Potential.<\/b> There is a rigorous, and perhaps needlessly complicated, way to define electrostatic potential. There’s a simple way too. Let’s take the simple way first:<\/p>\n
Positive charges create positive potentials and negative charges create negative potentials.<\/p><\/blockquote>\n
Suppose we have a neutral molecule. Equal number of positive and negative charges inside. If the potential near an atom is positive, the atom carries a partial positive charge. If the potential near an atom is negative, the atom carries a partial negative charge.<\/p>\n
How can atoms carry different charges? Simple. Electrons travel between atoms and they may spend more time near some atoms than others. The electron-poor<\/b> atoms carry a partial positive charge. The electron-rich<\/b> atoms carry a partial negative charge.<\/p>\n
Unfortunately, these simple rules break down for ions because an ion does not contain an equal number of positive and negative charges. The ion’s net charge determines the potential everywhere<\/i> around the ion. A positive ion creates a positive potential everywhere. An electron-poor atom in a positive ion creates really large positive potentials. These potentials are so large, in fact, that the potential is positive everywhere around the ion, even near the electron-neutral and electron-rich atoms. However, the potential is largest near the electron-poor atom so it is easily identified.<\/p>\n
OK, now the complicated definition of potential.<\/p>\n
The potential is the potential energy experienced by a +1 point charge.<\/p><\/blockquote>\n
Some folks have trouble visualizing “potential energy” so it may help to think about work. When you move a particle between positions of different potential energy, you either need to do work or the move can happen spontaneously and energy is released. Unfortunately, mentioning “work” causes its own problems so maybe its best to go back to the simple definition.<\/p>\n
Electron-deficient atoms.<\/b> This is a relative term. For example,<\/p>\n
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- if an atom is surrounded by fewer than 8 electrons in a Lewis structure, it could be regarded as “electron-deficient”<\/li>\n
- if an atom carries a positive formal charge, it could be regarded as ‘electron-deficient”<\/li>\n
- if a potential map shows that an atom carries a partial positive charge, the atom could be regarded as “electron-poor” and “electron-deficient”<\/li>\n<\/ul>\n<\/blockquote>\n
These definitions can often lead to contradictory conclusions. For example, B in BH3 is electron-deficient in the Lewis octet sense and in the potential map sense, but not in the formal charge sense. O in H3O+ is electron-deficient in the formal charge sense, but not in the octet or potential map sense. And C in H3C+ is electron-deficient in all three ways. When terminology is confusing like this, you need to pay close attention to context.<\/p>\n
Potential Map Quirks – BH3\/AlH3.<\/b> The blue zones on the BH3 and AlH3 potential maps show that these atoms are electron-poor and (partially) positively charged. Since these are neutral molecules, positive and negative charges must balance. If B and Al are electron-poor and positively charged, the 3 H must be electron-rich and negatively charged. Here are their potential maps.<\/p>\n
![\"model_2_capture.png\"](\"https:\/\/blogs.reed.edu\/chem201202\/files\/2010Lectures\/model_2_capture.png\")
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