Humboldt State University ® Department of Chemistry

Richard A. Paselk

	Science 331
Fall 2004	Lecture/Activity	Office: SA560a
Notes: 1 November		Phone: x 5719 Home: 822-1116
		e-mail: rap1

What is the basis of the periodicity of properties?

Demo spectroscopes and gratings - rainbows vs. line spectra.

In any atom under Earth surface conditions the electrons will be in the lowest possible energy state (as close as possible to the nucleus). An atom with all of its electrons in this lowest energy state is said to be in its ground state.

In atoms, electrons are held in shells. Shells correspond to average distances of the electron from the nucleus.

• 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,
    • thus the outer 2s electron "sees" only a single charge.
    • But its out further then 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 in F 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 (cancelling all nuclear charges,
    • so no charge is visible from the outside),
    • and the electrons of the atom are strongly attracted to the nucleus (they can all "see" many proton charges),
    • so Ne will be 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 Cl will be less reactive etc.
    • and each halogen below will be less reactive.
  • So we will expect the most reactive elements to be on the opposite corners of the table - lower left and upper right.

Just for fun and your greater wisdom - The modern atom.

Lewis Structures

Lewis structures are a very simple model for representing some chemical properties of atoms. Lewis structures are generally only used for Representative elements ("A" Group elements). For Lewis structures we assume

1. all atoms above B (C and beyond, Z>5) want eight electrons in their outermost shells.
2. We represent the nucleus and all inner electrons as the symbol which is referred to in Lewis structures as the "kernal" of the atom.
3. Outer, valence, electrons are then represented as dots for individual electrons.
4. Bonding pairs of electrons may be also represented as lines to simplify structure representations for molecules.

Lewis Structures for Atoms: Here we just show the inner "kernel" where the symbol stands for the nucleus and all inner shell electrons. Only good for Representative elements. For ions the charge is always shown. Thus for metal ions such as calcium the Lewis Structure simply becomes the symbol for the ion. For negative ions such as we see for oxygen (2-) we enclose the ion and its electrons in brackets to indicate that the electrons are all "owned" by the oxygen - it does not share.

Note that for elemental ions:

• positive ions will lose the number of electrons cooresponding the their grop number,
• negative ions will have 8 - the Group number of electrons and a charge = Group number - 8.

Examples:

• Sodium:
• Sodium ion:
• Phosphorus:
• Bromine:
• Bromide ion:
• Sodium chloride:

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© R A Paselk

Last modified 1 November 2004