1. Chapter 8: Periodic Properties of the Elements
Lizzie Rosenzweig

2. Periodic Property
Periodic property is a property that is predictable based on an element’s position within the periodic table
Ex. atomic radius, ionization energy and electron affinity
The arrangement of elements in the periodic table reflects how electrons fill quantum mechanical orbitals

3. The Development of the Periodic Table
Johann Döbereiner ( )
German chemist
Grouped elements into triads (threes with similar properties)
John Newlands ( )
English chemist
Grouped elements into octaves (in analogy to music notes), these were groups of eight that had similar properties

4. The Development of the Periodic Table (cont)
Dmitri Mendeleev ( )
Russian chemist
Periodic Law is when elements are arranged in order of increasing mass, and because of that, certain properties recur periodically
The elements were arranged in increasing mass from left to right so, elements with the same properties fall into the same columns
Mendeleev’s organization lead to the prediction of undiscovered elements and their properties
There were some problems with atomic mass order
Henry Moseley ( )
Showed elements arranged in atomic number

5. Electron Configuration
Electrons exist in orbitals
Electron configuration shows the particular orbitals that electrons occupy for that atom
Electrons generally occupy the lowest energy level orbitals available
Ground state is the lowest energy state

6. Electron Spin and the Pauli Exclusion Principle
An orbital diagram symbolizes the electron as an arrow and the orbital as the box
The direction of the arrow represents the electron’s spin
Electron’s spin is quantized as mₛ=½ or mₛ=-½ spinning up or down, respectively
Pauli exclusion principle: no two electrons in an atom can have the same four quantum numbers
This principle implies that each orbital can have a maximum of only two electrons, with opposing spins

7. Sublevel Energy Splitting in Multielectron Atoms
When orbitals have the same energy they are called degenerate
The energy of an orbital depends on the value of l
The energies of the sublevels are split
The lower the l value, the lower the energy of the corresponding orbital
s<p<d<f

8. Coulomb’s Law
Coulomb’s Law describes the interactions between charged particles
Coulomb’s Law states the potential energy (E) of two charged particles depends on their charges (q₁ and q₂) and on their separation
The amount of potential energy depends inversely on the separation between the charged particles

9. Conclusions from Coulomb’s Law
For like charges, the potential energy is positive and decreases as the particles get farther apart. Since systems move toward lower potential energy, like charges repel each other.
For opposite charges, the potential energy is negative and becomes and becomes more negative as the particles get closer together. Therefore, opposite charges attract each other.
The interaction between charged particles increases in size as the charges of the particles increases

10. Shielding
Shielding describes how one electron can shield another electron from the full charge of the nucleus
Shielding is the repulsion of one electron by another electrons, screening that electron from the full effects of nuclear charge
Effective nuclear charge is the actual nuclear charge experienced by an electron, defined as the nucleus plus the charge of the shielding electrons
Inner electrons shield the outer electron from full nuclear charge

11. Penetration
Penetration describes how one atomic orbital can overlap spatially with another, thus penetrating into a region that is close to the nucleus
Penetration is when an outer electron occupies the region with the inner electrons and experiences full charge

12. Electron Spatial Distributions and Sublevel Splitting
Splitting is a result of the spatial distributions of electrons within a sublevel
Because of penetration, the sublevels of each principal level are not degenerate for multielectron atoms
In the fourth and fifth principal levels, the effects of penetration become so important that the 4s orbital lies lower in energy than the 3d orbitals and the 5s orbital lies lower in energy that the 4d orbitals
The energy separations between one set of orbitals and the next become smaller for 4s orbitals and beyond

13. Electron Configurations for Multielectron Atoms
Aufbau principle is the pattern of orbital filling from lowest energy level to highest
Hund’s rule states that when filling degenerate orbitals, a single electron fills the orbitals first
Summarizing orbital filling
Electrons occupy orbitals so as to minimize the energy of the atom; therefore, lower energy orbitals fill before higher energy orbitals
One orbital has two electrons that spin in opposite directions
When orbitals of identical energy are available, electrons first occupy these orbitals singly with parallel spins rather than pairs

14. Sublevels The s sublevel has 1 orbital and can hold 2 electrons
The p sublevel has 3 orbitals and can hold 6 electrons
The d sublevel has 5 orbitals and can hold 10 electrons
The f sublevel has 7 orbitals and can hold 14 electrons

15. Practice Problems Write electron configurations for each element Mg PBr Al
1s² 2s² 2p⁶ 3s²
1s² 2s² 2p⁶ 3s² 3p³
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁵
1s² 2s² 2p⁶ 3s² 3p¹

16. Practice Problem
Write the orbital diagram for sulfur and determine the number of unpaired electrons.
1s² 2s² 2p⁶ 3s² 3p⁴
Two unpaired electrons