49.3 ZNK8.3 mmol/l <0.2 <0.2 <0.2 <0.2 <0.2 - - - KNK4.5 mmol/l 1.91 5.14 6.21 5.78 6.27 2.45 2.6 3.72 total hardness mmol/l 1.43 2.99 3.48 3.15 3.38 1.35 1.7 2.65 Na mg/l 2.1 2.1 1.9 2.3 3 3.8 3.5 - K mg/l 0.86 0.57 0.77 0.71 <0.5 0.5 1.2 - Ca mg/l 54.7 117.5 137.2 124.1 133.5 47.5 61.9 - Ca from the cave mg/l 52.8 124 144 - 137.6 52.8 65.6 - Mg mg/l 1.6 1.4 1.4 1.3 1.2 4 3.8 - S04 mg/l 40.1 38.6 45.6 30.6 30.2 13.9 43 - CI mg/l 9 7 4 5 3 7 4 - N03 mg/l 6.5 3.7 <3 <3 <3 - 6.5 - Sum of cations 2.97 6.08 7.06 6.42 6.89 2.88 3.59 - Sum of anions 3.1 6.2 7.27 6.56 6.98 2.94 3.71 - HC03- mg/l 117 314 379 353 382 149 159 - alkalinity from the cave meq/l 1.83 5.28 6.25 - 6.153561 2.585 3.446667 - HC03- mg/l 111.8 322.1 381.6 - 375.5 157.7 210.3 - Al Mg/" <20 <20 <20 <20 <20 - 155 - Mn mg/l <0.05 <0.05 <0.05 <0.05 <0.05 - - - Fe mg/l <0.1 <0.1 <0.1 <0.1 <0.1 - 0.17 - Sr Mg/l 77.6 77.9 86 75.9 87.2 77.9 80.8 - Si Mg/l 2055 1805 2225 1782 1847 4038 3787 - Li mg/l <0.1 <0.1 <0.1 <0.1 <0.1 - - - pH 7.7 7.92 7.78 7.89 7.61 8.34 8.5 8.33 drops/ Surface Surface Surface discharge min 51 39 117 167 flow flow flow 8 PC02 ppm 640 640 630 820 720 730 710 - Reconstruction of P ■Winter cave CO ♦ TC1 Winter ♦ TC1 Summer O TC2 Winter • TC2 Summer □ CP1 Winter ■ CP1 Summer A CP2 Winter A CP2 Summer O CP3 Winter ♦ CP3 Summer - ZD1 -Degassing (PHREEQC) Degassing (linear) — ~ Precipitation Trendline o 75 >-c LU 0 8 10 12 14 Fig. 4.12 Approximate locations of selected natural systems as a function of reduction-oxidation potential and pH. Note that redox units are given in terms of Eh (left) and pe. The thin line bounding the natural environments indicates the limits of nearly all natural waters (aller Bass Becking et al. I960). Taken from Ryan (2014) O 2 4 6 8 10 12 14 pH Appelo & Postma (2005) At the entrance to a wetland, sampled water has an Eh of 0.9 V and a pH of 2 At the output, the values change to Eh 0.2 V and pH 8.5 Describe, using diagrams, what happens in the wetland. How will the situation change if the decay in the wetland intensifies and Eh drops even more to, -0.2 V? LU 1.2 - 1.0-0.8 - Fg+3 X O ED k Ll_ 0.6 - 0.4- +2 Fe 0.2 - 0.0 - -0.2 - -0.4 - -0.6 - -0.8 - 1 2 Fig. 4.15 Fh pi I diagram for iron and aluminum. For Hie iron diagram, activity of Fe = 10" mol/L, S = 1()~2 mol/L and C02 - 10~3,4 aim. For aluminum diagram, Al = mol/L. Note that iron speciation is influenced by both Eh and pll; conversely, aluminum speciation is controlled by pll but not Eli. The arrow A to B represents oxidation as shown in Plate 5. Taken from Ryan (2014) REDOX CONDITIONS IN NATURAL WATERS Photosynthesis • Photosynthetic organisms convert C02to organic matter and release oxygen sunlight co2-> corg + 02 • Thanks to energy from the Sun, thermodynamically stable C02 is converted into unstable organic substances - If photosynthesis stopped, all oxygen would be consumed overtime and Corswould decompose Photosynthesis • Plants need more substances - Nitrates and phosphates - a limiting factor for algae — Trace elements - necessary, but not usually limiting • Averaging of photosynthesis to the average plankton composition: 106C02 + I6NO3 + HPO|" + 122H20 + 18H+ + Energie + trace elements = C106H2630110N16P + 13802 • The importance of phosphorus in water Respiration and decomposition processes • The opposite process of photosynthesis - when oxygen is available • The release of C02 increases the pH of the water • Without oxygen, the decomposition is a series of reactions representing gradually decreasing pe • Organic oxidation is generally: Corg + 2H20 = C02 + 4H+ + 4e~ • The main difference is in the final acceptor of e" Decomposition processes • Various oxidizing agents: A. Aerobic metabolism 02 + 4H+ + 4e" = 2H20 B. Denitrification 2NO3 + 12H+ + lOe" = N2 + 6H20 C. Fe3+reduction FeOOH + 3H+ + e~ = Fe2+ + 2H20 D. Sulphate reduction SOl" + 10H+ + Se~ = H2S + 4H20 Nitrate reduction • Bacterial decomposition: N+Vas e~ acceptor • Denitrification = reduction to biologically inert N2 Corg + 4NO3 + 4H+ = 2N2 + 5C02 + 2H20 • Other bacteria reduce to N02~or up to NH3 • Ammonia reacts with water NH3 + H20 = NH4 + 0H~ — Alkalization of the aquatic environment - Possible precipitation of CaC03 Iron reduction • In particular oxides and hydroxides with Fe3+ Corg + 4Fe(0H)3 + 8H+ = C02 + 4Fe2+ + 10H2O • Little significance in surface water - important in groundwater (order of magnitude more FeOOH) • The resulting Fe2+is transported or precipitated (eg siderite) • Very reducing conditions (reduction of sulphates) - iron sulphides are formed — Microbial interaction Sulphate reduction • Bacteria use S042~ as an oxidizing agent 2Corg + SOl" + 2H20 = H2S + 2HCO3 • Various S species may be formed - the final product is always sulphides/sulphane (according to available cations) • Bacteria can use only simple molecules (up to 20 C) — Additional fission processes producing shorter chains • H2S and HS~toxic - effect on biota Consequences of Fe3+and S042~ reduction • Fe reduction - color change from red and brown to gray and black • Release of substances sorbed on iron oxohydroxide surfaces (arsenic in Bangladesh) • Heavy metals soluble in oxidizing conditions (Cu, Zn, Mo, Pb, Hg) insoluble in the presence of S2" (forms sulphides) Fermentation and methanogenesis • Without external e" acceptors, bacteria use the transformation of organic matter as a source of energy • Complex organic substances to simpler to C02 and H2 • Other organisms use fermentation products to obtain energy producing CH4 • We will simplify both processes to general: 2Cors + 2H20 = C02 + CH4 Anaerobic decomposition • In general, the way in which microorganisms obtain energy by decomposing the products of photosynthesis • Organisms catalyze the decomposition of thermodynamically unstable substances into more stable ones • In general, each reaction has its own organisms and takes place gradually (with some overlap) • First, the most energy efficient processes: Consumption of02=> consumption of SO42~ => formation ofCH4 • 02toxic to many bacteria - must be consumed first Reductions Reduction < □enitrification Mn(4) oxide -* Mn{2) Fe(3) oxide Fe(2) Reduction Fermentation -20 Oxidat org. mat. > Sulfide -i so Oxidat of Fe (2) > Oxidations -10 Qxidat. of Mn(2) ^> ~~I— 0 pe 10 20 Figure 9.15. Sequences of important redox processes at pH = 7 in natural systems (modified and corrected after Stumm and Morgan. 1996). Buffering redox conditions • Observation: pe of groundwater decreases very slowly, then jumps to a lower level and stagnates again. • The system buffers redox if oxidizable and reducible substances are present which would prevent a significant change in pe with the addition of a small amount of strong oxidizing or reducing agent • pe remains stable while consuming a certain oxidant Buffering redox conditions In natural surface waters, 02 maintains oxidizing conditions During consumption, they fall sharply and are held by another redox pair (especially S042~, because there usually are very few nitrates) - The values of pe in water are usually within the ranges given by the buffers redox pairs - stable values Amount of organic matter reacted (units are mmol C/C H20) Buffering redox conditions • The solid phase also plays an important role in groundwater and sediments - especially oxides of 15 10 Ü Mn and Fe —I Oa H20 Mn02 Mn '2-t- FeOOH - Fe2 + SO; -+H2S er>tati0n A pH N03", MM Mft mM £ 75 8,0 0 20 40 0 10 20 20 Ü- U 40 Q Amount of organic matter reacted (arbitrary scale) 60 TTTT-| o o 0 " o o T—I—Ti T-1-1-1 1 I I Froehlich et al. (1979) Redox conditions in a lake environment Given by the balances of the decomposition of organic matter and the supply of oxygen Water circulation in lakes controlled by differences in density (consequence of temperature) Epilimnion Metalimnion (thermocline) Hypolimnion 0 5- 10h 15 20- 7h 30 T Epilimnion MetaNmnion or Thermocline zone 10 15 Temperature i°C[ Redox conditions in a lake environment When equilibrating (mixing) the 02 content is given by equilibrium with the atmosphere During stratification, the 02 content in