1. Hydrogeochemistry “Geochemistry of Natural Waters”
No wastewater, water resources
Class not related directly to social aspects of water
Study natural controls of chemistry of rivers, lakes, ground water, oceans etc.

2. Questions considered:
Why do different waters have different chemical compositions?
What controls the compositions?
This will lead to discussion of chemical reactions between water, rocks, and atmosphere
How do compositions vary with setting?
How do they vary with time?

3. Why consider water compositions?
Can provide information on basic geological processes:
Diagenesis ≡ All chemical (and physical) alteration of solid material (low T and P)
Weathering ≡ alteration of mineral phases at earth surface conditions
Biological activities in certain environments
Exchange with atmosphere (C cycle – global climate change)

4. Reactants include Water and…
Inorganic solids: rocks, sediments, minerals
Biota: plants, animals, bacteria, archea
Gas phases: Earth’s atmosphere, biologically produced gas (CO2, H2S)

5. Concentration terminology
We will need a way to discuss what is in water
Dissolved components (ions, gasses, complexes etc.)
Solid components
The following is a lot of definitions, terminology, and algorithms to calculate solutes

6. Dissolved concentrations
Fresh water
Potable, generally < 1000 mg/L solids per liter of water (TDS)
Brackish
Non-potable, but < seawater
Seawater, salinity 34 to 37 (defined soon)
97% of free water on earth
Concentration important threshold for Thermo
Saline water/brine > seawater salinity

7. Total dissolve solids
Mass of solid remaining after evaporation of water
Note – in GWB, often just “dissolved solids”
Bicarbonate converted to carbonate
Units of mass/volume (e.g. g/L, kg/L, etc.)
Commonly used in fresh water systems

8. Chlorinity (Cl) Definition Determined by titration with AgNO3
Mass of Cl in one kg of seawater equivalent to Cl, Br, and I in seawater
Determined by titration with AgNO3
Precipitate Ag(Cl,Br,I) with indicator
AgNO3 calibrated with seawater with known chlorinity ( g/kg)
Commonly used in seawater systems

9. Salinity Operationally – all salts in seawater
Originally (Knudsen, 1901) defined as
With Cl = g/kg measured with AgNO3
Intercept because not all salts measured
Precision to 3 decimal places
S(‰)= Cl(‰)
From Millero, 2011, Marine Chemistry

10. Salinity In 1960’s, electrical conductivity of seawater of known Cl
More precise, easier to measure
Conversion adopted
More precise, dropped intercept value
Referred to now as “Reference Salinity” SR
Units of g/kg or ‰
S(‰)= Cl(‰)

11. Practical Salinity Scale
Based on conductivity measurements to determine Reference Salinity (SR)
Practical Salinity (SP) is defined as
Since ratio, it is unitless
Originally had units of PSU (Practical Salinity Units)
SP= ∗( SR g/kg )

12. Absolute Salinity Scale (SA)
Defined by
SCOR (Scientific Committee on Oceanic Research)
IAPSO (International Association of Physical Sciences of the Ocean)
Endorsed by IOC (Intergovernmental oceanographic Commission)

13. Defined as:
where dS represents addition in deep ocean water from dissolution of minerals (e.g., calcite, diatoms etc) and organic carbon (e.g., non-conservative solutes)
SA g/kg =SR+δS

14. Other measures of TDS Refractive index Conductivity/resistivity
Amount of refraction of light passing through water
Linearly related to concentrations of dissolved salts
Conductivity/resistivity
Current carried by solution is proportional to dissolved ions

15. Conductivity Inverse of resistance Units of Siemens/cm
Siemen = unit of electrical conductance
1 Siemen = 1 Amp/volt = 1/Ohm = 1 Mho
Conductance is T dependent
Typically normalized to 25º C
Called Specific Conductivity

16. Reporting units
Need to report how much dissolved material (solute) in water, two ways:
Mass
Moles (& equivalents)
Need to report how much water (solvent)
Volume of water, typically solution amount – analytically easy
Mass of water, typically solvent amount – analytically difficult

17. Weight units Mass per unit volume If very dilute solution
For example: g/L or mg/L
If very dilute solution
Mass per unit volume about the same as mass per mass
1L water ~ 1000 g, BUT variable with T, P and X as density changes

