Chapter 1.2: The Table and Periodic
1.2.2 The First Periodic Table
The first table was made my Dimitiri Mendeleev.
Arrangement is based on physical properties such as density and reactivity.
In this chapter we learn about the periodic trends in terms of atomic size, ionization energy,
electron affinity and electronegativity.
PERIODIC LAW: when the elements are arranged in order of increasing atomic mass, certain sets of
properties recur periodically.
1.2.3 The Modern Periodic Table
Elements listed in order of increasing atomic number
They group is divided up into s block, d block, p block and f block.
This means that the last electron (valence electron) will be in the s,p,d or f orbital.
1.2.4 What are the Main Group Elements? (main group elements)
They have either one valence electron (group 1 alkali metals) or 2 valence electrons (group 2
alkaline earth metals).
All elements in the same group have an identical valence electron configuration, the only
difference is that each successive heavier elements has a higher principal quantum number (n).
Therefore group 1 has an ns1 while group 2 has an ns2 valence configuration.
N is simply the row number.
P block elements are from group 13 to group 18
Group 13 is trieles
Group 14 is tetrels
Chalcogens are group 16
Halogens are group 17
Noble gasses are group 18
1.2.5 What are the Transition Elements?
The other 2 sections are known as the d block and the f block elements. These are the transition
elements and the inner transition elements.
They have d or f orbitals GROUP 3 – 12 are the d block elements. This is 10 groups wide because the outermost energy
holds 10 d electrons.
2 f block rows have a maximum of 14 electrons. WE DON’T NEED TO WORRY ABOUT THIS IN
1.2.6 How do we write Electron Configurations using the Periodic Table
We can use the periodic table to help us write the electron config for each element.
When we have anions or cations, take away an electron from the highest n level electron.
Fe = [Ar]4s 3d 5
1.2.7 What are Periodic Trends
Effective Nuclear Charge
Data shows that not every electron in an atom has the same electrostatic attraction with the
positively charged nucleus. The charge felt from the nucleus between electrons in the same
atom is not the same
They nuclear charge is different because inner electrons can shield the outer ones.
EFFECTIVE NUCLEAR CHARGE: the actual attractive force felt by any electrons. The symbol Z*
Z* = Z – ZPrevious noble gas
As the atomic number increase across a row in the table, Z* also increase. The higher the Z*
value, the stronger the attractive between outer electrons and the nucleus.
EXAMPLE: Determine Z* for Ca…
Ca has 20 electrons. The closest noble gas is Argon which has a total of 18 electrons.
20 – 18 = 2
Effective nuclear charge helps rationalize reactivity and may change an elements structure and
bonding ( in some cases only)
For example Be and B+ are isoelectronic (same e- config) yet they have different chemistry
because of difference in nuclear charge.
1.2.8 Bond Lengths and Size of Atoms and Ions
An atomic radius is defined as half the distance between the nuclei in a homoatamic bond. It is
expressed in pictometers (1pm = 1 x 10 m)
As Z* increases the atomic value decreases, this means that there is a inverse relationship
between the two.
As we go across a table (left to right) the effective nuclear charge is increased, and the strength
of the nuclear attraction on the outer electrons increases across a period, therefore the
electrons are pulled closer the nucleus. This means a smaller atomic radius. As we move down a group the size of the atom increase. The atoms are bigger because there is
more protons which means there are more electrons to balance the charge. These electrons are
placed in orbitals that are further and further away from the nucleus.
The atomic radius versus atomic number graph can be produced.
There are some anomalies (not valid with the general trend)
Since the d and f orbitals (transition metals) have poor shielding properties (p electrons of d
electrons are not shielded as mu