Lecture 5 Stable Isotopes
Isotopes of Elements
Types of Isotopes
Chart of the Nuclides
Measurements
Delta Notation
Isotope Fractionation
Equilibrium
Kinetic
Raleigh
See E & H Chpt. 5
Key questions:
What are isotopes?
What are the types of isotopes?
How do we measure isotopes?
How do we express measurements of isotopes?
What is isotope fractionation and how do we express it?
What is equilibrium isotope fractionation?
What is kinetic isotope fractionation?
What is Raleigh distillation?
What are some applications of stable isotopes?
Isotopes of Elements
Atomic Number = # Protons = defines which element and its chemistry
Atomic Weight = protons + neutrons = referred to as isotopes
Different elements have different numbers of neutrons and thus atomic
weights.
Example: Carbon can exist as 12C, 13C, 14C
How many protons and neutrons in each of the C isotopes?
12C = 6P, 6N
13C = 6P, 7N
14C = 6P, 8N
1 chemical, many isotopes!
Where do Isotopes come from?
In the beginning (Big Bang), light elements of H and He were
formed (and a little bit of Li)
Nuclear reactions (ie: fusion) in stars created the remaining
elements (and are still creating), some of which have since
decayed to more stable elements
There are 92 naturally occurring elements – some are stable,
some are not
Types of Isotopes
Isotopes can be categorized into 2 categories:
Stable isotopes – Isotopes that do not decay over the timescale of
earth history (4.5 billion years)
Radioactive isotopes – Isotopes that spontaneously convert into
other nuclei at a discernable rate
The chart of the nuclides (protons versus neutrons)
for elements 1 (Hydrogen) through 12 (Magnesium).
Valley of Stability
Most elements have more
than one stable isotope.
1:1 line
b decay X
X
a decay
Number of neutrons tends
to be greater than the
number of protons
Full Chart of the Nuclides
1:1 line
Valley of Stability
Examples for H, C, N and O:
Atomic Protons
Neutrons
Weight (Atomic Number)
Hydrogen H
1P
0N
D
1P
1N
Carbon 12C
6P
6N
13C
6P
7N
14C
6P
8N
Nitrogen 14N
7P
7N
15N
7P
8N
Oxygen 16O
8P
8N
17O
8P
9N
18O
8P
10N
% Abundance
(approximate)
99.99
0.01
98.89
1.11
10-10 1/2 = 5730 yr
99.6
0.4
99.76
0.024
0.20
% Abundance is for the average Earth’s crust, ocean and atmosphere
Isotope Ratio Mass Spectrometer (IRMS)
How we measure stable isotopes – the IRMS
1. Input as gases
2. Gases Ionized
3. Gases/ions accelerated in vacuum
4. Gases bent by magnetic field
according to mass
5. Gases detected
1.
2.
5.
3.
4.
Isotopes are measured as ratios of two
isotopes.
Standards are run frequently to correct
for instrument stability
Nomenclature – δ Notation
Report stable isotope abundance as ratio to most abundance isotope (13C/12C)
- Why? The ratio can be measured very precisely.
BUT – any differences in the isotope ratio can be very very small so we use δ (“del”) notation
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δ = “delta” or “del” (if you’re real savvy), units are per mil (‰)
Where
H = moles of heavy isotope
L = moles of light isotope
R = H/L
δ tells us how much the sample deviates from the
standard
The sign of δ
Compared to the standard
Standards Vary
Each standard has a well defined H/L ratio
Example 1:
The IRMS standard for C is PDB (13C/12C = 0.011237)
Your sample has an 13C/12C = 0.010957.
What is δ13C in ‰ for the standard?
For the sample?
Isotopic Fractionation
• All isotopes of a given element have the same chemical properties
• Small differences in the distribution of the isotopes in materials because heavier
isotopes form stronger bonds and move slightly slower
A heavier mass = stronger bond
You would need a stronger string to hold two bowling balls together than you would
need to hold two golf balls together
Isotope Fractionation = process that results is differences in delta values in products
and reactants
Example: condensation of water vapor
H2O(g) <=> H2O(l)
In a closed water sample:
δ 18O of H2O(g) = -1‰ (Atlantic)
δ 18O of H2O(l) = -10‰ (Atlantic)
Because of isotope fractionation!
a and ε Nomenclature
Fractionation Factor = a
a=
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a is unitless
If a = 1, no fractionation
If a >1, more heavy in product
If a <1, more heavy in reactant
Difference fractionation Factor = ε
e =d H
products
dH
-
reactants
If ε = 0, no fractionation
If ε > 0 , more heavy in product
If ε < 0 , more heavy in reactant
ε is in permil (‰)
e »1000´(a -1)
Example #2: condensation of water vapor
H2O(g) <=> H2O(l)
In a closed water sample:
δ 18O of H2O(g) = -10‰
δ 18O of H2O(l) = -1‰
What is ε and a of this reaction?