the hypolimnion gradually decreases (aerobic decomposition of organic matter from the epilimnion) The amount of organic matter is limited by nutrients - especially phosphates Large supply of organic matter - achieving anaerobic conditions Spring turnover 02 very slow pe decrease In particular Fe(OH)3, Mn02and Fe203 Groundwater flow rate Bacterial redox processes are slow-the effect of residence time and reservoir size on the resulting pe -10 pH Natural conditions -20 15 - 10 - CD cl +5 + 1.0 + 0.5 -o -0.5 Appelo & Postma (2005) > in 0 8 10 12 14 Fig. 4.12 Approximate locations oi' selected natural systems as a function of reduction-oxidation potential and pH. Note that redox units are given in terms of Eh (left) and pe. The thin line bounding the natural environments indicates the limits of nearly all natural waters (aller Bass Becking et al. I960). Taken from Ryan (2014) Summary • Redox conditions of natural waters are usually determined by the balance of 02 supply from the atmosphere and its consumption by microbial decomposition of organic matter. • After 02 consumption , pe decreases progressively with the advancing sequence of reactions (reduction of Fe3+, S042_and fermentation). • Organic production in lakes and the ocean is significantly controlled by the availability of inorganic nutrients (phosphates and nitrates). • Changes in pe have a great influence on the solubility and mobility of many metals. EVROPSKÁ UNIE Evropské strukturální a investiční fondy Operační program Výzkum, vývoj a vzdělávání MINISTERSTVO ŠKOLSTVÍ, MLÁDEŽE A TELOVÝCHOVY Tento učební materiál vznikl v rámci projektu Rozvoj doktorského studia chemie č. CZ.02.2.69/0.0/0.0/16_018/0002593 References • APPELO, CAJ and Dieke POSTMA. Geochemistry, groundwater and pollution :. _ 2nd ed . Leiden: AA Balkema publishers , c2005. ISBN 0-415-36421-3. • CLARK, I. (2015): Groundwater Geochemistry and Isotopes.BocaRaton , Florida: CRC Press. • DREVER, James I. The geochemistry of natural waters: surface and groundwater environments . 3rd ed. Upper Saddle River, NJ: Prentice Hall, cl997. ISBN 0-13-272790-0. • Ryan, P. (2014). Environmental and low temperature geochemistry. John Wiley and Sons. 402p. ISBN 978-1-4051-8612-4 ( pbk .) • Froelich , PN, et al. "Early oxidation of organic matter in pelagic sediments of the eastern equatorial Atlantic : suboxic diagenesis ." Geochimica et Cosmochimica Acta 43.7 (1979): 1075-1090. References • Unmarked images are author's or taken over with the permission of doc. Zemana and doc. Faimona • Clark, I. (2015). Groundwater Geochemistry and Isotopes . CRC Press . 442p. ISBN 978-1-4665-9174-5 ( eBook - PDF) • Drever, J. (1997): The geochemistry of natural waters : surface and groundwater environments . (I have a book from our library with me on a long-term loan) • Dreybrodt, W. & Eisenlohr, L. (2000): Limestone dissolution rates in karst environments . - In: Klimchouk, AB ( ed .): Speleogenesis evolution of karst aquifers . - National speleological society. Huntsville , Alabama. 527 s. • Holland HD, Kirsipu T V., Huebner JS, Oxburgh UM (1964) On Some Aspects of the Chemical Evolution of Cave Waters. J Geol 72: 36-67. doi : 10.1086 / 626964 • Plummer LN, Parkhurst DL, Wigley TML (1979) Critical review of the kinetic of calcite dissolution and precipitation. • Plummer LN, Wigley TML, Parkhurst DL (1978) The kinetics of calcite dissolution in CO 2 -water systems at 5 degrees to 60 degrees C and 0.0 to 1.0 atm CO 2. Am J Sei 278: 179-216. doi: 10.2475 / ajs.278.2.179 • Pracný P, Faimon J, Všianský D, Handbag L (2017) Evolution of Mg / Ca Ratios During Limestone Dissolution Under Epicarstic Conditions. Aquat Geochemistry 1-21. doi : 10.1007 / sl0498-017-9313-y • Pracný P, Faimon J, Kabelka L, Hebelka J (2016) Variations of carbon dioxide in the air and dripwaters of Punkva Caves (Moravian Karst, Czech Republic). Carbonates and Evaporites 31: 375-386. doi : 10.1007 / S13146-015-0259-0 • Stumm , W. and Morgan, J. (1995): Aquatic chemistry: chemical equilibrium and rates in natural waters.