Home Up E1 E2 E3 E4 Topics - E1 Topics E2 Topics E3 Topics E4 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 27 28 29 30 31


Periodic Trends

Atomic radii - closest 2 atoms can approach - distance between nucleus and outer electrons.

Might predict that atoms get larger as you add electrons - not entirely the case.

Notice, it is periodic!

As the primary energy level increases (going down a group), the radius increases.  Remember we said the primary energy level (n) determines the size of an orbital so this makes sense.

Why, however, does the radius decrease within a period (row) ?

1. Recall - electron  (-) charge    proton (+) charge

2. Let's go back to the Bohr model and ignore the different types of energy sub-levels (s,p,d,f)







Smaller moving across a period (left to right) due to higher nuclear charge --------->

1s electrons feel full nuclear charge and screen outer electrons


2s,p electrons "feel" screened or "effective" nuclear charge.








Going across a period (row), the same primary energy level is being filled, but the nuclear charge is getting higher.  The increased attraction to the nucleus results in them being more tightly held resulting in a smaller atomic radius. 


Elements in a given group feel the same effective nuclear charge, but as one goes down a group electrons occupy orbitals further from the nucleus in higher primary energy levels.



Larger down a group ----------------------->


All of this is related to chemistry - ionization energies

As the outer electron experiences less attractive force on it, it is easier to lose.

In fact, some lose it so easily they "explode."

Compare alkali metals:

To get them to react you have to give them something to give their electron to - H2O

Li + H2O

Na + H2O (electron further out than with Lithium, less tightly held, more reactive)

K + H2O

Na vs. Mg - higher nuclear charge of magnesium means less reactive.

Ionization of atoms

Atom -> Positive Ion + e-

Ions - atoms that lose or gain an electron and become charged

There is an energy, termed the ionization energy, associated with this process.


Ionization energy

Li --> Li+ + e- 

124 kcal = 519 kJ

Na --> Na+ + e-

118 kcal = 494 kJ

K --> K+ + e-

100 kcal = 418 kJ

Be --> Be+ + e-

215 kcal = 900 kJ

Ionization energy - energy required to remove an electron from an atom in the gaseous state.

1st ionization energy - 1st e- removed

2nd ionization energy - 2nd e- removed.


Halogens and noble gases hard to ionize!


Summary of atomic structure, its relation to the periodic table, and the properties of elements:

1. Nucleus consists of positively charge protons and uncharged neutrons.  Nucleus very small, high mass density.  Atoms are mainly empty space.

2. Electrons (much smaller than protons and neutrons) occupy orbitals of discrete energy and classified by a primary energy level (n) and a sub-level (s,p,d,f).

3. Electrons fill orbitals from the lowest energy up (Aufbau), two electrons per orbital with opposite spins (Pauli exclusion principle), and when given a choice of orbitals of the same energy singly before pairing up.

4. Groups in the periodic table share the same outer (valence) energy level electronic structure structure. 

5. Point four implies elements of the same group have similar chemical and physical properties (as empirically observed by Mendeleev for instance). Periodicity!

6. Examples: relating ionization energy (chemical property) to electronic structure, relating atomic radii (physical property) to atomic structure.