Two kinds of Isotope Fractionation Processes
1. Equilibrium Isotope effects
Occurs in equilibrium reactions (reactions can go both ways) if the system is in
equilibrium
Chemical equilibrium
Phase changes (closed system)
Distributes isotopes in a system so that the total energy of the system is minimized
Heavier isotope equilibrates into the compound or phase in which it is most stably
bound
Within a molecule (CO2 vs HCO3-)
Between molecules (CO2(g) vs CO2(aq) )
Usually applies to inorganic species. Usually not in organic compounds
Due to slightly different free energies for atoms of different atomic weight
Usually temperature dependent!
Differences in vibrational energy is the source of the fractionation.
Heavier isotopes wind up in the compound where it is bound more strongly
Example #3: Condensation of Water Vapor in a closed container
H2O(g) <=> H2O(l)
H216O(l) + H218O(g) ↔ H218O(l) + H216O(g)
In a closed container:
δ18O of H2O(g) = -10‰
δ18O of H2O(l) = -1‰
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Is this reaction an example of an equilibrium isotope effect? How can
you tell?
Does the 18O “prefer” to be in the gas or liquid phase? Why?
Example #4: Bicarbonate system
The carbonate buffer system involving gaseous CO2(g), aqueous
CO2(aq), aqueous bicarbonate HCO3- and carbonate CO32-.
One step of that reaction:
CO2(aq) + H2O ↔ HCO3- + H+
δ 13C of CO2(aq) = 1‰
a = 1.0092 at 0ºC and 1.0068 at 30ºC
(The IRMS standard for C is PDB (13C/12C = 0.011237))
Is this reaction an example of an equilibrium isotope effect? How can
you tell?
What is the final δ13C of HCO3- at 0ºC at 30ºC?
Is 13C more stable as CO2(aq) or HCO3-?
Is there more or less fractionation at higher temperatures?
2. Kinetic Fractionation
Occurs in
unidirectional (irreversible) reactions
reversible reactions that are not yet at equilibrium
diffusion or differential bond breaking
Heavier isotopes move more slowly (KE = ½ mv2)
Therefore react more slowly
Reaction products are depleted in the heavy isotope relative to the
reactants
All isotopes effects involving organic matter are kinetic
Why do heavier isotopes move more slowly?
Same kinetic energy, despite isotope
E = ½ mv2
If E is the same and mass increases, the v must decrease
Examples of Kinetic Fractionation
Three types of kinetic fractionation:
1. Unidirectional reactions
Example:
Carbon fixation via photosynthesis:
12CO + H O -> 12CH O + O
2
2
2
2 faster
13CO + H O -> 13CH O + O
2
2
2
2 slower
Organic matter gets depleted in 13C during photosynthesis (decreases in 13C)
2. Reversible reactions that are not yet at equilibrium
Example:
Evaporation of water vapor if not in equilibrium (net evaporation ie: N .Atlantic)
H216O(l) -> H216O(g) faster
H218O(l) -> H218O(g) slower
Water vapor gets depleted in 18O during net evaporation (decreases in 18O)
3. Diffusion
Example:
Diffusion of H2Oacross a cell membrane
H216O(l) outside cell -> H216O(l) inside cell faster
H218O(l) outside cell -> H218O(l) inside cell slower
Water vapor gets depleted in 18O during net evaporation (decreases in 18O)
Equilibrium Fractionation vs Kinetic Fractionation
The difference depends on the reason for the fractionation
Equilibrium fractionation occurs so that the total energy of the system is
minimized via forming the most stable bonds possible
Equilibrium is related to bond stability of the isotope
Kinetic fractionation occurs because smaller molecules move faster than
heavier molecules and therefore react more slowly
Kinetic is related to the speed of the isotope
13C in carbon
reservoirs
E & H Fig. 5.6
13C of atmospheric CO2 versus time
See Quay, 1992, Science
Raleigh Fractionation
A combination of kinetic and equilibrium isotope effects
• Kinetic when water molecules evaporate from sea surface (net
evaporation b/c system is not in equilibrium)
• Equilibrium effect when water molecules condense from vapor
to liquid form
A isotope fractionation reaction where products are isolated
immediately from the reactants will show a characteristic trend in
isotopic composition.
