Download Basic Background Review: Acid-Base , Redox, and Stable Isotopes

Basic Background Review:
Acid‐Base , Redox, and Stable Isotopes
1) Basic Physical Chemistry
• Thermodynamics
• Potential for reaction to proceed
– Expected state at equilibrium
• Kinetics
– Reaction rates
– Reaction inhibition, catalysis
What makes reactions go?
• Maximum entropy, minimum free energy (= equilibrium state of system)
•Gibbs Free Energy for a chemical substance: G=H‐TS
•ΔH largely function of strengths of chemical bonds in products vs, reactant (more strong bonds in product than reactants (negative ΔH) contributes to spontaneous reaction)
•ΔS measure of relative degree of disorder in product and reactant—greater disorder in product than reactant (positive ΔS) contributes to spontaneous reaction
`
What makes reactions go?
•If T of product and reactants are equal, ΔG=ΔH‐TΔS
•For a chemical reaction, ΔG=Gproducts – Greactants
•Reaction has thermodynamic potential to react if ΔG is negative; larger ΔG, more likely spontaneous reaction. `
How fast will they go?
• RATE LAW, reactants A, B, C
dA/dt = ‐kAmBnCp
Where: •k is rate constant, units depend on concentration units and exponents
•m, n, p define order of reaction with respect to each reactant
•order of overall reaction is m + n + p
How fast will they go?
• Rate constant can be temperature dependent
•
Arrhenius Equation K = A exp(‐Ea/RT)
•Ea is ACTIVATION ENERGY
•Catalyst for a reaction lowers the
activitation energy
How fast will they go?
•Reaction Mechanisms
•
Multi‐step series of simple reactions leading to final product are common in natural systems
•
Typically have “rate‐limiting”
step determining kinetics
•
Catalyst is something which
lowers the activation energy (and increase rate) without being changed itself
Example: ozone depletion
•“Oxygen only”mechanism—Chapman Model
•
O2 + hν → O + O
J1
•
O + O2 + M → O3 + M
k2
•
O3 + hν → O + O2
J3
•
O + O3 → 2O2
k4
•Balance of first and fourth reactions determines amount of O3, solve at steady‐state, consistently overestimates concentration of ozone in stratosphere‐ Therefore, must be other ozone‐
depleting reaction
ozone depletion
•“series of catalytic ozone‐destroying reactions with NO/NO2 as catalyst:
•
•
NO + O3 → NO2 + O2
NO2 + O → NO + O2
k5
k6
•
O3 + O → 2O2
NET REACTION
•Each NO, NO2 molecule goes through cycle hundreds to thousands of times before being rained out
Biogeochemical Systems and Physical Chemistry
• Non‐equilibrium states are typical
– Biological input of energy
– Typically boundaries
– Often very slow kinetics
• Always consider possible role of catalysts
• Environmental gradients typical
– Temperature, pressure, Ph, Eh
2) Acids and bases
pH = ‐log10[H+]
ACID‐BASE REACTIONS
 Involve exchange of protons in liquid phase Bases dissociate to give [OH‐] (are bitter, slippery, and also turn litmus paper blue!)
Acids dissociate to give [H+] (are sour, non‐slippery, turn litmus paper red!)
Lots of complexity: Strong vs. weak, hard vs. soft, etc.. Dissociation constant of water
•H2O → H+ + OH‐, log K = 10‐14
•pH of a neutral solution of water = 7
•Acid‐base balances within Earth system generally involve elements of relatively high abundance
ACID‐BASE REACTIONS
•C, N, S are especially important in Earth System Acid‐base Chemistry
All undergo ~ same oxidation rx to acids‐
R+ H20 => HxR0x Eg: H2CO3, H2S04, H2NO3
3) Redox Reactions
Redox and Life
“As unstable chemical entities fueled by a continual diet of decomposing compounds, sustained by a constant energy flux from oxidation reactions, we are motivated by more than pure academic curiosity to study redox chemistry; it is a simple (or not so simple) matter of life and death….”
Anonymous (to me) Redox: super basics
 OXIDATION STATE OF AN ELEMENT
Example: C(‐IV)
C(0)
C(+IV)
methane
C, CH2O
CO2 and other dissolved
inorganic carbon species
In any compound, assume O(‐II) and H(+I)
Redox: super basics
 Oxidation‐reduction reactions or redox reactions—
= exchange of electrons “ LEO the lion says GERRRR*”
Loss of electrons is oxidation
Gain of electrons is reductant
Reductant=electron donor
Oxidant=electron acceptor
Reducing force=electron supply
Oxidizing force=electron drain
*anonymous, strikes again..
Super important: redox of C, N, O, S, plus Fe, Mn…
“Because life is a matter of
stochiometry as well as one of energetics, the cycles of these elements are then tightly coupled in two different ways.*”
Anonymous; speaks again
Except for P: Phosphorous has no redox chemistry.
EQUILIBRIUM CONSTANTS FOR REDOX REACTIONS
 Redox reactions generally expressed as individual half redox reactions
 (Really aren’t free electrons in solution in way there are H+)
Half reactions written by convention as reductions
BY CONVENTION,
H+ + e‐ = ½ H2(g), ΔG0= 0. K = 1
COMPLETE REDOX REACTION
Ox1 + Red2 = Red1 + Ox2
Transfer of n electrons from Red2 to Ox1
Complex possible notations, all same idea:  FOUR DIFFERENT, EQUIVALENT SCALES TO DEFINE THERMODYNAMIC CHARACTERISTICS OF REDOX REACTIONS

