Coverage of the recent problems with Japanese nuclear reactors has increased public awareness of radioactive isotopes of cesium, iodine, and uranium, but it hasn't helped people understand what makes a given isotope dangerous. It's no surprise, really; the threat posed by a particular isotope depends on a combination of factors, including its half-life, mode of decay, and what happens to the isotope once it gets inside the body. We'll look at each of these issues separately to help clear up some of the confusion.
Half-life isn't just a game
Radioactive decay is largely a random process. There is no way to predict when a specific atom will decay, but it is possible to get a sense of how often an average atom will survive before decaying. The most common measurement for this average is a half-life: the amount of time it takes for half the atoms in a sample to undergo decay. For some isotopes, the half-life is a fraction of a second; within a few seconds, nearly all of it will be gone. For other isotopes, a half-life can be hundreds of thousands of years or more, so you need a substantial amount of the material for the radiation to really register. If you only had 100,000 atoms of a long-lived isotope, chances are low that there would be any decays during a short exposure.
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