Raleigh Fractionation - Concept
• Vapor depleted in 18O
compared to ocean water
• Air masses transported to
higher latitudes where it is
cooler.
• Rain enriched in 18O, removed
from system (cloud)
• Cloud gets lighter
• Rain enriched in 18O, removed
from system (cloud), but less
enriched
Raleigh Fractionation – Characteristic trend
Example: Evaporation – Condensation Processes
18O in cloud vapor H2O(g) and condensate (H2O(l) rain)
plotted versus the fraction of remaining vapor for a
Raleigh process. Idealized:
• 20ºC – All vapor -9‰
• Just colder than 20ºC – Condensate starts to form,
more enriched in 18O, but is removed from the
system (rained out)
• The vapor continues to condense as the temperature
decreases – becoming more and more depleted in
18O
• Fractionation increases with decreasing temperature
• Same pattern for D/H isotopes - different scale
because more fractionation during the condensation
(ε = +78‰ rather than +9‰)
• This trend is used to reconstruct local
paleotemperature from in Antartica and
Greenland from ice cores
18O variation with time in Camp Century
ice core.
18O was lower in Greenland snow
during last ice age
15,000 years ago 18O = -40‰
10,000 to present 18O = -29‰
Reflects
1. 18O of precipitation
2. History of airmass – cumulative
depletion of 18O
http://www.youtube.com/watch?v=nZC5EMPZDFA
Applications of Stable Isotopes
There are many applications of stable isotopes – especially in the study of past
conditions on earth
Three case studies in oceanography:
1.
2.
3.
Paleothermometer from foraminifera shells
Origin of organic matter
Estimate primary production in marine systems
Case study: 18O of forams in sediment to reconstruct
paleotemperature
HCO3- + Ca2+ ↔ CaCO3(s) + H+
Fractionation of 18O is temperature
dependent and well quantified in labs
The 18O of CaCO3 precipitated in forams
reflects the temperature
Preserved in marine sediments
Complicated because although this
relationship is well defined, depends on a
known 18O of water . That may change
due to ice volume.
Case Study: Estimation of temperature in
ancient ocean environments
CaC16O3(s) + H218O  CaC18O16O2 + H216O
The exchange of 18O between CaCO3 and H2O
The distribution is Temperature dependent
last
interglacial
Holocene
last glacial
18O of planktonic and benthic foraminifera
from piston core V28-238 (160ºE 1ºN)
Planktonic and Benthic differ due to differences
in water temperature where they grow.
Planktonic forams measure sea surface T
Benthic forams measure benthic T
Assumptions:
1. Organism ppted CaCO3 in isotopic equilibrium
with dissolved CO322. The δ18O of the original water is known
3. The δ18O of the shell has remained unchanged
Case study: 18O of forams in sediment to reconstruct
paleotemperature
Does the 18O of water in the ocean change over time?
Large scale Raleigh distillation
Net transfer of water from ocean to continental ice sheets make ice
very depleted in 18O and the oceans enriched in 18O, increasing the
18O of water about 1‰
In some cores, pore water can
be measured directly, which
gets around this issue.
Case study: 13C of bulk organic matter to determine source
Many people are interested in
the preservation in organic
matter in marine sediments
By looking at the 13C of an
organic material, it can say
something to how it was
produced (marine or
terrestrial) because the starting
material is so different in 13C
Complicated by C4 and CAM
plants.
C4 = grasses
Crassulacean acid metabolism, also known as CAM photosynthesis, is a carbon fixation pathway
that evolved in some plants as an adaptation to arid conditions
Case study: Profiles of DI13C and 18O to estimate primary
productivity
The profiles of DI13C and 18O
can be used to estimate primary
productivity
More photosynthesis in surface
results in a heavier DI13C,
resulting in a more positive 13C
in surface DIC
During respiration, 16O is
preferentially taken up, resulting
in a more positive 18O “left
over” in the water (obvious at O2
minimum)
Why does the 13C decrease
slightly at the O2 minimum?
North Atlantic data