2023 SL12009 SL12101 Lecture 3 Introduction to Radioactivity, Radio Isotopes and Radioactive Decay PDF

Lecture 3: Introduction to isotopes, radioactivity and radioactive decay Department of Dr Tim Woodman Pharmacy & Pharmacology  Go to menti.com and use code 2398...

Lecture 3: Introduction to isotopes, radioactivity and radioactive decay Department of Dr Tim Woodman Pharmacy & Pharmacology  Go to menti.com and use code 2398578 Why is knowledge about radiation important to pharmacists and pharmacologists? a. It’s not – it’s just more stuff to learn b. It’s used in the imaging of disease c. It’s used in pharmacology to study drug mechanisms d. It’s used in metabolic studies of new drugs The neutron… So far we have seen the proton and the electron. The Rutherford model (see in the last lecture) featured the small, compact nucleus with all the protons, surrounded by the electron clouds (in orbitals). There was a problem with this – the positively charged protons repel each other – what was keeping the nucleus intact? The idea of ‘nuclear electron’ was proposed, but had obvious problems – notably the observed atomic and molecular properties did not fit for the nuclear spin for such a system. In 1931 it was discovered that alpha particles interact with beryllium, boron and lithium to produce a very penetrating radiation, that was not influenced by an electric or magnetic field. Initially assumed to be gamma (γ) radiation, Chadwick showed that in fact the radiation consisted of uncharged particles of very similar mass to the proton – the neutron. Models of the atomic nucleus using the neutron were now introduced and satisfied the nuclear spin results. Later the neutron was shown to the source of beta (β) radiation – the neutron decays to a proton, emitting a high energy electron and a neutrino. The neutron keeps the nucleus intact by binding to protons via the nuclear force. This usually overcomes the proton-proton repulsion in stable nuclei. Radioactivity… Radioactivity was first discovered in 1896 by Henri Becquerel whilst working with phosphorescent materials. These materials glow in the dark after exposure to light, and Becquerel suspected that the glow might be related to the glow produced in cathode ray tubes by X-rays. He wrapped photographic plates in black paper and placed various phosphorescent salts on it. All were negative until he used uranium salts, which caused a blackening of the plate in spite it being wrapped in paper. He termed the source of the blackening “Becquerel Rays”. It became clear that the effect was not linked to phosphorescence, as non phosphorescent uranium salts and metallic uranium produced the same effect. At first it was assumed that the rays were similar to X-rays, more research showed that the picture was much more complicated. For instance all radioactive elements decay with the same mathematical exponential law. In addition different types of radiation were observed, some that would not penetrate paper, some that could only be stopped by thick lead plates. Rutherford and Soddy realised that some decay processes resulted in the transmutation of elements. Types of radiation – alpha (α), beta (β) and gamma (γ) By placing sources of radiation in a lead box with a hole, a beam can be produced (as the sources emit radiation in all directions). Becquerel showed that different sources produced beams that behaved differently when exposed to a magnetic field – deflecting in different (opposite) directions, or not at all – showing three classes of radioactivity (negative, positive and neutral). Rutherford termed these alpha, beta and gamma, based on their ability to penetrate matter. Magnetic field The alpha source was stopped by only a few Photographic plate centimetres of air – only in a vacuum was a spot observed. By contrast beta sources can To vacuum pump be stopped by aluminium sheet (1 mm thick). Gamma sources were only stopped by lead (several cm, or a metre of concrete). Gamma sources did not deflect, alpha particles were shown to positive and beta Radioactive particles negative. element Alpha particles Alpha particles were shown to be positively charged helium nuclei by trapping alpha particles in an emission chamber (the radiation passed through very thin glass into the evacuated chamber). Electrons were introduced and then studying the electronic spectra of the product showed that helium atoms had been formed. The best known source of alpha particles is from alpha decay of heavy atoms (heavier than 106 mass units, 106Te). In the process the parent atomic nucleus transforms (decays) into a different atomic nucleus, with a mass number reduced by four and an atomic number reduced by two. For example 238U decays to form 234Th and an alpha particle. 