1. Groundwater Chemistry Evolution
Evolutionary sequence controlled by mineral identity, availability, and solubility
High availability: carbonates and felsic minerals
High solubility: gypsum/anhydrite, evaporites
Deeper groundwater is a closed system with respect to gases
Water is isolated from the atmosphere
If gases are consumed, their concentrations decrease; if generated, concentrations increase

2. Evolution of Groundwater Chemistry

3. Trends with age/depth
As groundwater migrates, concentration of TDS and most major ions increases
Anions
HCO3-  HCO3- + SO42-  SO42- + HCO3-  SO42- + Cl-  Cl- + SO42-  Cl-
Cations
More difficult to generalize trends
Most common trend: Ca2+, then Ca-Na, Na-Ca, finally Na+
Driven by cation exchange and CaCO3 precipitation

4. Evolution of Groundwater Chemistry
Low TDS
Intermediate TDS
Aquitard: TDS high relative to aquifers
High TDS

5. Water Chemistry: Information on Weathering Reactions
Knowing starting and ending solution chemistry of a system, we can infer what reactions have taken place to produce the ending solution
Reaction-Path Modeling
In addition to water chemistry, need information on minerals present
As groundwater migrates along a flow path, reactions occur:
Dissolution adds ions
Mineral precipitation removes ions
The change in water chemistry = the sum of all dissolution/precipitation reactions

6. Reaction Path Models
Good for simple systems where flowpaths are well defined
The larger and more complex the systems, the harder it is to constrain potential reactions
Can consider redox reactions, gas exchange, isotopic reactions, mixing of waters, etc.
N.B.: there is no unique solution
Modeler determines which phases to consider
Based on available data and “intuition”

7. Redox reactions in Groundwater
Redox reactions are extremely important in groundwater and soil water
Water tends to become more reducing as it moves along a flow path
Almost all redox reactions in groundwater are biogeochemically mediated
DO typically consumed in the soil zone and shallow groundwater, resulting in anoxic groundwater

8. Groundwater Chemistry: Redox Evolution
After DO is consumed, other TEAPs are used by microbes based on thermodynamics
NO3- reduction (denitrification)
MnO2 [Mn(IV)] reduction
Ferric [Fe(III)] mineral reduction
SO42- reduction
Fermentation and methanogenesis (CO2 reduction)
“Redox ladder”
The order of the reactions based on obtainable energy for the microbes
Kinetics: the less the energy, the slower the reaction

9. Organic Matter Oxidation
Aerobic
CH2O + O2  CO2 + H2O
Denitrification
5 CH2O + 4 NO H+  5 CO2 + 2 N2 + 7 H2O
Ferric iron [Fe(III)] reduction
CH2O + 4Fe(OH)3 + 8 H+  CO2 + 4 Fe H2O
Sulfate reduction
2CH2O + SO42- + H+  2 CO2 + HS- + 2 H2O

10. Redox Ladder: electron acceptors and donors

11. Fermentation and Methanogenesis
Reactions that occur when all external electron acceptors have been used; methane (CH4) is produced, CO2 both produced and consumed
Transformation of complex organics into simpler compounds
Fermentation:
CH3COOH  CH4 + CO2
CH3COOH = Acetic acid
Also produces H2
CO2 + H2O  HCO3- + H+
2 H+ + 2 e-  H2
Fermentation byproducts are used by methanogenic microbes

12. Fermentation and Methanogenesis
Methanogenesis: CO2 + 4 H2  CH4 + 2 H2O
Methanogens need fermenters
H2 is a reactive intermediate product, produced and consumed by metabolic processes
Low at high Eh, higher at lower Eh
H2 is best indicator of dominant TEAP, but difficult to measure (field GC)

13. Use of H2 to delineate redox processes
Chapelle et al. (1996)
Hypothesis:
Fermentative microbes continually produce H2
Fe(III), SO42-, and CO2 reducing microbes use H2 as TEAP at different efficiencies
Fe(III) reduction: 0.2 – 0.8 nM
SO42- reduction: 1 – 4 nM
CO2 reduction (methanogenesis): 5 – 15 nM

14. Redox predictions based on Eh

15. Redox predictions based on H2

16. Other Redox Data

17. TEAPs in Groundwater
Contaminated
Uncontaminated
FLOW

18. Mason Tree Nursery

19. Mason Co. Tree Nursery

20. Mason Co. Tree Nursery

22. TEAPs
While thermodynamics predicts an orderly progression of the dominance of individual TEAPs, it’s not so simple in nature
Often have 2 (or more) TEAPs active in same part of aquifer
e.g., often have Fe(III)-reduction and SO42--reduction occurring together, even though Fe(III)-reduction more thermodynamically favorable
Due to: micro-environments, different microorganisms responsible, solid vs. aqueous environments
Where there’s energy to be gained, microbes are working

23. TEAPs and Eh Ranges

24. Determining predominant TEAP

25. Defining Redox Zones
From McMahon, P.B. and F.H. Chapelle Redox processes and water quality of selected principal aquifer systems. Ground Water 46(2):

