|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.