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
1
Evolution of Groundwater Chemistry
2
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
3
Evolution of Groundwater Chemistry
Low TDS
Intermediate TDS
Aquitard: TDS high relative to aquifers
High TDS
4
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
5
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”
6
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
7
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
8
Organic Matter Oxidation
• Aerobic
– CH2O + O2  CO2 + H2O
• Denitrification
– 5 CH2O + 4 NO3- + 4 H+  5 CO2 + 2 N2 + 7 H2O
• Ferric iron [Fe(III)] reduction
– CH2O + 4Fe(OH)3 + 8 H+  CO2 + 4 Fe2+ + 11 H2O
• Sulfate reduction
– 2CH2O + SO42- + H+  2 CO2 + HS- + 2 H2O
9
Redox Ladder: electron acceptors and
donors
10
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
11
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)
12
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
13
Redox predictions based on Eh
14
Redox predictions based on H2
15
Other Redox Data
16
TEAPs in Groundwater
Contaminated
Uncontaminated
FLOW
17
Fertilized fields
Tree Nursery
+
a.
Road
Mason Tree Nursery
Gradient
Mason Co. Tree Nursery
0
10
Depth (ft)
20
30
40
50
MLS-3
MLS-18
60
70
0
2
4
6
NO3-N (mg/L)
8
10
12
Mason Co. Tree Nursery
10
DO
Depth (ft)
20
NO3-N
30
40
Fe
50
0
2
4
6
8
10
NO3-N and DO (mg/L)
0.0
0.2
0.4
Fe (mg/L)
0.6
0.8
Fertilized fields
Road
a.
Tree Nursery
+
Gradient
b.
+
c.
+
MLS
d.
MLS
MLS
+
Reducing zone
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
22
TEAPs and Eh Ranges
23
Determining
predominant
TEAP
24
Defining Redox Zones
From McMahon, P.B. and F.H. Chapelle. 2008. Redox processes and water quality of
selected principal aquifer systems. Ground Water 46(2):259-71.
25
McMahon, P.B., and F.H. Chapelle.
2007. Redox Processes and Water
Quality of Selected Principal
Aquifer Systems. Ground Water
46:259–271
26
Principal TEAPs in U.S. Aquifers
27
Principal TEAPs in U.S. Aquifers
28
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
29
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°
30
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
31
Redox Buffering
32
Computed vs. Measured Field Eh
- Vertical bands indicate
buffered ranges; reflect
the standard E°
33
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
34
Evaluating Water Chemistry Data
35
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)
36
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
37
Charge Balance Errors
Ca
79.2
60.2
0.7
76.7
42.2
83.9
72.0
Mg
31.2
34.5
31.5
30.0
17.2
37.4
32.0
Na
HCO3
25.4
413
36.2
450
15.1
399
18.1
421
10.3
434
27.1
456
18.0
440
Cl
0.7
2.0
1.0
1.0
0.9
1.1
0.9
SO4
18
1.5
26
22
5.8
28
0.85
CBE
3.2
-0.3
-36.8
-2.1
-29.3
2.1
-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)
38
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
39
Graphs useful for Water Quality Data
•
•
•
•
•
Histograms
Scatterplots
Box and whisker plots
Piper diagrams
Stiff diagrams
40
Histograms
• Bars drawn to indicate number of samples in a
certain interval
– Visual impression depends on number of intervals
41
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)
42
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
43
Scatterplots (x-y plots)
• Very common, easily made
• Illustrates the relationships between 2 (or
more) variables
• Can perform linear regressions
44
Scatterplot
1400
TDS (mg/L)
1200
1000
800
600
400
200
400
800
1200
1600
2000
2400
Specific Conductance (S/cm)
45
Scatterplot
300
Tazewell Co.
Champaign Co.
