Electron Configuration

and Periodic Properties

Atomic Radii

-The atomic radius is half the distance between the nuclei of two identical atoms that are bonded together. An atom’s size is defined by its orbital’s edge.

-There are several period trends in the periodic table. The atomic radius decreases across a period because of the increasing positive charge of the nucleus. Electrons are pulled closer to the nucleus when they add to the s and p sublevels of the same main energy level. This results in an increase of atomic radii, so the difference in the radii becomes smaller in between neighboring atoms in each period.

-There are also group trends in the periodic table. The atomic radii increases as you go down a group. When electrons occupy sublevels in higher main energy levels that are farther away from the nucleus, the atoms’ sizes increase.


Ionization Energy

A+energy ->A++e-, where A is an atom of any element in the periodic table, and A+ is an ion of that element.

ion: an atom or group of bonded atoms with a positive or negative charge.

ionization: processes that result in the formation of ions.

ionization energy: the energy required to form a neutral atom of an element.

If there is enough energy supplied in an atom, electrons can be removed.



Electron Affinity

Valence Electrons

  • Valence Electrons are when electrons are available to be lost, gained, or shared in the formation of chemical compounds.

  • They are often located in incomplete filled main-energy levels.

  • Main- group elements valence electrons are located in the outermost s and p.

Group number 1 2 13 14 15 16 17 18

Valence electrons 1 2 3 4 5 6 7 8



Electronegativity

  • America’s most famous chemist Linus Pauling came up with a scale electronegativity which is the tendency of an atom or radical to attract electrons in the formation of an ionic bond.

  • Electronegativity tends to increase across each period and decrease or stay the same down a group. But there are always exceptions.

  • Noble gases are usually part of the exceptions because they do not form compounds.



Periodic Properties of the d- and f-Block Elements

  • d- block has 0 to 2 electrons and s orbital has 1 to 10 electrons in the d sublevel.

  • Because ns sublevel can then interact with their surroundings, electrons in the d sublevel that is incompletely filled are the ones that have the characteristic properties.

  • Atomic radii decreases across the periods but this is generally a lot less decrease the the main-group elements.

  • d-block elements can decrease slightly and then increase forming a dip because of their electron reactions.

  • f- block elements (period 6) have a fall in some element due to atomic number increase. The radii is not a rule following pattern and is all over the place for f- block elements.

  • Ionization energies of d-block and f-block elements increase across the periods.

  • This happens because electrons are available for ionization in the outer sublevels and are less shielded from the increasing nuclear charge in the incomplete (n-1)d sublevels.

  • The highest occupied sublevels always remove electrons first in d and f blocks.

  • d sublevels, the first electrons to be removed are in the outermost sublevels.

  • Most d-block elements commonly form 2+ ions in compounds.

  • d- block elements follow general trend for electronegativity values to increase as radii decrease.

  • d-block electronegativities range from 1.1 and 2.54 and f-block 1.1 to 1.5.


Ionic Radii

Positive ions are called cations and negative ions are called anions. Within a period, metals on the left tend to form cations and nonmetals on the upper right usually form anions. Cationic radii decreases across a period and anionic radii decreases across each period for elements in Groups 15-18. This is because cationic radii decrease from left to right across each period. In a group, there is an increase of ionic radii going down.