1. Periodic Properties of the Elements
Chapter 8 Notes
Periodic Properties of the Elements
Sections 8.1 – 8.9

2. I. Nerve Signal Transmission
Right now, tiny pumps in your cells are working hard to transport Na+ ions out of your cells and K+ ions into your cells. Your cells are able to differentiate between the ions due to their size: Na+ ions have a radius of 95 pm while K+ ions have a radius of 133 pm. The size of atoms and ions is a predictable property based on their location on the periodic table.
a property that is predictable based on
an element’s location on the periodic
table
Periodic Property:

3. II. The Development of the Periodic Table
The modern periodic table is primarily credited to the Russian chemist Dmitri Mendeleev ( ).
What did Mendeleev arrange the elements on
his periodic table according to?
By increasing mass
Mendeleev’s table did have predictive properties, but encountered
problems as well.
Henry Mosely ( ) corrected these problems by
organizing the periodic table according to this:
By increasing atomic number

4. Electron Configurations: Valence Electrons and the Periodic Table
We saw that Mendeleev arranged elements with similar periodic
properties in the same column. Also, notice that as you move down a
column, the number of electrons in the outermost principle energy level
stays the same. The key connection between the macroscopic world
(an element’s chemical properties) and the microscopic world (an
atom’s electronic structure) lies in the outermost electrons.

5. For multielectron atoms, not all electrons experience the same
attractive charge from the nucleus to the same degree. The force that each electron
feels from the nucleus is called the effective nuclear charge (Zeff).
The nucleus attracts electrons (because it’s positive) but other electrons
repel one another (because they are negative). This negative effect is
called shielding.
The electrons furthest from the nucleus experience the weakest effective
nuclear charge and are held least tightly within the atom; therefore, they are
the easiest/ first electrons to be lost or shared

6. Valence electrons:
Located in highest energy s & p orbitals
The e− involved in bonding / ion formation
Determine element’s chemical properties
For elements in Groups 1 – 8:
For transition metals:
# valence e─ = group #
count the highest energy s, p, and d e─
Core electrons:
inner e─ ; exist in completely filled energy
levels

7. Periodic Trends in the Size of Atoms and Effective Nuclear Charge
Atomic radius:
- used to represent the size of an atom,
assuming it is roughly spherical
- measured as half the distance between two
nuclei of bonded atoms
DOWN a column: radii INCREASE
RIGHT across a row: radii DECREASE
Trend:

9. Explaining the Trend:
Down a column: e─ occupy larger orbitals further from nucleus
so e─ are more spread out  radius increases
Right across a row: e─ enter orbitals in the same energy level
(equal distance from nucleus) ; nuclear charge also increases
(an attractive force) so e─ are being pulled closer to the nucleus
 radius decreases

10. VII. Ions: Electron Configurations, Magnetic Properties,
Ionic Radii, and Ionization Energy
Magnetic Properties
Paramagnetic:
an atom/ion that contains unpaired e─ and is
attracted by a magnetic field
Diamagnetic:
an atom/ion that contains no unpaired e─ and is
not attracted by a magnetic field
Note that an element may be diamagnetic but an ion of that element may be paramagnetic.
In order to determine if an atom/ion is paramagnetic or diamagnetic, we must look at its orbital diagram.

11. Ionic Radii
Cations:
positive ions are smaller than neutral atoms
Ca2+ smaller than Ca
WHY?
Anions:
negative ions are larger than neutral atoms
F─ larger than F
WHY?

12. Ionization Energy
Ionization energy:
energy required to remove an electron
from an atom or ion in its gaseous state
Ex. Na (g)  Na+ (g) e─ IE1 = 496 kJ/mol
The e─ is a product (it was lost)
DOWN a column: IE generally DECREASES
RIGHT across a row: IE generally INCREASES
Trend (1st Ionization Energy):

13. VIII. Ions: Electron Affinities and Metallic Character
Electron Affinity:
energy change associated with gaining
an electron in the gaseous state
Ex. Cl (g) + 1 e−  Cl − (g) EA = − 349 kJ/mol
The e─ is a reactant (it was gained)
EA’s are usually negative because energy is released when an atom or
ion gains an electron
* NO regular trend for columns / rows
* In general, nonmetals tend to have larger negative EA’s
(more readily accept e─)
Trend:

14. Metallic Character:
metals are good conductors of heat and electricity, are malleable
(hammered into sheets) and ductile (drawn into wire), and readily LOSE e─.
DOWN a column: M.C. INCREASES
RIGHT across a row: M.C. DECREASES
Trend: