|Notes: 1 November
||Phone: x 5719
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
- 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
- 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
- it wants to steal electrons instead of giving them up.
- for neon all of the orbitals will be filled (cancelling all
- 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
- 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
- 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
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
- all atoms above B (C and beyond, Z>5) want eight electrons
in their outermost shells.
- We represent the nucleus and all inner electrons as the symbol
which is referred to in Lewis structures as the "kernal"
of the atom.
- Outer, valence, electrons are then represented as
dots for individual electrons.
- 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.
- Sodium ion:
- Bromide ion:
- Sodium chloride:
© R A Paselk
Last modified 1 November 2004