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Elements: A Visual Exploration of Every Known Atom in the Universe

Elements: A Visual Exploration of Every Known Atom in the Universe PDF

Author: Theodore Gray

Publisher: Black Dog & Leventhal


Publish Date: October 1, 2009

ISBN-10: 9781579128142

Pages: 240

File Type: PDF

Language: English

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Book Preface

Hang on tight, we’re going to explain quantum mechanics in one page. (If you find  this section too technical, feel free to skim it— there isn’t going to be a quiz at the end.) Every element is defined by its atomic  number, the number of positively charged  protons in the nucleus of every atom of that  element. These protons are matched by an  equal number of negatively charged electrons,  found in “orbits” around the nucleus. I say  “orbits” in quotes because the electrons are  not actually moving around their orbits like  planets around a star. In fact, you can’t really  speak of them as moving at all.

Instead, each electron exists as a  probability cloud, more likely to be in one  place than another, but not actually in any  one place at any given time. The figures below  show the various three-dimensional shapes of  the probability clouds of electrons around a  nucleus.

The first type, called an “s” orbital, is  totally symmetrical—the electron is not  any more likely to be in one direction than  another. The second type, called a “p” orbital,  has two lobes, meaning the electron is more  likely to be found on one side or the other of  the nucleus, and less likely to be found in any  direction in between.

While there is only one “s”-type orbital,  there are three “p” types, with lobes pointing  in the three orthogonal directions (x, y, z) of  space. Similarly there are five different types  of “d” orbitals and seven different types of “f”  orbitals, with increasing numbers of lobes.  (You may think of these shapes as a bit like  three-dimensional standing waves.) Each shape of orbital can exist in  multiple sizes, for example the 1s orbital is a  small sphere, 2s is a larger sphere, 3s is larger  still, and so forth. The energy required for an  electron to be in any given orbital increases  as the orbit becomes bigger. And all else  being equal, electrons will always settle into  the smallest, lowest-energy orbit.

So do all the electrons in an atom  normally sit together in the lowest-energy  1s orbital? No, and here we come to one of  the most fundamental discoveries in the  early history of quantum mechanics: No two  particles can ever exist in exactly the same  quantum state. Because electrons have an  internal state known as “spin,” which can be  either up or down, it turns out that exactly  two electrons can reside in a given orbital— one with spin up and one with spin down. Hydrogen has only one electron, so it sits  in the 1s orbital. Helium has two, and they  both fit into 1s, filling it to its capacity of two.  Lithium has three, and since there is no room  in 1s anymore, the third electron is forced  to sit in the higher-energy 2s orbital. And so  on— the orbitals are filled one at a time in  order of increasing energy.

Look at the Electron Filling Order  diagram on the right side of any element page  in this book, and you’ll see a graph of the  possible orbitals from 1s to 7p, with a red bar  indicating which ones are filled with electrons  (7p is the orbital of highest energy occupied  by electrons of any known element). The exact  order in which orbitals are filled turns out to  be surprisingly subtle and complex, but you  can watch it happen as you flip through the  pages of this book. Pay particular attention  around gadolinium (64)—if you think you’ve  got it figured out, your confidence might be  shaken by what happens there.

It is this filling order that determines  the shape of the periodic table. The first  two columns represent electrons filling “s”  orbitals. The next ten columns are electrons  filling the five “d” orbitals. The final six  columns are electrons filling the three “p”  orbitals. And last but not least, the fourteen  rare earths are electrons filling the seven “f”  orbitals. (If you’re asking yourself why helium,  element 2, is not above beryllium, element  4, congratulations—you’re thinking like a  chemist rather than a physicist. Eric Scerri’s  book, referenced in the bibliography, is a good  start toward answering such questions.)

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