pe = ‐ log10[e‐1], nondimensional scale like pH for electron activity

E0H, redox potential (volts)

K, equilibrium constant

K=[products]/[reactants]

ΔG0, free energy change

ΔG0=‐RTlnK
BIOGEOCHEMICAL CYCLES AND REDOX
• Photosynthesis
• Photochemical reactions
• Oxidation of organic matter
• Reduced compounds from Earth’s interior (e.g., hydrothermal vents)
A most fundamental biogeochemistry example:
Organic matter reduction/oxidation
1) Photosynthesis captures sunlight,
(increasing the free energy of the planet)
CO2 + H2O  CH2O + O2
= a reduction of carbon
C(+2)  C(0) Delta G = ‐686 kcal/ mol
CH2O
POM‐sink
2) Organic matter oxidation, electron acceptor series
Redox + Acid/Base: mechanistic control of most biogeochemical cycle reactions
ACID‐BASE: In natural systems, some ‘master’ system (often dissolved inorganic carbon systematics) sets the overall pH. Then can use pH as “master variable” to figure out degree of dissociation of other acids/bases. (pH diagrams)
REDOX: In natural systems, some ‘master’ system (typically either the presence of dissolved oxygen or dissolved sulfide), sets the overall redox state. Then can use a ‘master variable’ to describe the state of the system, figure out the poise of other redox couples. (pE‐pH diagrams or Eh‐pH diagrams)
4) Stable Isotopes
I. Introduction
A. What are stable isotopes?
 Recall: Stable isotopes vary in mass, but chemistry is unchanged (ie outer shell electron configuration)
Eg: Stable isotopes of Oxygen: Isotope
# Protons
#Neutrons
Atomic Wt
Atom%
16
8
8
16
99.759
17
8
9
17
0.037
18
8
10
18
0.204
O
O
O
Distribution and Occurrence
Main Organic Stable Isotopes
1. All the major bioactive elements (except P) have multiple stable isotopes.
2. Within this group, the light isotope (L) is consistently more abundant than the heavy (H) counterpart(s).
3. It is very small (ppt) differences in (H/L) that constitute the basis of using stable isotope signatures as geochemical source and process indicators
Isotope
1
H
2
H (D)
12
C
13
C
14
N
15
N
16
O
17
O
18
O
32
S
33
S
34
S
36
S
Atom%
99.985
0.015
98.89
1.11
99.63
0.37
99.759
0.037
0.204
95.00
0.76
4.22
0.014
Why do useful differences occur in Nature?
 Reactions with kinetic component, always end up enriching in LIGHT isotope. • Note: Natural processes: often may have both a passive (diffusion) and active (biochemical reaction) component. BOTH parts will fractionate for light isotope. •
In fact, most natural fractionations = “chain of processes”
A bit more background:
 Isotopes of different mass have slightly different bond strengths. This affects their distribution because:
1) Lighter Isotopes tend to move faster
 rates of diffusive transport are different! (eg: 12CO2(g) ~1% faster than 13CO2(g))
2) Lesser effect: equilibrium distributions in a reversible reaction, where heavier isotope is typically concentrated in more strongly bonded form:
13CO2(g) + H12CO3‐(aq) = 12CO2(g) + H13CO3‐(aq) favored (1)
RESULT:
 Virtually ANY and EVERY chemical process (both biological and abiotic) fractionates stable isotopes to some degree