26. McMahon, P. B. , and F. H. Chapelle. 2007
McMahon, P.B., and F.H. Chapelle Redox Processes and Water Quality of Selected Principal Aquifer Systems. Ground Water 46:259–271

27. Principal TEAPs in U.S. Aquifers

28. Principal TEAPs in U.S. Aquifers

29. Redox Conditions in Aquifers
Shallow groundwater usually low but detectable DO
Most deeper aquifers are anoxic
Key variables:
Organic matter
Hydraulic conductivity
Mineralogy
Recharge rates (climate)
Most aquifers have a dominant TEAP, but most (if not all) TEAPs active

30. Redox Buffering
Observation: the Eh of groundwater does not linearly decline as oxidizers are consumed along a flow path
The Eh remains relatively constant as a particular oxidizer is consumed, then the Eh drops and stabilizes again
Similar to pH buffering in that a reaction is preventing a rapid change even though e- (vs. H+) are being produced/consumed
For pH buffers, occurs around pKa of conjugate acid/base pair (e.g., 6.35 for H2CO3/HCO3-)
For Eh buffers, occurs around E°

31. Redox Buffering
System is buffered if oxidizable or reducible compounds are present that prevent a significant change in Eh when strong oxidizing/reducing agents added
Expect Eh of natural waters to generally be in buffered ranges
Values in unbuffered ranges unstable

32. Redox Buffering

33. Computed vs. Measured Field Eh
- Vertical bands indicate buffered ranges; reflect the standard E°

34. Redox Buffering
Example: recharging water has dissolved O2, Eh will remain high until O2 is consumed; after O2 gone, Eh drops rapidly and stabilizes at the value determined by next oxidizer
Buffers can be dissolved species or solid matter
Dissolved species: usually limited in concentration and consumed rapidly (if right conditions exist)
Solid matter: can provide large buffering capacity
e.g., Fe(OH)3 can provide buffering until equilibrium is reached with dissolved Fe concentration

35. Evaluating Water Chemistry Data

36. Evaluating Water Chemistry Data
When we collect a sample, we trust that the lab analyzes and reports the results correctly
Need to do appropriate field and lab QA/QC (more on this later)
One test of analytical integrity is the charge balance error (CBE)

37. Charge Balance Error (CBE)
Based on the concept all ions in water are charge balanced, i.e., Σanions = Σcations
Calculate using equivalents (molarity x charge)
Can use all ions, but often only major ions considered
A positive value = excess of cations
A negative value = excess of anions
A value < 5% is usually considered adequate

38. Charge Balance Errors Ca Mg Na HCO3 Cl SO4 CBE 79.2 31.2 25.4 413 0.718
3.2
60.2
34.5
36.2
450
2.0
1.5
-0.3
31.5
15.1
399
1.0 26
-36.8
76.7
30.0
18.1
421 22
-2.1
42.2
17.2
10.3
434
0.9
5.8
-29.3
83.9
37.4
27.1
456
1.1 28
2.1
72.0
32.0
18.0
440
0.85
-1.8
When you get a water sample composition, the first thing you should do
is calculate the CBE (assuming it’s a complete analysis)

39. Graphical Data Analysis
Graphs are essential for two purposes:
To provide insight into the data under scrutiny
To illustrate important concepts when presenting results
Graphing should be done before any other analysis
See patterns
Guide further analysis

40. Graphs useful for Water Quality Data
Histograms
Scatterplots
Box and whisker plots
Piper diagrams
Stiff diagrams

41. Histograms
Bars drawn to indicate number of samples in a certain interval
Visual impression depends on number of intervals

42. Histograms
For sample size of n, number of intervals (k) should be the smallest integer for 2k ≥ n
So for n = 100, k = 7 (27 = 128)

43. Histograms
Not best for data measured on continuous scale (such as concentration)
Best when displaying data which have natural categories or groupings
e.g., number of wells contaminated with bacteria based on land use or rock type

44. Scatterplots (x-y plots)
Very common, easily made
Illustrates the relationships between 2 (or more) variables
Can perform linear regressions

45. Scatterplot

46. Scatterplot

47. Boxplots Boxplots provide a visual summary of:
The center of the data (the median – the center line of the box)
The variation or spread (interquartile range – the box height)
The skewness (quartile skew – the relative size of box halves)
Presence or absence of unusual values ("outside" and "far outside" values)
Can easily compare more than one dataset

48. Box Plots 10
Concentration
IQR

49. Box Plots

50. Making Box Plots: calcium data
30.7
58.7
73.7
77.2
82.2
32.5
60.2
73.9
77.4
82.4
33.5
60.3
78.4
83.9
40.1
60.9
74.0
79.0
42.2
62.5
74.2
79.2
42.3
65.2
74.4
84.3
45.8
65.9
74.9
79.3
85.5
51.8
71.5
75.6
89.8
52.0
72.0
76.7
79.4
93.2
52.6
76.9
80.5
93.6
55.9
73.5
77.0
80.7
120.0
77.0
79.4
72.0
73.5
82.2
89.8
74.4
60.2
83.9
93.6
65.9
79.2
74.9
76.9
71.5
80.5
73.9
79.0
75.6
65.2
42.3
73.7
80.7
77.2
58.7
51.8
60.9
85.5
40.1
55.9
33.5
45.8
82.4
79.3
62.5
52.6
42.2
93.2
76.7
120.0
52.0
77.4
30.7
74.2
74.0
84.3
32.5
78.4
60.3
Outliers
Outliers
10th
25th
Median
75th
90th
Whiskers
Box

