The Foundations of Physical Organic Chemistry - ACS Publications

The Foundations of Physical Organic Chemistry: Fifty Years of the James Flack Norris Award Downloaded from pubs.acs.org by UNIV OF CALIFORNIA SAN DIEGO on 12/05/15. For personal use only.

Chapter 10

60 Years of Research on Free Radical Physical Organic Chemistry K. U. Ingold* National Research Council, Ottawa, ON K1A 0R6, Canada *E-mail: [email protected]

In this Chapter the reader is conducted on a gentle cruise through the green and pleasant land of mechanistic radical chemistry. My love of chemistry arrived with a bang (potassium metal + water) and I was already enamored by “how fast?”. These two loves combined at the NRC as I helped to convert radical chemistry from a tarry, black art, into clean, practical, useful, predictable and comprehensible chemistry.

My father, Christopher Kelk Ingold, Professor and Head of the Chemistry Department at University College London (UCL, 1930-61), was the first recipient of the James Flack Norris Award for Physical Organic Chemistry (1965). This was a great honor and one that he very much appreciated. (To both of us, this was a very personal example of Americans’ generosity of spirit, particularly since it was a very distinguished American, Louis Hammett, who coined the words “Physical Organic Chemistry”. Hammett’s seminal role in this subject was honored by his receipt of the second J. F. Norris Award in 1966). I went to the ACS meeting in Atlantic City to see the Award presented and then drove my father, and two chemists from UC Santa Barbara, Clifford Bunton (ex-UCL) and Tom Bruice (Norris Award winner in 1996), back to my home in Ottawa. I live on the Rideau River which is at the bottom of my garden (in Summer, in Spring the garden is at the bottom of the river!) and had a ski boat (woefully underpowered relative to later boats). My father surprised all of us by expressing a wish to water ski though he’d never done it before. With many misgivings I agreed and even drove the boat. To our joint surprise, my father got up the first time he tried, but then he had always kept fit. He used to do, and encouraged me

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to do, a lot of walking, rock climbing (Wales), and mountaineering (Alps), sports he continued into his late 50s (Figures 1-4).

Figure 1. Getting ready for the great water-ski caper: (right to left) John Ingold (age 5), Bruice, Bunton, CKI, and KUI.

Figure 2. CKI skiing for first time at age 72

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Figure 3. KUI skiing at 85. One of the advantages of living with a river!

Figure 4. Other than the first two pictures in this Chapter, this is the only known snap of myself and my father together. We are sitting with Lionel Jones, husband of my sister, Dilys, in their garden in Sussex.

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C. K. Ingold was chosen to be the 1st recipient of the Norris Award largely for his pioneering mechanistic and kinetic studies of heterolytic chemical reactions. The statement: “I should tell you that in this Department, homolysis, even between consenting adults, is grounds for instant dismissal.” was not made by CKI but by Peter de la Mare as comment on Alwyn Davies’ wish to start work on peroxide homolysis (1). That I would also become a research chemist, like my father (and mother) was unsurprising, as was my very conscious decision to steer well clear of heterolytic chemistry. Emigration to Canada (with a freshly minted D.Phil from Oxford) in 1951 at the age of 22 may have been one of my wiser decisions since my father and I had, perhaps, too much in common to get on as well as we actually did if we’d lived closer to one another, cf., pictures of the two of us through the years (Figures 5 and 6). I believe my love of chemistry was sparked by an event in 1941-2. World War II disrupted everyone’s life. My father and mother were evacuated with half of UCL’s Chemistry Department to Aberystwyth (Aber), Wales. Eventually, (summer 1940) they were joined by me and my two sisters. I had attended at least 3 schools after evacuation and before my arrival in Aber in June, 1940. While living in Aber I attended 3 more schools (and went to yet another upon our return to London in 1944) before my entry as a chemistry undergraduate into UCL. My frequent school changes were not due to (sometimes well-deserved) expulsion, they were just a consequence of Hitler’s disruptions. In Aber there was little to do on a weekend. In a forlorn attempt to keep me out of trouble, my father arranged for the Department’s glass blower to teach me some of his skills on Saturday mornings (which proved very useful later). However, time hung heavy (until I discovered the joys of rock climbing). One Sunday my father took me into the labs and, after finishing whatever work had dragged him there, he decided to show me some interesting chemistry. Interesting was an understatement! He wanted to show me what happened when a small piece of sodium was dropped into water. Fortunately for me, he didn’t know where the sodium was kept. However, he found the potassium and, being a bit rattled by his failure to find the sodium in “his” new-to-him lab, cut off several grams of potassium and threw it straight into a sink prefilled with water. Spectacular! There was an enormous bang, far louder than any I later heard from V2 rockets landing in London. The sink was instantly dehydrated, small puddles by the hundred appeared on the floor, the bench-tops, everywhere, each one with little bits of purple-burning potassium speeding around on its surface. We were not wearing safety glasses, but survived unharmed. Even though we had to clean the entire lab (to keep all others in the Department ignorant of the Professor’s sins) I was instantly and totally hooked on chemistry. Hooked on chemistry, yes, but I’d been hooked on kinetics almost a decade earlier. Well before I was a teenager, I overheard some adults describing a lady who lived near us as a “fast woman”. I’d seen the lady in question and she didn’t look fast to me, though I did consider the possibility that high-heels gave speed on the running track. Since this seemed a mite unreasonable, I decided that when I grew up I would adopt an experimental approach and find out what made a woman fast. The results of these experiments were so rewarding that I’ve remained an 226


