Wednesday, October 27, 2010

Periodic Trends on the Periodic Table

In class yesterday we covered the electronic structure of an atom. We Primarily focused on electron configurations and orbital diagrams (and many did very well on the simple quiz I gave midway through the period). Near the end of class I briefly talked about periodic trends, the topic I want to focus on in this blog post.

In our discussion about the periodic table we talked about how each period represents a primary energy level within an atom. Period one is the energy level closest to the nucleus while period 2,3,4, etc are further and further from the nucleus. We also talked about how each period represents a fixed energy level. This returns us back to the concepts of electromagnetic radiation we covered earlier in the chapter. Fixed energy levels were first observed in heat radiating from everyday objects. The energy levels of the measured heat were discontinuous. This same observation holds true for electrons surrounding the nucleus of an atom.

If we describe Arsenic's electron structure in words we would say it includes 4 energy shells. Its 33 electrons (ground state, stable, chargeless atom) are distributed in these 4 shells in the following way: In the first shell there is one subshell (s) and there are 2 electrons included in that orbital; the second shell contains 2 subshells (2s, 2p) and there are two electrons in each orbital. Since the p subshell has 3 orbitals there are 6 electrons included in this subshell. The third energy level contains 3 subshells (s,p,d- five orbitals) and there are 2 electrons in each orbital of these subshells (2, 6, 10). In the fourth energy level we have s,p,d and f subshells however we are running out of electrons to distribute. We put 2 electrons in the s orbital, and 3 in each of the 3 p orbitals in the fourth level.

Ok, I digressed a bit to summarize the way the PT is set up. The general trends in terms of the order of filling of atomic orbitals cause certain patterns to occur that help chemists predict how elements might behave. The three trends I mentioned in class are as follows: chemical reactivity, ionization energy, and the size of atoms and ions.

The clearest way to understand trends in chemical reactivity is to look at the outer electron structures of elements within each family or group. Each family has the same number of electrons in the highest principal energy level. Because chemical bonding occurs in this outermost energy level it makes sense that elements from the same group bond to others elements in the same ratio. The best example is the metal oxide MgO. We created MgO in the lab for our empirical formula laboratory. By calculation you figured out the smallest whole number ratio of magnesium to oxygen is 1:1. This makes sense by looking at the expected charge of magnesium based on its electron configuration: [Ne]3s2 (sorry, no superscripts here). Logically it would lose its 2s electrons and become more like Neon with a perfect noble gas formation. As if it is magic, it becomes Mg2+. And- we know oxygen forms a 2- for the very same reason- it gains to electrons easily to become more like Neon (gosh doesn't neon win the periodic table popularity contest today?)

Ionization energy is defined as the amount of energy required to remove an outer electron from the element in gas form. Naturally you can see how electron configuration greatly influences this trend. On the metallic side of the PT we have elements just dying to donate their heart and soul (not to mention their electron) while on the right side of the PT we've got stingy stooges that will protest the removal of their electrons to the hilt. In fact, they'd much rather snatch up extra electrons- the reason they so readily complex with the generous electron donors on the left side.

What element has the highest ionization energy? Helium. The figure on page 274 shows helium above all other elements in terms of first ionization energy. The electrons in the 1s orbital are closest to the nucleus and very low in overall energy. This means they are relatively stable compared to the electrons further out. Helium is more stable than hydrogen because it has a filled shell of electrons. This explains a lot about why helium is used in hot air balloons rather than hydrogen. Hardly anything can occur to make helium react- it is too stable.

In general, the amount of energy required to remove an electron increases as you move from bottom to top of the PT and from left to right. This makes sense because atoms near the bottom of the PT have outer electrons that are further from the nucleus than elements near the top. More principal energy levels means the radius of the atom is simply much larger. An electron that is further from the nucleus feels less of the positive charge of the nucleus than one that is close to it. The same thing applies as you move across the table from left to right. However, because you are adding electrons to the same principal energy level in this case the amount of energy required to remove an electron increases per electron added. Each subsequent electron feels a strong pull of positively charged nucleus. (lovingly called "effective nuclear charge" in chemistry)
The third important trend is atomic size. Judging by the information already provided would you expect the trend in ionization energy to parallel or mirror the trend in atomic size? As you move from top to bottom you add electrons to shells that are further and further from the nucleus. As the electrons get easier to remove (lower ionization energy) the overall atom gets larger and larger. Likewise, as you move left to right, however, the electrons feel a greater and greater effective charge from the nucleus. Electrons are added to the same principal energy as the number of protons in the nucleus increases. This makes the atom get smaller and smaller as the nucleus pulls the electrons in more tightly.

Congrats! You just survived my mini lesson in periodic trends. Now go read CChapter 8 and Experiment 18 for tomorrow's lab. Good luck.