|Notes: 8 September
||Phone: x 5719
Let's look at the Periodic Table of the Elements again. Last
time we looked at some examples of element properties, particularly
for the alkali metals (group IA) and the halogens (group VIIA),
as well as the variation in properties for the elements of Period
3 (Na - Cl). Now let's look at the last group of elements, the
Noble Gases and then some terms. We will then look at more general
- Group VIII is known as the Noble Gases, or sometimes the
Inert Gases because until the 1960's they had no known compounds.
Very unreactive and only known compounds are with very reactive
elements like F and O, and even they don't form compounds with
smaller Noble gases such as He and Ne.
- Period: the rows of elements showing a repeating pattern
of properties (e.g. Na - Ar).
- Group: a vertical column of elements on the table sharing
a family resemblance of properties (e.g. Li - Fr).
- Representative elements: the elements of the s-block and
p-block (blue and green on the table below).
- Transition metal elements: the elements of the d-block (yellow
in the table below).
- Inner-transition metal elements: The Lanthanides and Actinides
(not shown on the table below)
- IA = alkali metals;
- IIA = Alkaline earth metals;
- VIIA = Halogens (note the generic symbol of X standing for
- VIIIA = Noble gases (older = inert gases).
You should know the terminology above and memorize the
names and symbols for the elements shown in the table below.
What is the basis of the periodicity
Periodic Table of the Elements
Electrons are held in shells.
Just for fun and your greater wisdom
- 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 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 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
- 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
- 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.
Trends: Note the trends for
- atomic size: decreases going from left Æ
right and from bottom Æ top.
- Size goes up with atomic number for any individual group.
- Size decreases irregularly as atomic number increases for
any given period (more charge pulls electrons in to nucleus,
but shielding reverses as subshells [s or p orbital sets] file.
- ionization energy: increases from left Æ
right and from bottom Æ top.
- Ionization energy goes down with atomic number for any individual
- Ionization energy increases irregularly as atomic number
increases for any given period (more charge pulls electrons in
to nucleus, but shielding reverses as subshells [s or p orbital
Lewis strucutres are a very simple model for representing some
chemical properties of atoms. For Lewis structures we assume all
atoms above Be want eight electrons in their outermost shells.
We represnt the nucleus and all inner electrons as the symbol
or "kernal." Outer electrons are then represented as
dots for individual electrons (electron bonding pairs may be represented
Lewis Structures for Atoms: Just show inner "kernel"
where symbol stands for 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. Examples:
- Sodium ion:
- Bromide ion:
- Sodium chloride:
Atoms and molecules can be held together by Strong bonds
or Weak bonds. We are first going to look at strong bonds.
Covalent vs. Ionic compounds: There are two kinds of
strong bonds: ionic bonds and covalent bonds.
- In covalent compounds atoms have a definite relationship
to each other, they are "married." Thus for water,
H2O, the smallest particle is a water molecule
containing one oxygen and two hydrogens.
- In ionic compounds ions of opposite charge attract each other,
but there is no definite attachment. Thus in sodium chloride
crystals each sodium ion is surrounded by six chloride ions and
each chloride ion is surrounded by six sodium ions and they are
equally attracted by each - there is no one-to-one relationship.
We will begin our discussion with ionic bonds since they are
easier to understand.
With the representative elements bond formation generally results
in the formation of "octets" of electrons in the outermost
Ionic bonds and Ionic Compounds:
- Formation: 2 Na + Cl2 Æ
NaCl. Can think of as composed of two equations, the oxidation
of Na and the reduction of Cl:
- Na + energy Æ Na+
- Cl + e- Æ Cl-
- Lewis Structure
- Crystal Structure (model)
- Ionic Compound Lewis Structure Examples:
- Potassium bromide
- Aluminum chloride
Covalent Compounds and Covalent Bonds:
- Formation: H2 + Cl2 Æ
2 HCl. In this case can consider that we get two equations each
involving a homo dissociation to give radicals, that is
atoms with unpaired electrons:
- H2 Æ 2 H.
- Cl2 Æ 2 Cl.
These radical then combine to form a bond with these two
electrons shared between the two atoms.
- Lewis Structure:
- Covalent Compound Lewis Structure Examples:
- Ammonium ion
- Hydrogen sulfide
- Carbon dioxide
- Carbon monoxide
© R A Paselk
Last modified 8 September 2004