Ch107_Lec

*Humboldt State University* ® *Department of Chemistry*

Richard A. Paselk
Chem 107	Fundamentals of Chemistry	Fall 2008
Lecture Notes: 9 September	© R. Paselk 2005

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**Periodic Table of the Elements**
	IA	IIA		IIIA	IVA	VA	VIA	VIIA	VIIIA
H	He
2	Li	Be		B	C	N	O	F	Ne
3	Na	Mg	IIIB	IVB	VB	VI	VIIB	VIIIB	IB	IIB	Al	Si	P	S	Cl	Ar
4	K	Ca	Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn	Ga	Ge	As	Se	Br	Kr
5	Rb	Sr	Y	Zr	Nb	Mo	Tc	Ru	Rh	Pd	Ag	Cd	In	Sn	Sb	Te	I	Xe
6	Cs	Ba	Lu	Hf	Ta	W	Re	Os	Ir	Pt	Au	Hg	Tl	Pb	Bi	Po	At	Rn

First I want to do a brief historical and qualitative overview of the structure of the atom. This information is for your general understanding and to provide some underpinnings for future discussions. Refer to your in-class notes.

I want to talk about one key type of evidence about atomic structure (spectroscopy), and then look at how electrons are arranged in atoms using a low resolution (low level of detail) picture.

Spectra

Atomic spectra

Atoms display line spectra as seen below for Potassium (complex) and the much simpler example of Hydrogen (compared to a continuous spectra):

Spectra like these were among the most important evidence for the inner structure of atoms. The discreet lines indicate that the electrons reside in "orbitals" of discreet energies, with the colored emission lines indicating the the differences in energies between the different levels (ladder model - steps analogous to energy levels, damage from a fall analogous to "color" of light). The "bluer" the light the more energy/packet or energy/photon, the 'redder" the light the less energy/photon.

Today I want to look at a low resolution picture of the electronic structure within atoms where we will indicate the approximate relative energies. These electronic configurations indicate geometries of distribution as well, but we'll worry about that later.

For chemistry on Earth we can assume that all atoms are in their ground states. That is, each of the electrons in an atom will be at the lowest energy it can attain.

What is the basis of the periodicity of properties?

As we noted above, electrons are arranged in an atom into specific energy levels. These energy levels are called shells.

A shell indicates the average distance of its electrons from the nucleus, since higher energy electrons are more loosely held (much like planets in the Solar system, where higher energy = faster speed moves planets away from the Sun).

Within the shells electrons occupy geometrical regions of space called orbitals.

The lowest energy orbital within each shell is spherically symmetrical and is called an s-orbital.
At a slightly higher energy, but still in the same shell, are the p-orbitals. The p-orbitals show planar symmetry with two lobes. There are three p-orbitals arrayed along mutually perpendicular axis (x, y, and z), giving p_x, p_y, and p_z orbitals.

So let's look at the elements in the Periodic Table in light of this model.

The first shell holds only 2 electrons in what is called the 1s orbital. Thus helium has its first shell filled. There is no more room for electrons, so it can't react by picking up another electron. On the other hand, as a crude thought model, we can consider that each electron is held by both charges in the He nucleus, so they are much more tightly held than the electron in H, so He won't give up an electron either - its inert.
The second shell is larger (its out further from the nucleus) so holds 2 electrons in a 2s orbital, but there is now room for an additional three 2p orbitals. Thus 8 electrons can be accommodated in the second shell. Note the consequences:
- for lithium (Li) the inner two electrons of the 1s shell cancel the attraction of two of the three protons, so the outer 2s electron "sees" only a single charge. But its out further than the electrons in the 1s shell were, so its not held as strong, so Li loses its outer electron more readily than H and is more reactive.
- for Fluorine on the other side of the chart we can think of the outer shell electrons being attracted to the nucleus by 9 - 2 = 7 charges, so the last open space in an orbital will be super attractive to an outside electron, so F will be be very reactive, but in an opposite way to Li - it wants to steal electrons instead of giving them up.
- for neon all of the orbitals will be filled, and the electrons will be strongly attracted to the nucleus, so Ne will again be inert, like He above it.
The third shell is larger yet (further from the nucleus), but still crowded, so initially it can only accommodate another eight electrons.
- Of course the electrons in the 3s orbitals are even farther out from the nucleus, so we would expect Na to be even more reactive than Li, and so on for K, Rb, etc. each giving up its outermost electron more readily than the element above it in the Periodic table.
- On the other hand Cl will also attract electrons less than F, so it will be less reactive etc. for the halogens.
- So we will expect the most reactive elements to be on the opposite corners of the table - lower left and upper right.

The second Period is special, having greater tendency to form covalent bonds (share electrons). In particular the second period elements tend to from the strongest covalent bonds - C is the only element to form strong, stable, multiple covalent bonds to itself, making C based polymers possible.

Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chemical Periodicity, cont.

Trends

Electrons in Atoms

Spectra

Atomic spectra

What is the basis of the periodicity of properties?