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Does anyone have any innovative ways of teaching electron shielding

I am trying to teach patterns of reactivity to BTEC (vocational based learning) in groups 1 and 7.

Meenhaz, I teach a remedial HS level chemistry class to community college students. Because I want my students to be able to "think chemically", I find the concept of shielding electrons useful. I helps then reason through the fact that atomic radius decreases across a period as the effective nuclear charge increases. I develop this idea through the following sequence: Basic atomic structure Electron configuration (I do NOT use that little diagonal chart but rather teach then to use the period table to figure this out. The periodic table is your chemical Rosetta stone and you need to learn the "code". I make a very clear distinction between shells (corresponding to principal quantum number) and orbitals (subdivisions of shells). I do introduce them to the concept of quantum numbers but do not use the exact technical terminology. We talk about "exceptions occurring due to less energy separating orbitals as one gets father from the nucleus and the extra energy need to spin pair in a single orbital (the bus filing analogy works well - no one sits next to anyone on a public bus until they have to) Once we have electron configuration down we move to periodic trends. Here is where the idea of shielding is very useful I more or less follow the progression at the Clackamass Community College Website (not where I teach) http://dl.clackamas.cc.or.us/ch104-06/shielding_electrons.htm When I do this well, my students ask very insightful questions. (a sort of formative assessment)

I know of a teacher who used an analogy that the nucleus was a like a guy with a lot of cash (1 proton = 1 millions dollars) and the electrons represented gold digger women, with the energy levels representing the amount of access they had to the guys cash. First level girl meant nearly full access to the money, whereas second level girl meant you only got money after the first level girls had what they wanted, and third level girls had to wait for the first two levels to get what they wanted. Hence the outer level girls were "shielded" from the cash by inner level girls. A guy with 11 millions dollars is more attractive than a guy with 10 million dollars if you are a first level girl, but not if you have to be a third level and are shielded from the money by other girls. How much competition you were going to put up with before looking for a better guy depended on how much money the guy had. This analogy can be extended a number of ways. I have been hesitant to introduce it to my students, for obvious reasons. But you asked for something different, and here it is.

Eric Since the energies of electrons are quantized none of those girls are going to get access to the cash (probability of being in the nucleus for an electron is essentially nil). I am wondering how quantum levels were handled in this analogy? Thanks Pam

Well, every model has its limits. I suppose you could try to extend this to quantum states by saying that within every tier of girlfriend there were subtiers, with s girls getting first dibs on cash, followed by p girls, then d girls. But I think the analogy was originally devised just for shielding, the idea being that girls aren't going to be attracted to a guy who is giving attention to other girls, especially if those girls are getting more attention than they are. Electrons are more attracted to a sodium nucleus than to a fluorine nucleus, but when they have to compete for the space with other electrons, that changes everything. Feel free to develop the model beyond shielding and let me know how it works for you. As I said, I've been hesitant to employ it in my class, but if you find kids are connecting with it, then maybe I'll give it a go.

Eric " Electrons are more attracted to a sodium nucleus than to a fluorine nucleus," The eletronegativity of sodium is 0.93 and of fluorine 3.98 meaning that it is MUCH harder to pull an electron away from fluorine than from sodium. Sodium easily ionizes by loosing an electron to form a positive ion. Fluorine on the other hand gains an electron to form a negative ion. Fluorine is in fact the most electronegative element in the periodic table. Electronegativity increases across a row because the atomic radius decreases. With greater nuclear charge the electrons in the valence shell are more strongly held. Since one is adding electrons to the valence shell there is NO increase in shielding by core electrons as one moves across a row. Then dropping down a row the atomic radius increases because one is now adding electrons to the next valence shell.

I think of electronegativity as an atom's ability to attract electrons towards itself when a compound has been formed and ionization energy as the measure of how hard it is to strip an electron away. That said, fluorine's ionization energy is much higher than sodium's. My point was more about how a stronger nuclear charge doesn't always mean a stronger ability to hold on to all of your electrons because of the shielding effect. My comment about the sodium nucleus attracting atoms better was meant to convey that if a sodium nucleus and a fluorine nucleus both have the same number of electrons on them, the electrons on the sodium atom will be held more tightly.

I teach this to my AP Chemistry students, and they grapple with it pretty well. We do it when we talk about anti-bonding and bonding levels. Anyhow, my students have learned that when all else fails, "blame shielding." They are debating getting a t-shirt made for the class that says "when all else fails, its shielding" on one side, and "remember your organic chemistry, save you it can" on the other.