238 4 234 92 U 2 He + 90 Th (By convention the charge is ignored – we know that the alpha particle is doubly positively charged). Beta particles In beta decay a neutron transforms into a proton, with the emission of an electron and a neutrino. Effectively the mass number of the parent atom remains the same, but the atomic number increases by one unit. For example 14C decays into 14N: 14 14 6 C 7 N + e- + neutrino The beta particles produced (i.e. electrons) are energetic, and as we have seen easily pass through paper, and need aluminium sheet to be stopped. Gamma decay Gamma rays are usually produced after alpha or beta decay. The daughter nucleus that is formed is typically left in an excited state, and decays to a lower state by emitting a gamma ray photon. 60 60 An example of gamma ray production is the decay of Co. Initially Co 27 27 60 * * decays to Ni ( denotes an excited state), with the emission of an 28 60 electron. Then the excited Ni* decays to the ground state by emitting 28 two gamma ray photons (1.17 MeV and 1.33 MeV) Radioactive decay and the half life We have already seen that all radioactive elements decay with the same exponential law. The processes that govern whether an individual atom will undergo a decay cannot be predicted, however for a collection of such atoms the decay process can be calculated. The probability that a given radioactive nuclei will decay per unit time is a constant – called the decay constant, λ. The radioactive decay law is given as: N(t) = N0e-λt An easier way to view radioactive decay is the half-life – the time taken for half the radioactive nuclei in a given sample to undergo decay. For instance if a given radioactive isotope has a half-life of 14 days, then exactly half the atoms will decay during the 14 days. Half-lives vary extremely widely, ranging from millionths of a second all the way to billions of years. We will see the importance of the half-life in medicine shortly. Decay of unstable atomic nuclei First-order decay / exponential decay Nt = N0 e-t Nt = No. radionuclides left Nt N0 = No. of radionuclides at start Differentiating t = time elapsed  = decay constant = -N0 e-t t Nt Differential is rate of change - = At t Substituting At = A0 e-t At = radioactivity at time t A0 = Radioactivity at start 0.693 Half-life t½ =  1 Bq = 1 disintegration per second 1.0 Ci = 3.7  1010 Bq Decay of unstable atomic nuclei 0.693 0.693 Since t½ = we also know that  =  t½ 18F decays to 18O, with a half life of 109.7 minutes. What is the decay constant for 18F? 0.693  = t½ 0.693  = 109.7 min  = 0.00631 min-1 ( = 6.31 x 10-3 min-1) Isotopes, and why some nuclei are stable and others not… We are now familiar with the atomic structure of a compact nucleus, with protons and neutrons bound together by the nuclear force (and protons repelling each other as they are all positively charged), surrounded by a cloud of electrons, in defined atomic orbitals. Isotopes are atoms of the same element having the same number or protons (atomic number) but different numbers of neutrons (hence a different atomic mass). They have the same chemical properties (since this is governed by the charge of the nucleus and the electronic configuration, but different physical properties. We’ve already seen the notation commonly used has the mass number, atomic number and symbol arranged with the mass number superscript, the atomic number subscript in front of the chemical symbol: 12 6C Neutrons stabilise the nucleus in two ways. Firstly with the nuclear force, which binds the protons and neutrons together. Secondly the presence of neutrons helps to reduce the electrostatic repulsion of the protons for each other, by spreading them out. A minimum of one neutron is required to stabilise two or more protons in a nucleus. As the number of protons increases the ratio of neutrons to protons needed to stabilise the nucleus also increases. Whereas for light elements the ratio of A/Z approximates to 2, for everything heavier than 40Ca there are always more neutrons that protons if the nuclide is to be stable. Stability and instability Having too few or too many neutrons can render a nuclide unstable. The chart to the right plots the number of neutrons (y-axis) against the atomic number (x-axis) and also shows the half-lives of the nuclides. Unstable nuclei will attempt to reach lower energy states – and they do this by nuclear decay (as we’ve seen above). Stable isotopes Many common elements have at least two stable isotopes. For instance hydrogen has three isotopes: 1 2 3 1H 1H 1H The most common hydrogen, by