18. 1 mole sodium sulfate makes 2 moles Na+ and 1 mole SO42-
Molar units
Number of molecules (atoms, ions etc) in one liter of solution
Most common – easy to measure solution volumes
Units are M, mM, µM (big M)
Example
Na2SO4 = 2Na+ + SO42-
1 mole sodium sulfate makes 2 moles Na+ and 1 mole SO42-

19. Molal Units Number of molecules (atoms, ions etc) in one kg of solvent
Abbreviation: m or mm or mm (little m)
Difficult to determine mass of solvent in natural waters with dissolved components
not used so often in natural waters
Useful for physical chemistry because doesn’t depend on T, P or X (composition of dissolved constituents).

20. Simple to convert between
Why use molar units?
Reaction stoichiometry is written in terms of moles, not mass
CaCO3 = Ca2+ + CO32-
100 g =
1 mole
40 g =
1 mole
60 g =
1 mole
Simple to convert between
mass (easily measured) and moles

21. Example Nitrate a pollution of concern Commonly measured as mass
Reported as mass of N in NO3
E.g., g N in NO3
NO3 is measured
Weight concentration of NO3 is 4 X bigger than weight concentration of N
Molar units of NO3 and N are identical

22. Mass – Mole conversion Conversion from weight units to molar units
Divide by gram formula weight
Molar units to weight units
Multiply by gram formula weight

23. Alternative – charge units
Equivalents – molar number of charges per volume
eq/L or meq/L
Used to plot piper diagrams
Used to calculate electrical neutrality of solutions

24. Na2SO4 = 2Na+ + SO42- Calculation: 1 mole of Na = 1 eq Na solution
Moles (or millimoles) of ion times its charge
Na2SO4 = 2Na+ + SO42-
1 mole of Na = 1 eq Na solution
1 mole of SO4 = 2 eq SO4 solution
Although different number of moles, solution is still electrically neutral

25. Example

26. Charge Balance SmcZc - SmaZa CBE = SmcZc + SmaZa
Electrical neutrality provides good check on analytical error
Charge Balance Error – CBE
Note: - = excess anions; + = excess cations
SmcZc - SmaZa
CBE =
SmcZc + SmaZa
Where: m = molar concentration of major solutes
z =charge of cation (c) or anion (a)

27. Possible causes of errors
Significant component not measured
Commonly alkalinity – can be estimated by charge balance
Analytical error
+5% difference OK – acceptable
+ 3% good
0% probably impossible

28. Graphical data presentation
Cross plots (XY plots)
Time series
Ternary diagrams
Stiff diagram – geographic distribution
Piper diagram – comparison of large numbers of samples
Many others, not used widely

29. Stiff Diagrams Cations on left Anions on right
Top K+Na & Cl = seawater
Middle Ca & HCO3 (alkalinity) most fresh water
Carbonate mineral reactions
Bottom Mg & SO4 – other major component
4th optional – redox couples
Drever, 1997
Depth (ft)
Zaporozec, 1972, GW

30. Piper Diagrams Two ternary diagrams Projected on quadralinear diagram
Very useful figure for comparing concentrations of numerous water samples

31. Construction Convert concentrations to meq/L
Use major cations and anions concentrations
Cations = Ca, Mg, Na + K
Anions = SO4, Cl, HCO3 + CO3 (or alkalinity)

32. Calculate %’s of each element on ternary diagram For example Ca is:
Plot %’s on ternary diagrams
Project each % onto diamond diagrams
[Ca]
*100
[Ca] + [Mg] + [Na + K]

33. Composition:
Ca = 22.3%
Mg = 13.7%
Na+K = 64%
Alkalinity = 31.3%
Sulfate = 54.5%
Chloride = 14.2%

34. Santa Fe water chemistry
Mixing between three sources of water:
1 = shallow groundwater
2 = deep groundwater
3 = dilute Surface water
Moore et al., 2009, J Hydro

35. Fortunately – computer programs available to make plots for you
Stiff diagrams:
Piper plots: http//water.usgs.gov/nrp/gwsoftware/GW_Chart/GW_Chart.html
Geochemist workbench
Download GWB