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.

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?

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)

Reactants include Water and… Inorganic solids: rocks, sediments,

minerals Biota: plants, animals, bacteria, archea Gas phases: Earth’s atmosphere,

biologically produced gas (CO2, H2S)

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

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

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

Chlorinity (Cl)

Definition 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 (19.374 g/kg) Commonly used in seawater systems

Salinity Operationally – all salts in seawater Originally (Knudsen, 1901) defined as

With Cl = 19.374 g/kg measured with AgNO3

Intercept because not all salts measured Precision to 3 decimal places

S(‰)=0.03 + 1.805 Cl(‰)

From Millero, 2011, Marine Chemistry

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(‰)=1.80655 Cl(‰)

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)

S P=( 35.00035.165 )∗( S Rg /kg

)

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)

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)

S Ag /kg

=SR+δ S

Other measures of TDS

Refractive index 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

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

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

Weight units

Mass per unit volume 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

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 SO4

2-

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).

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 betweenmass (easily measured) and moles

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

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

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

Calculation: Moles (or millimoles) of ion times its

chargeNa2SO4 = 2Na+ + SO4

2-

1 mole of Na = 1 eq Na solution1 mole of SO4 = 2 eq SO4 solution

Although different number of moles, solution is still electrically neutral

Example

Charge Balance Electrical neutrality provides good

check on analytical error Charge Balance Error – CBE

Note: - = excess anions; + = excess cations

CBE =SmcZc - SmaZa

SmcZc + SmaZaWhere: m = molar concentration of major solutes

z =charge of cation (c) or anion (a)

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

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

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

Zaporozec, 1972, GW

Depth

(ft

)

Piper Diagrams

Two ternary diagrams

Projected on quadralinear diagram

Very useful figure for comparing concentrations of numerous water samples

Construction

Convert concentrations to meq/L Use major cations and anions

concentrations Cations = Ca, Mg, Na + K Anions = SO4, Cl, HCO3 + CO3 (or

alkalinity)

Calculate %’s of each element on ternary diagram

For example Ca is:

Plot %’s on ternary diagrams Project each % onto diamond

diagrams

[Ca]

[Ca] + [Mg] + [Na + K]*100

Composition:Ca = 22.3%Mg = 13.7%Na+K = 64%

Alkalinity = 31.3%Sulfate = 54.5%Chloride = 14.2%

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

Fortunately – computer programs available to make plots for you Stiff diagrams:

http://www.twdb.state.tx.us/publications/reports/GroundWaterReports/Open-File/Open-File_01-001.htm

Piper plots: http//water.usgs.gov/nrp/gwsoftware/GW_Chart/GW_Chart.html

Geochemist workbench http://student.gwb.com/ Download GWB