the other two are also commonly called deuterium and tritium. It is often common to use only the superscript to denote the isotope – 2H and 3H. Chemical and molecular properties of isotopes Different isotopes for a given element share the same number of electrons and the same electronic structure. As the chemical behaviour is largely governed by the electronic structure, different isotopes exhibit virtually identical chemical behaviour (reactivity). There is one main exception to this – the kinetic isotope effect. Bond energies are affected by the mass of the atoms that form the bond – heavier isotopes tend to react slower than lighter isotopes of the same element. The most pronounced differences occur for hydrogen. It is not hard to see why – a deuterium has twice the mass of a proton. Thus if a bond involving either a proton or deuterium is made or broken during a reaction then there will be significant (and measurable) differences. For heavier atoms the difference between the isotopes is much less, so the any potential kinetic isotope effect is much smaller, so the effects on reactions are usually negligible. A physical effect of having deuterium rather than proton is observed in the IR (and/or Raman) spectra – as the bond energy is also affected by the masses involved. Isotopic labelling One application of isotopes is in drug metabolism. Since changing a normal nuclide for a stable isotope in a drug molecule is expected to have a negligible effect on the chemistry of the drug, this approach can be used to investigate how the drug is metabolised in the body. Commonly used isotopes are carbon-13 (13C), carbon-14 (14C), deuterium (D or 2H) and tritium (T or 3H). The isotopes are usually introduced into the drug during chemical synthesis. (Synthetic routes will be designed to take advantage of commercially available isotopically labelled building blocks). Administration of the drug is followed by studies to see where the label ends up. For instance, carbon-13 may end up as 13CO2 and could be detected by mass spectrometry as the subject exhales. Alternatively the label may end up as a metabolite that is excreted in urine. Several methods are used to identify the labelled metabolites, of which mass spectrometry is probably the most widely used and most sensitive. Mass spectrometry is able to discriminate between molecules that differ by just one mass unit, such as either containing a proton or a deuterium. Radiolabelling If instead of a stable isotope, a drug, or radioactive tracer, is labelled with an unstable radionuclide, then the resulting radiation can be used to follow the tracer in the body. The most common application of this is called Positron emission tomography, and usually makes use of fluorine-18, a positron emitting radionuclide. The fluorine-18 is introduced into the body on a biologically active molecule, typically fludeoxyglucose (FDG). The resulting positron emissions are all directions, and are observed all round the body, then reconstructed into a 3D image of the patient with varying concentration of the fluorine-18 source. 18F has a half-life of 109.8 minutes. 18F is usually prepared by proton bombardment of 18O enriched water, leading to 18F- ions in the water. The short half life requires rapid chemistry (automated) from this point. 18F- ions can be isolated using ion- Fludeoxyglucose Glucose exchange columns, then eluted with potassium salts and a 2,2,2-cryptand. This is then used to synthesise the FDG. FDG is a glucose analogue and is rapidly taken up by high-glucose using cells. Typically these are the brain, brown adipocytes, the kidneys and cancer cells. The replacement of the 2-hydroxyl with fluorine prevents glycolysis, and thus the FDG remains in the cell. As a result the distribution of FDG is a good reflection of glucose uptake by cells in the body. The radioactivity of 18F-FDG splits into two main routes. Around 20 % is rapidly excreted through the renal system, and is presumed to be no longer attached to glucose, as glucose is not normally excreted in urine. This is why the bladder etc is usually an area of high concentration. The remainder undergoes nuclear decay via positron emission, converting to 18O (not radioactive). With the relatively short half-life of 18F, within 24 hours radioactivity will have decayed to around 1/8192 of the initial amount. In practice patients who are injected with 18F are told to avoid contact with radiation sensitive people (e.g. children, pregnant women) for at least 12 hours. Summary – key facts Nuclear stability and instability Types of radiation (and decay) Radiolabelling – mechanistic studies and imaging (using 18F)