Central Valley
Arsenic (g/L)
250
200
150
100
50
0
0
50
100 150 200 250 300 350
SO42- (mg/L)
46
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
47
Box Plots
IQR
Concentration
10
Outlier
90th percentile
75th percentile
Median
25th percentile
10th percentile
0
48
Box Plots
Concentration (mg/L or g/L) or mv
1000
TOC < 2 mg/L
_ 2 mg/L
TOC >
100
10
1
DL
DL
DL
0.1
DL
0.01
0.001
As
Fe
Mn
NH4-N
SO42-
-
HCO3
ORP
49
77.0
30.7
79.4
58.7
72.0
73.7
73.5
77.2
82.2
89.8
32.5
72.0
60.2
74.4
73.9
60.2
77.4
83.9
82.4
93.6
33.5
65.9
60.3
79.2
73.9
74.9
78.4
83.9
76.9
40.1
71.5
60.9
80.5
74.0
73.9
79.0
79.0
83.9
75.6
42.2
65.2
62.5
42.3
74.2
73.7
79.2
80.7
83.9
77.2
42.3
58.7
65.2
51.8
74.4
60.9
79.2
85.5
84.3
40.1
45.8
55.9
65.9
33.5
74.9
45.8
79.3
82.4
85.5
79.3
51.8
62.5
71.5
52.6
75.6
42.2
79.3
93.2
89.8
76.7
52.0
120.0
72.0
52.0
76.7
77.4
79.4
73.9
93.2
79.2
52.6
79.3
72.0
30.7
76.9
74.2
80.5
74.0
93.6
84.3
55.9
83.9
73.5
32.5
77.0
78.4
80.7
120.0
60.3
10th
25th
Median
Whiskers
Box
75th
Outliers
Outliers
Making Box Plots: calcium data
90th
50
Box Plot
84.1
79.35
74.4 (median)
60.6
44.05
51
Plots Specific to Groundwater
Chemistry
• Piper diagrams
• Stiff diagrams
52
Piper Diagrams
• Shows relative ratios of major ions
• Uses
– 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
53
Ca
+C
l
Piper Diagram
SO
4
g
+M
SO4
Mg
100
0
0
CO
40
60
60
80
20
40
80
20
100 100
0
100
Ca
3
+K
60
40
80
+H
Na
40
60
20
CO
3
20
80
100
80
60
40
20
0
Na+K
0
0
20
HCO3 + CO3
40
60
80
100
Cl
54
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
55
Constructing Piper Diagrams
• Now combine (Ca2+ + Mg2+) and (SO42- + Cl-)
• Divide these values by Σcations or Σanions, as
appropriate
– Multiply by 100
– (Ca2+ + Mg2+) = 100 - (Na+ + K+)
– (SO42- + Cl-) = 100 - (HCO3- + CO32-)
• Plot these values and (HCO3- + CO32-) and (Na+
+ K+) values on middle diagram
• Excel can do these calculations easily
56
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
57
Stiff Diagrams
58
Stiff Diagrams
Group 2
Group 1
Group 3
Sample 10
Sample 14
Sample 1
-15 -10
-5
0
5
10
15
-15 -10
-5
meq/L
0
5
10
15
-15 -10
-5
meq/L
0
5
10
15
5
10
15
meq/L
Sample 15
Sample 11
Sample 2
1.0000 vs Col 1
-15 -10
-5
0
Col 2 vs Col 1
-15 -10
-5
0
5
10
15
-15 -10
-5
meq/L
0
5
10
meq/L
15
meq/L
Sample 16
Sample 12
Sample 3
-15 -10
-5
0
5
10
15
meq/L
-15 -10
-5
0
5
10
15
-15 -10
-5
meq/L
0
5
10
15
meq/L
Sample 17
-15 -10
Sample 13
-5
-15 -10
0
5
10
15
5
10
15
meq/L
Sample 4
-5
0
meq/L
5
10
15
-15 -10
-5
0
5
10
15
Sample 18
meq/L
-15 -10
-5
0
meq/L
59
Constructing Stiff Diagrams
• Again, basically start as with CBE
– Convert all concentrations to meq/L
– Sum (Na+ + K+)
– Plot
60
Plots to avoid
• Stacked bar charts
• Pie diagrams
6
4
mg/L
C
Al
B
Ba
Li
5
D
B
E
F
3
2
A
1
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
61
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
62
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
63
Isoelectric Points
Phase
SiO2
MnO2
Kaolinite
Fe2O3
FeOOH
Fe(OH)3(am)
Al2O3
MgO
iso. pt.
2.0
2.8
4.6
6.7
7.8
8.5
9.1
12.4
• Below this point,
charge is positive,
above is negative
• For most common
solid phases at
natural pHs, the
surface charge is
negative
64
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
65
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
66
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
– http://www.cee.vt.edu/ewr/environmental/teach
/gwprimer/sorp/sorp.html
67
_
_
_
_
Solid _
_
Phase _
_
Fixed Surface _
Charge
_
Counterions
Adsorption
_
+
+
+
_
_
_
+
Solution
+
_
68
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)
69
Cation exchange
This is basically how water softeners work
70
Cation exchange
71
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
72