experimentalist who measures how fast things go, though today it is only chemical reactions, unfortunately.

Figure 5. CKI and KUI as ca. 12 year olds.

Figure 6. CKI and KUI in their 70s. 227


Four years (and 2 Post-Docs) after I arrived in Canada, I abandoned gas phase chemistry to accept a position in the Division of Applied Chemistry at the National Research Council in Ottawa (despite snide comments from many of my ‘pure’ chemistry friends). I wanted to work in the liquid phase because there is so much more interesting chemistry there than there is in the gas phase (witness, the entire field of biochemistry). I was hired to investigate the mechanism of oxidative degradation of automobile engine oils and to try to find ways to retard such degradation. A few years of glacially slow progress as a “real” applied chemist working on “real” engine oils convinced me to go back to basics. To this end, during 1960 I wrote my first Chemical Review, Inhibition of the Autoxidation of Organic Substrates in the Liquid Phase, in order to acquaint myself fully with the current state of knowledge (2). My career is now “book-ended” by my second Chemical Review (2014): Advances in Radical-Trapping Antioxidant Chemistry in the 21st Century. A Kinetic and Mechanistic Perspective (3). Although antioxidants remain my “first love”, the present Norris Chapter will show that my research has extended into a few other areas of physical organic chemistry.

Mechanism of Inhibition of Hydrocarbon Oxidation by Phenols In the early 1960s I started work on this topic. Literature results were very confusing. Controlled oxidation of a hydrocarbon, RH, is achieved by thermal decomposition of an azo-initiator, iN=Ni, and occurs by a two-step chain reaction:

Such oxidations are retarded by phenols, ArOH, but not by anisole, PhOMe. This implied that the phenolic OH group was required for inhibition. The obvious inhibition step is an H-atom transfer by which a chain carrying peroxyl radical is converted into a non-chain carrying phenoxyl radical:

However, if this reaction was responsible for the inhibiting effect of ArOH, it would be expected to show a deuterium kinetic isotope effect, DKIE, but none had been found. The DKIE experiments involved measuring the rates of azo-initiated oxygen uptake by some readily oxidizable hydrocarbon (e. g., cumene or Tetralin) under an atmosphere of O2. To the hydrocarbon was added, in mM concentrations, a phenol, ArOH, and in a matched experiment, the same concentration of the corresponding pre-formed O-deuterated phenol, ArOD. In such experiments the reduction in the rate of oxygen uptake induced by the ArOH 228


and ArOD were identical, indicating that there was no DKIE for inhibition. This negative result was confirmed by many authors (myself included) and it led to some extraordinary mechanistic suggestions in the literature. During the writing of my first Chemical Review, I realized one evening that these inhibition experiments actually had no chance whatever of uncovering a DKIE (if there really was one). Rates of oxidation were followed in sealed systems under pure O2 by measuring the consumption of O2. What I belatedly realized was that to measure such a rate required the absorption of an appreciable amount of O2. In the equipment of the day, this measurement required the uptake into the hydrocarbon of a higher concentration of O2 than the concentration of ArOH or ArOD that had been added. This meant that in order to measure the rate of O2 uptake it was necessary to make a higher concentration of the substrate’s hydroperoxide, ROOH, than the concentration of added ArOD. I had recently completed some IR studies on phenols, incidentally demonstrating (contrary to then accepted wisdom) that the O-H bond in 2,6-di-tert-butyl phenols lay in the aromatic ring plane and that 2-tert-butylphenols existed as an equilibrium mixture of syn and anti structures. This work had taught me that OH-containing compounds underwent proton exchange with ArOD extremely rapidly. I concluded, therefore, that in all the experiments searching for a DKIE in reaction 5, essentially all ArOD had been converted to ArOH before the rate could be properly measured.