51. Box Plot
84.1
79.35
74.4 (median)
60.6
44.05

52. Plots Specific to Groundwater Chemistry
Piper diagrams
Stiff diagrams

53. Piper Diagrams Shows relative ratios of major ions Uses Pros and cons
Visually describes the differences in major ion chemistry in groundwater flow systems
Can indicate mixing between 2 water sources
Pros and cons
Can show large number of samples
Do not show ion concentrations
Symbol size sometimes made proportional to TDS or some other ion/species

54. Piper Diagram

55. Constructing Piper Diagrams
Need to have concentrations for all major ions
Basically start as with CBE
Convert all concentrations to meq/L
Separately sum all the cations and anions
Sum (Na+ + K+) and (HCO3- + CO32-) (if measured)
CO32- only important at pH > 8.3 or so, so in groundwater can usually ignore
Divide Ca2+, Mg2+, and (Na+ + K+) by Σcations
Divide HCO3- , Cl-, and SO42- by Σanions
Multiply these values by 100
Plot these values on the 2 trilinear diagrams

56. Constructing Piper Diagrams
Now combine (Ca2+ + Mg2+) and (SO42- + Cl-)
Divide these values by Σcations or Σanions, as appropriate
Multiply by 100
(Ca2+ + Mg2+) = (Na+ + K+)
(SO42- + Cl-) = (HCO3- + CO32-)
Plot these values and (HCO3- + CO32-) and (Na+ + K+) values on middle diagram
Excel can do these calculations easily

57. Stiff Diagrams Shows absolute values of major ions Pros and cons
Easy to see differences/similarities between samples
Widths show ion concentrations
Can add additional ions as desired
Difficult to show a lot of samples

58. Stiff Diagrams

59. Stiff Diagrams

60. Constructing Stiff Diagrams
Again, basically start as with CBE
Convert all concentrations to meq/L
Sum (Na+ + K+)
Plot

61. Plots to avoid
Stacked bar charts
Pie diagrams

62. Solute Transport
Ions and molecules being transported in the subsurface often travel at rates slower than water
The migration is “retarded” primarily due to their interactions with mineral surfaces
Adsorption and ion exchange are driven by electrical interactions

63. Surface Charge Solids typically have an electrically charged surface
There are 2 main sources of surface charge
(1) Chemical reactions
At low pH, surfaces tend to have positive charges, at high pH, negative charges
Where the net surface charge = 0 is the isoelectric point, or point of zero charge
This is a function of the solid identity

64. Isoelectric Points
Phase
iso. pt.
SiO2
2.0
MnO2
2.8
Kaolinite
4.6
Fe2O3
6.7
FeOOH
7.8
Fe(OH)3(am)
8.5
Al2O3
9.1
MgO
12.4
Below this point, charge is positive, above is negative
For most common solid phases at natural pHs, the surface charge is negative

65. Surface Charge
The second main source of surface charge is lattice imperfections and substitutions in the solid
e.g., Al3+ commonly substitutes for Si4+ and Mg2+ for Al3+ in clay layers
Charges resulting from these are not pH dependent

66. Surface Charge
As with other systems, the interfacial system (surface – water) must be electrically neutral
Electrical Double Layer
Fixed surface charge on the solid
Charge distributed diffusely in solution
Excess of counterions (opposite charge to surface) and deficiency of ions of same charge as surface
Counterions attracted to the surface

67. Adsorption
Adsorption refers to a dissolved ion or molecule binding to a charged surface
Reversible reactions; i.e., if conditions change, the ion can desorb
An important process for removing some ions, such as heavy metals, from solution
Heavy metals are typically present in groundwater in much lower concentrations than predicted by solubility calculations

68. Adsorption
Counterions
Fixed Surface
Charge

69. Ion Exchange
Ion exchange refers to exchange of ions between solution and solid surfaces
It differs from adsorption in that an ion is released from the surface as another is adsorbed
AX + B+  BX + A+
X refers to a mineral surface to which an ion has adsorbed
Most important for cations, anions less so, because most mineral surfaces are negatively charged
Primarily occurs on clay minerals of colloidal size (10-3 – 10-6 mm)

70. Cation exchange
This is basically how water softeners work

71. Cation exchange

72. Ion Exchange Ion size (radius) and charge affect how they exchange
Smaller ions from stronger bonds on surfaces
e.g., Ba2+ is smaller than Mg2+, more likely to be adsorbed to surface
Ions with more positive charge form stronger bonds on surfaces
e.g., Ca2+ more likely to be adsorbed, Na+ more likely to go into solution
Reversible reactions
Cation exchange often cited as reason for evolution of Ca  Na dominant waters