Eg: Photosynthesis, Respiration, Biochemcial reaction steps, Evaporation, Gas dissolution, Precipitation, etc etc etc.  OFTEN these fractionations are very predictable‐
 thus Stable Isotopes represent an imbedded tag in every molecule for 1) Source and 2) process history
Main Organic Stable Isotopes
Stable Isotope signatures of organic matter are among most powerful off all Organic geochemical tools!
A few examples of uses/ processes:
C 12/13: ultimate origin of fixed carbon
N 15/14: Trophic level processing
O & H ‐ Hydrologic cycles: eg: PaleoRainfall
S: ancient sulfer cycles. * C and N are most common (and easy) isotopes measured mainstays . Isotope
1
H
2
H (D)
12
C
13
C
14
N
15
N
16
O
17
O
18
O
32
S
33
S
34
S
36
S
Atom%
99.985
0.015
98.89
1.11
99.63
0.37
99.759
0.037
0.204
95.00
0.76
4.22
0.014
Some Terms and definitions:
“fractionation” = The amount of change in light vs heavy (Measured in “del” notation) created by a given process. “isotope effect” (typically means “kinetic isotope effect”) = fractionation occurs in the process of interest. 12
12
C
C
12
12
13
C
13
12
13
An Isotope Effect
(a phenomenon)
C
Fractionation
(observable quantity)
How isotope ratios are expressed:
1) Del (δ) notation: measure of relative ratio of heavy vs light isotopes a) expressed as ‰ ( “per mil”, or part per thousand) relative
to a (semi‐ arbitrary) standard. (why relative to a standard?)
b) Written : “δ 13C of sedimentary bulk OC was –27.5 ‰”
or “δ 15N of the bulk plankton tow from 100m was + 8 ‰”
Formal Definition:
H = (H/L)spl - (H/L)std x1000
(H/L)std
where H= heavy isotope and L= light isotope
Reference Standards:
Selected: To reflect the major source reservoir that processes are fractionating. EG: Ocean water, Atmospheric N, Carbon in deep space
Primary Standards
Standard mean ocean water
“
“
PeeDee belemnite (PDB)
Air
Canyon Diablo meteorite
Isotope Ratios
2
H/1H
18 16
O/ O
17 16
O/ O
13 12
C/ C
15 14
N/ N
32 34
S/ S
Ratios x 10-6
155.76
2005.20
373
11237.2
3676.5
22.22
How to make stable isotope measurements?
Basic Problem: ‰ Differences are VERY SMALL! Recall: the “per mil” is not a “direct” measure, but is a relative measure‐
IE actual carbon Std Ratio 12C/ 13C = already a BIG excess
 Boils down to measuring absolute differences in a small populations of atoms of the Heavy Isotope. (Analogy: being able to measure relative differences in distance of a cm or two, accurately and reproducibly of the scale of here to Paris) How its Done:
1) As with CHN measures: Oxidize OM fully to CO2, NOx, etc.  now, only have one kind of molecule to worry about detecting. 2) Now, need to measure RATIO of 13CO2 vs 12CO2 , or the ratio of masses 44 vs 45‐ at exquisite sensitivity!  Special kind of Mass Spec: “Isotope Ratio Mass Spectrometer” or “IRMS”
EA‐ IRMS setup:
Roast
(as w/ CHN)
sample
1050 deg
CO2 CO2
CO2
Molecular
Leak
So Ion
ur
ce
CO2 CO2
CO2
Ratio Mass
Spectrometer
o
60
Ionizing
Electron
Beam
Magnet
Detectors
Signal
Processor
Analyzer
Tube
Precision attained: typically to better than 0.1 ‰ !!!!
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