Marilyn and the ill-fated oxygen line, 1981 |
As a new staff member, it was exciting
to move into my own dedicated laboratory space. The original Geophysical
Laboratory on Upton Street in Washington DC was built in 1908, had 18-inch walls
and big hallways. Director Yoder assigned me a laboratory to share with
retiring staff scientist, Gordon Davis. Davis’ career focused on radiogenic
dating of rocks. His laboratory was primarily a preparatory, “semi-clean” lab
where he purified lead and other radiogenic isotopes for dating. For about a
year, we “shared” the space with my culturing equipment on benches in the clean
lab. It was an arrangement that would never really work, because I did not have
full control over my lab space. Only many years later did I realize that this
arrangement put the start of my career at a moderate disadvantage. Fortunately
for me, after one year Gordon fully retired and I was able to renovate the
laboratory, move in two isotope mass spectrometers, and build my first original
vacuum line.
After investigating hydrogen isotopes
for several years, my interests turned to oxygen in organic matter with an
experimental plan similar to what I had carried out for hydrogen isotopes.
Oxygen has three stable isotopes: 16O, roughly 99.6% of all oxygen,
with six protons and six neutrons; 17O, at <0.1% the rarest
isotope, has six protons and seven neutrons; 18O, 0.3% of all of the
oxygen isotopes, has six protons and eight neutrons. Our isotope ratio mass
spectrometers (IRMS), then, were set to measure 16O and 18O,
and were not tuned to measure 17O.
Oxygen enters in the biological realm
from three sources: water that is taken up directly into cells, oxygen in food
for animals and heterotrophic microbes, and oxygen in the air for aerobic
organisms. Oxygen in air is metabolized to water during aerobic respiration in
animals and microbes and during photorespiration in plants. Because the
analytical methods were limited to the measurement of carbohydrates, almost
nothing was known about how oxygen and its isotopes were incorporated into
other important biological molecules such as collagen, the protein in bones and
teeth, and keratin, the protein in feathers.
Oxygen isotopes in water have very
different compositions relative to the oxygen isotope composition of oxygen in
air. I took advantage of these natural differences in designing my first
experiments. Using microalgae growing in cultures, I wanted to know whether any
of the oxygen in air was permanently incorporated into the cell’s organic
substances. If it did, I should have a way of determining the oxygen isotope
composition of atmospheric oxygen over geological time by analyzing kerogen or
fossil proteins. The ultimate goal was to understand long-term global oxygen
cycles, which operate in tandem with global carbon cycles that are linked
through living organisms’ respiration processes.
I had modified a method for my hydrogen
studies that enabled me to separate lipids from carbohydrates and proteins in
cells. I cultured cyanobacteria and green algae in liquid broth under several
conditions. I varied the oxygen isotope composition of the water, an easy thing
to accomplish by adding “heavy” water composed of H218O,
while keeping the isotopic composition of air constant. I modified the
microalgae’s ability to carry out photorespiration by varying carbon dioxide
concentrations. High concentrations of carbon dioxide inhibit photorespiration.
I thought I was particularly clever.
The analysis of organic oxygen was, at
this time, limited to carbohydrates. The Caltech group, Sam Epstein and his
students, was analyzing cellulose in plants and tree rings. John Hayes and his
student Kim Wedeking at Indiana University were revising a method, originally
described by Rittenberg and Ponticorvo (1956), for analyzing proteins and
kerogen—the intractable molecule that forms the bulk of organic matter
preserved in ancient sediments and rocks. These compounds contain nitrogen,
sulfur, and other elements besides carbon and hydrogen. Tom Hoering and I followed
their work closely. Their method, published only in Kim Wedeking’s
dissertation, was based on a reaction of organic matter with mercuric chloride
(HgCl2) to form a mixture of carbon monoxide (CO) and carbon dioxide (CO2). Tom
Hoering and I were working on a similar method. Organic matter was heated in an
evacuated sealed glass tube with HgCl2 at 500°C. The reaction of this compound with the
organic matter resulted in products that included both CO and CO2, along with
hydrochloric acid and other impurities.
Hoering and I built a system—a gas
chromatograph—to separate the CO from the CO2 and the other
impurities (Hoering and Estep, 1981). Gas chromatography is an analytical
method designed to separate mixtures of compounds based on their chemical and
physical characteristics, e.g., boiling point and reactivity. No one builds
their own gas chromatograph these days! We started out with a 5 meter long ¼”
copper tube, a funnel, and beaker full of molecular sieve, a compound known to
separate simple gases like CO2. The copper tube was suspended in the
stairwell of the Geophysical Lab. Hoering stood at the top with the funnel and
beaker of molecular sieve. I stood halfway down tapping the side of the tube to
promote even packing. Once filled, we rolled the tubing around a 2” diameter
pipe to form a neatly wound “column” that was then fastened to our analytical
system.
