Ionization Energy and Electron Affinity

IONIZATION ENERGY:

• A certain amount of energy needed to knock off the electrons from a neutral gaseous atom to form a positive ion (cation)
• Related to Atomic Size. "The bigger the atom, the farther is the outermost electron from the nucleus." 

Note: Ionization energy decreases with increasing atomic size :D 

• Atoms going across a period from left to right generally increases. The rise is due to the increasing nuclear charge (Atomic No.), although the electronsbelong to the same principal energy level.
• There are certain exceptions to this rule. The elements in Group IIIA have lower first ionization energies than those in Group IIA. The same is true for elements in Group VA and VIA.

EXAMPLES: 

Going down a column, ionization energy decreases. The ionization energy of halogens are given below: 

Fluoride (F) ----------> 1879 kJ/mole

Chloride (Cl) ----------> 1254 kJ/mole

Bromine (Br) ----------> 1139 kJ/mole

Iodine (I) ----------> 1031 kJ/mole

Astatine (At) ----------> 917 kJ/mole 

*If we would check the periodic table, we would find these in a column from up to down. This shows that from the top to the bottom, ionization energy decreases.

Going across a period, ionization energy increases. The increase is not regular but a general trend can be seen. This can be observed in the elements in Period 3 :)

Sodium (Na) ----------> 492 kJ/mole
Magnesium (Mg) ----------> 733 kJ mole
Aluminum (Al) ----------> 577 kJ/mole
Silicon (Si) ----------> 782 kJ/mole
Phosphorus (P) ----------> 1013 kJ/mole
Sulfur (S) ----------> 1004 kJ/ mole
Chlorine (Cl) ----------> 1254 kJ/mole
Argon (Ar) ----------> 1525 kJ/mole

* In period 3, it does not completely follow the rule on ionization energy. If you'd observe you could see that the I.E. increases, then decreases then increases again and decreases from left to right.
* The ionization energy decreases from one atom to the next as one goes down a group. This energy decrease is due to the fact that the valence electrons at a high energy level or a farther distance from the nucleus, which means the attraction of the electrons to the nucleus becomes smaller. It is easier to remove the electrons. 

Practice:

In each set, arrange the elements from highest ionization energy to the lowest.

A. He, Ne, Ar, Kr

B.Mg, As, Al, F
C. K, Sc, Ge
D. Zn, Cr, K, Br
E. Ti, Al,In Ga

ELECTRON AFFINITY:

•   The amount of energy release when a gaseous atom gains an electron.
• The more positive an atom's electron affinity, the greater its tendency to accept an electron to form anions. 

* A negative electoron affinity means that energy is released when an electron is added to the atom; if energy is required, the electon affinity is positive. The more negative the E.A , the greater is the tendency of the atom to attract an electron.

* Halogens have the most negative electron affinities.

ELECTRON AFFINITY TRENDS:

*The trend of the E.A is not as regular as for ionic size and ionization energy.
*Electron Affinity becomes more negative from left to right and less negative from top to bottom. This translates to stronger affinity for electrons increasing from left to right and decreasing from top to bottom. 

Note: 
 * E.A of Nobles gases have not yet been measured. They no longer show tendencies to attract electrons because they are already what we call the IDEAL Gases.

* Ionization Energy and Electron Affinity  are very important properties that determine how atoms interact and bond with each other.
* Electron Affinity is related to the formation of anions while Ionization Energy is related to the formations of cations.

Remember: Electrons with high affinities for electrons are found on the right side of the periodic table. These are the non metals. Non metals have a strong tendency to form anions
* Halogens are an exception.

Submitted By Group 5:

Mia San Juan

Alex Vergara

Jhoanne Sanchez

Elma Tejada

Cierra Mortega