Lesson 1: Periodic Trends
Image courtesy of Wikipedia
The distance from the atomic nucleus to the outermost stable electron orbital in an atom that is at equilibrium.
↓
Increases down a group (column of the periodic table)
←
Increases across a period (row of the periodic table) from right to left
Therefore atomic radii are largest in the bottom left corner of periodic table. e.g. Ce is the largest, and Fr has a larger radius than He.
Why does atomic radius decrease as you go across a period? (from left to right)
→
The effective nuclear charge increases
→ therefore attracting the orbiting electrons towards the nucleus and lessening the radius. Less distance between the electrons and the nucleus so the nuclei pull is stronger.
(See bottom of this post for more information on effective nuclear charge.)
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Why does atomic radius increase as you go down a group?
↓
The addition of a new energy level (shell)
This is the energy required to remove an electron from a gaseous atom or ion.
The closer / more tightly bound an electron is to the nucleus, the more difficult it is to remove, and therefore more energy is required to remove the electron/s from the shell.
1 st Ionisation Energy – the energy required to remove one electron from the parent atom.
2 nd Ionisation Energy – the energy required to remove a second valence electron
(from the univalent ion to form a divalent ion)
3 rd
4 th etc.
More ionisation energy is required each time. So taking one electron from the outermost shell is generally easier than taking one from closer to the nucleus.
(Think of the nucleus as having an almost magnet like attraction that the electrons are drawn to, though of course this is not the literal explanation.)
This is the ability of an atom to accept electrons. It is also the energy change that occurs when an electron is added to a gaseous atom.
Increases across a period (left to right)
→
Increases up a group (column)
↑
Therefore electron affinity is highest at the top right corner of the periodic table
(excluding the noble gases)
Take a look at your periodic tables:
The Alkaline Earth Metals have a low electron affinity. Why?
This is because they already have filled ‘s’ sub shells.
Halogens have a high electron affinity (they want electrons). Why?
This is because the addition of an electron/s to an atom of one of these elements results in a filled/complete sub-shell.
By Ruth N. www.ruthlearns.wordpress.com

Noble gases – These are generally excluded and have an electron affinity near zero. Why?
This is because they possess a ‘stable octet’ of electrons so will not readily accept electrons.
This is the ability of the atom to attract electrons in a chemical bond. (Similar to electron affinity, but electronegativity measures the attraction of electrons not the ability to accept electrons.)
Increases across a period (left to right)
→
Increases up a group (column)
↑
Therefore electronegativity is highest at the top right corner of the periodic table
(excluding the noble gases)
 The higher the electronegativity the greater the attraction for bonding electrons
 Related to Ionisation Energies because low Ionisation Energies have low electronegativity.
 Example: F has the highest electronegativity, whereas Ce has very low electronegativity.
Effective Nuclear Charge
This is the charge an electron experiences after accounting for the shielding attributed to other electrons.
Thus an electron in the 2p orbital of Nitrogen has a greater Z eff
(effective nuclear charge) than an electron in the 2p orbital of Carbon.
Every time you move right across the periodic table you add one proton. i.e. you add to the nuclear charge, and thus one more electron to the valence shell. But this additional electron will not be entirely shielded from the nucleus so the S (shielding constant) increases by less than one, so the Z eff increases.
Hence:
Z eff
= Z - S (where S = shielding constant)
How does shielding affect Effective nuclear Charge?
Think of an atom of Hydrogen, it has 1 single proton surrounded by 1 electron that resides in the 1s orbital (1 st shell). Electrons are negatively charged, and
By Ruth N. www.ruthlearns.wordpress.com

protons positively charged, so the electron is attracted to the proton, held closely to the nucleus.
Then view Helium; it contains 2 protons, surrounded by 2 electrons in the 1s orbital. As well as the electron-to-proton attractions however, we must also consider the electron-to-electron relationships.
This is where repulsion becomes relevant. Because of the repulsion, electrons experience nuclear charge somewhat less than the actual nuclear charge; this is because one electron shields the other electron from the nucleus.
Effective nuclear charge is therefore the positive charge that an electron experiences, depending on how shielded it is.
Other useful links: http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Descriptive_Chemistry/Peri odic_Table_of_the_Elements/Periodic_Properties_of_the_Elements
By Ruth N. www.ruthlearns.wordpress.com