These thoughts led me to suggest to my first Post-Doc, J. A. (Tony) Howard, that he add a drop of D2O to a hydrocarbon oxidation retarded by 2,6-di-tertbutyl-4-methylphenol (BHT, the major phenolic antioxidant used commercially) so as to make sure that the BHT remained fully deuterated during the reaction. I also suggested using styrene as the oxidizable substrate because it was known to yield a polyperoxide (with no exchangeable H-atoms) rather than a hydroperoxide. Tony’s results (4) were a game changer in antioxidant research. He found a DKIE of 10.6 at 65 °C, thus demonstrating that phenols are antioxidants because they donate their phenolic H-atom to an attacking peroxyl radical.

The length of time the oxidation is retarded by the phenol, known as the induction period, τ, is given by:

Work on phenolic antioxidants was continued by an examination of the effect of ring substituents on the abilities of phenols to retard the autoxidation of styrene. Electron donating (ED) para and meta substituents improved antioxidant activity while electron withdrawing (EW) substituents reduced activity. In these experiments, the rates we were measuring only allowed us to determine the rate constant ratio, k2/k5. A Hammett plot of log(k2/k5) against Brown and Okamoto’s σ+ values for the substituents yielded a straight line. This was completely unexpected because σ+ substituent constants had been derived from the relative 229


rates of a heterolytic reaction, the ionization (solvolyses) of m- and p-substituted cumyl chlorides in aqueous acetone.

The reason for the linear free energy relationship between log(k2/k5) and σ+ was only uncovered many years later. It arises because the XC6H4O-H bond dissociation enthalpies (BDEs) also correlate with σ+. This last correlation was recognized (by an outstanding Post-Doc, Gino DiLabio) to arise because the O• moiety is, like the +CMe2 group, an extremely powerful EW substituent. As a consequence of the EW effect of O•, ED substituents weaken and EW substituents strengthen the ArO-H BDE, making the antioxidant reaction 5 thermodynamically more and less favorable and, hence, faster and slower, respectively. In fact, the OH BDEs in substituted phenols and the N-H BDEs in substituted aromatic amines trend with σ+ due to the electron-poor nature of O• and N• as discussed in a 2004 Account: Bond Strengths of Toluenes, Anilines, and Phenols. To Hammett or Not (5). The mechanism of retardation of hydrocarbon oxidation by diarylamines, Ar2NH, was also contentious until we demonstrated an NH/ND antioxidant DKIE. This class of industrially important antioxidants therefore owed their activity to donation of their amino H-atom to ROO• radicals. Thus both phenols and aromatic amines are Radical Trapping Antioxidants, RTAs. If further progress in antioxidant chemistry was to be made it was necessary to know the magnitude of k5 rather than k2/k5 rate constant ratios. After all, my ultimate goal was to see whether a phenolic antioxidant could be designed that reacted with the first peroxyl radical it encountered! This Holy Grail was achieved several decades later by Derek Pratt, who had worked with me in 1998-99 as a NRC summer student, and there caught “grail fever”, see reference (3) for the full story. Such an antioxidant could not be bettered since reaction 5 would be diffusion-controlled (with k5 ~ 109 M-1s-1). It had already been shown that absolute rate constants could be determined for a number of free radical polymerization chain reactions. The main focus of research in my lab therefore shifted to the measurement of absolute rate constants for radical reactions, a subject that looked boundless and a subject to which I have (mainly) stayed faithful.

Absolute Rate Constants for Hydrocarbon Oxidation Some polymer chemists had been using photolysis to initiate their reactions and wanted to know how the rate of polymerization varied with the light intensity. Since no instrument was available to measure light intensities they inserted a metal disc, from which a 90° sector had been removed, between the light source and the reaction vessel. Rotation of this disc would block the light for 75% of the time and so the rate was expected to be 25% of the rate with full illumination. This was the case when the disc was rotated slowly but as the rotation rate was increased the rate of polymerization also increased up to a limit of 50% of the full-light rate. This odd behavior arose because chain termination in these polymerizations was a second-order process, Pn• + Pm• → Pn+m. Provided the rate of chain initiation, Ri, was known, the rate constant for chain termination could be calculated from the 230


duration of the flash at rotation speeds where the rate was >25% but

The Foundations of Physical Organic Chemistry - ACS Publications

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