We connected a series of valves to move
the gases from our reactions through the gas chromatograph. We had a simple
detector attached to a paper, strip-chart recorder, that I used to measure the
yields of CO and CO2. As the gases were pushed through the column
with helium gas under pressure, I manually flipped the valves to capture the CO
and CO2, while letting the impurities exit.
CO was converted to CO2 by reacting
it in a high-voltage discharge chamber where excess carbon is plated out on
platinum electrodes. The method was never robust, and as a byproduct of the
high-voltage discharge in the reaction chamber, nitric oxide (NO2) was
formed from the nitrogen from proteins in the organic matter. Multiple analyses
of biological samples introduced NO2 unwittingly into the metal guts of our
homemade IRMS. NO2 is notoriously sticky on metal surfaces. The vacuum pumps on the IRMS
couldn’t get rid of it.
Eventually the IRMS would not reach the
level of vacuum needed to operate and measure isotope compositions. This was a real
disaster. Tom Hoering removed the metal guts of the mass spectrometer called
the flight tube. He took it out onto the lawn of the Geophysical Laboratory and
sandblasted it to remove the contaminating nitric oxide. In the environment of
the Geophysical Lab, where everyone seemed to have research successes all the
time, I felt deflated, maybe even an “imposter.” I didn’t know anyone else who
caused such damage to an isotope ratio mass spectrometer. [With time, I did
hear of equally awful things--blowing up expensive metal “bellows” like balloons;
dripping strong acid onto expensive electronics, for example.] I immediately,
and sheepishly, moved on to other projects while Hoering reassembled the IRMS.
The oxygen isotope vacuum and
extraction line I designed for this project was my first. As a “girl”, I wasn’t
taught how to use tools, really. I figured things out on my own—sometimes not
until I did something stupid, did I learn to do it correctly. Parts of my line
were constructed using a type of metal compression fitting composed of four
important parts. The fittings, called Swagelock fittings, had a small cone, a
ring called a ferrule, a nut that screwed on, and the body of the fitting made
out of brass. Assembled correctly, the fittings made a water and gas tight
seal. The cone went on first, then the ferrule, then the body, and finally the
nut. I used these fittings to construct a water line designed to cool parts of
the reaction line. When the water was first turned on, the line dripped at
every connection. I did not know how to assemble a Swagelock fitting properly.
Tom Hoering strolled into my lab and snorted: “You have these in ass-backward.”
I had reversed the cone and the ferrule during assembly.
A lot of work- for not much, 1982 |
I gave myself a grade of “C-“—sloppy
work. On the science side, it was learned a few years later that all of the
oxygen in a living organism is in equilibrium with cellular water. My original
hypotheses wouldn’t have produced much. But, I did learn how to build a gas
chromatograph, a skill needed for later work on oxygen isotopes in a project
with molecular O2 that did work out. Even later, I used this knowledge to understand
much more complex continuous flow gas chromatography-combustion-IRMS
(GC-C-IRMS) techniques that I still use today.
The analysis of oxygen isotopes in
organic matter was my first failed project, but I learned a lot during the
process. The most important lesson was learning to deal with failure. As a high
school student, I got a couple “C”s in history classes to scuff up my
reputation as a nerd. In college, I breezed through with only a couple more
“C”s in genetics and statistics. In the early 1970s, Penn State taught genetics
through its Ag departments. The class concentrated heavily on traits of
chickens and corn. I was not at all interested and sloughed off. In grad
school, with the exception of a difficult biochemistry class on Intermediary
Metabolism in which I earned a “B”, I was a straight A student. I needed to learn
about real failure.
A scientist’s life is filled with a
certain amount of rejection: manuscripts, proposals, and ideas. Only rarely
does one get affirmation and accolades. Learning to not give up, but to keep
searching for the next promising idea, is key to a successful long-term career.
As a University professor, I counsel undergraduates to take difficult courses—quantum
physics, climate modeling, or Bayesian statistics—and figure out how to learn difficult
concepts and accept that you aren’t perfect. For those straight A students,
this can be frightening! What better time to know that part of yourself as to how
you deal with small failures.
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