{[ promptMessage ]}

Bookmark it

{[ promptMessage ]}

nplewisdoc - ` Lewis Dot Structures Prepared By Richard E...

Info icon This preview shows pages 1–3. Sign up to view the full content.

View Full Document Right Arrow Icon
```````` 1 Lewis Dot Structures Prepared By: Richard E. Bleil, Ph.D. Introduction: What can I say? There is more to life than receiving, and more than giving. There is sharing. Yep, I'm talking about, of course, covalent compounds. In ionic bonds, one element wants the electron so much more that it literally rips that electron from the valence shell of the other. It's not quite so mean as it might sound, because truth be told, the other atom is happy to lose that electron. This is an ionic bond, wherein there is a complete transference of at least one electron from one atom (to become the cation) to another (to become the anion). But, what happens when there is no great difference between how much the two electrons want that electron? Each of them want an electron, but neither is strong enough to just take it. When this happens, the atoms are forced to share electrons. Sharing electrons forms what we call a "covalent" bond. Just like in ionic bonding, atoms share electrons to complete their valence shell. The Octet Rule: Long before quantum theory, it was noted that most elements will lose or share electrons to get eight electrons in their valence shell. Now we know this is so as they are trying to fill their "s" and "p" orbitals, but then it was called the "octet rule". Lewis Dot structures are steeped in the "octet rule", which simply states that atoms will share electrons to gain eight electrons in their valence shell. An exception to this readily recognized is hydrogen. Since we now know that hydrogen's entire electronic configuration is confined to the 1s orbital, it is not surprising that hydrogen will want only two electrons, rather than eight. Early on, it was an empirical observation. Another early empirical observation was that electrons preferred to exist in pairs. I imagine it is the observation that atoms with unpaired electrons, called "radicals", are so reactive. This observation led to the way in which we draw the Lewis dot structure for an atom. We begin by writing the atomic symbol of the atom, and count the number of "s" and "p" orbital valence shell electrons (eight maximum). We then draw one dot for each of the s and p valence shell electrons in order in each of the four quadrants around the symbol, top, right, bottom, left. Do NOT pair up electrons until we have one electron in each of these quadrants. We then continue around the circle pairing up electrons until all of these valence shell electrons are accounted for.
Image of page 1

Info icon This preview has intentionally blurred sections. Sign up to view the full version.

View Full Document Right Arrow Icon
```````` 2 Each of the four quadrants, then, represents one of the four orbitals, the one s or three p orbitals. Now it becomes a game, a puzzle if you will, to match up the electrons so that there are no unmatched electrons. The electrons with a partner are already happy, so we need worry about the unpaired electrons only. Any bond must contain two electrons, but between two atoms we can only form either single, double, or triple bonds. Since each bond must contain two electrons, a single bond involves two electrons. Each of the two
Image of page 2
Image of page 3
This is the end of the preview. Sign up to access the rest of the document.

{[ snackBarMessage ]}

What students are saying

  • Left Quote Icon

    As a current student on this bumpy collegiate pathway, I stumbled upon Course Hero, where I can find study resources for nearly all my courses, get online help from tutors 24/7, and even share my old projects, papers, and lecture notes with other students.

    Student Picture

    Kiran Temple University Fox School of Business ‘17, Course Hero Intern

  • Left Quote Icon

    I cannot even describe how much Course Hero helped me this summer. It’s truly become something I can always rely on and help me. In the end, I was not only able to survive summer classes, but I was able to thrive thanks to Course Hero.

    Student Picture

    Dana University of Pennsylvania ‘17, Course Hero Intern

  • Left Quote Icon

    The ability to access any university’s resources through Course Hero proved invaluable in my case. I was behind on Tulane coursework and actually used UCLA’s materials to help me move forward and get everything together on time.

    Student Picture

    Jill Tulane University ‘16, Course Hero Intern