1. Electron Configurations
And
Periodic Properties

2. Atomic Radii
Atomic radius – one-half the distance between the nuclei of identical atoms that are bonded together.
The left hand diagram shows bonded atoms.

3. As you move across the period on the periodic table the size of the atoms gets smaller.
The trend to smaller atoms across a period is caused by the increasing positive charge of the nucleus. The electrons are pulled closer to the more highly charged nucleus.

4. As you read down a group on the periodic table the size of the atoms increase.
Electrons are farther from the nucleus.

5. Ionization Energy
An ion is an atom or group of bonded atoms that has a positive or negative charge.
Sodium (Na), for example, easily loses an electron to form Na+.
Any process that results in the formation of an ion is referred to as ionization.
The energy required to remove one electron from a neutral atom of an element is the ionization energy (IE).
To compare the ease with which atoms of different elements give up electrons, chemists compare ionization energies.

6. The ionization energies of the main-group elements (the s-block and p-block elements) increase across each period.
This is due to the increasing nuclear charge.
A higher charge more strongly attracts electrons in the same energy level.

7. Among the main-group elements, ionization energies generally decrease down the groups.
Electrons are farther from the nucleus and removed more easily.

8. Electron Affinity
The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity.
Electron affinity is, essentially the opposite of the ionization energy. Instead of removing an electron from the element we add an electron to create an anion.

9. Electron affinities become more negative across a period (easier to gain electrons).
Electrons add with more difficulty going down a group.

10. Ionic Radii
A positive ion is known as a cation (formed by the loss of one or more electrons).
A negative ion is known as a anion (formed from the addition of one or more electrons).
The metals tend to form cations and the nonmetals tend to form anions.
Cationic radii decrease across a period due to increasing nuclear charge and anionic radii decreases for the elements in groups

11. There is a gradual increase of ionic radii down a group.
The outer electrons in both cations and anions are in higher energy levels as one reads down a group and are less affected by the nuclear charge.

12. Valence Electrons
The electrons available to be lost, gained, or shared in the formation of chemical compounds are referred to as valence electrons (in the outer s and p levels).
For main-group elements, the valence electrons are the electrons in the outermost s and p sublevels.

13. Electronegativity
Valence electrons hold atoms together in chemical compounds.
In many compounds, the negative charge of the valence electrons is concentrated closer to one atom than to another.
This uneven concentration of charge has a significant effect on the chemical properties of a compound.

14. Electronegativity
Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons (the most electronegative element is fluorine).
Electronegativity increases across each period.
Electronegativity decreases or stays the same down a group.

15. Writing Prompt #13 Using 5 to 10 sentences, answer the following:
Electronegativity, Ionic Radii, Electron Affinity, Ionization Energy and Atomic Radii are often referred to as Periodic Trends. Why is that an appropriate name for these properties? What does each trend explain? How does each property change across the periodic table? 1 2 3