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I know that strontium ($\ce{Sr^90})$ has a higher mass number than its counterpart, but strontium ($\ce{Sr^86}$) is non-radioactive and not deadly to humans.

What makes $\ce{Sr^90}$ so different that it is cancerous and lethal to humans? Is it just the chemical makeup of the element that makes the difference?

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    $\begingroup$ wikipedia "While $\ce{^90Sr}$ (half-life 28.90 years) [$\beta^-$ emitter] has been used similarly, it is also an isotope of concern in fallout from nuclear weapons and nuclear accidents due to its production as a fission product. Its presence in bones can cause bone cancer, cancer of nearby tissues, and leukemia." en.wikipedia.org/wiki/Strontium See also en.wikipedia.org/wiki/Isotopes_of_strontium $\endgroup$ – MaxW Mar 25 '17 at 15:44
  • $\begingroup$ Related: chemistry.stackexchange.com/questions/20146/… $\endgroup$ – TAR86 Mar 25 '17 at 20:47
  • $\begingroup$ The big deal is that Sr is similar enough chemically to Ca to be taken up by the body. This means that the energetic betas will damage living tissue. If external to the body, or not taken up by the body, beta emitters are less of an issue (depending on the exact energy of the beta, of course). $\endgroup$ – Jon Custer Mar 26 '17 at 2:11
  • $\begingroup$ 4 neutrons$%and a MathJax comment$ $\endgroup$ – Jan Mar 29 '17 at 0:14
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For all practical purposes, Sr90 is chemically identical to Sr86. However, the extra neutrons crammed into Sr90 make it radioactive.

As mentioned in the comments above, alkaline earth metals (e.g. Ca, Sr, Ba) behave chemically in a similar fashion, and tend to deposit in bone like calcium does. This keeps the radioactive Sr90 in contact with the rapidly-growing hematopoietic material in bone, increasing the risk of cancers such as leukemia.

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