Download Lizzie Rosenzweig 8.1-8.3- Periodic Properties of the Elements

Chapter 8: Periodic Properties
of the Elements
Lizzie Rosenzweig
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
The Development of the Periodic Table
● Johann Döbereiner (1780-1849)
○ German chemist
○ Grouped elements into triads (threes with similar properties)
● John Newlands (1837-1898)
○ English chemist
○ Grouped elements into octaves (in analogy to music notes), these were
groups of eight that had similar properties
The Development of the Periodic Table (cont)
● Dmitri Mendeleev (1834-1907)
○ 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 (1887-1915)
○ Showed elements arranged in atomic number
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
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
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
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
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
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
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
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
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
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
Practice Problems
Write electron configurations for each element
1. Mg
2. P
3. Br
1.4.
2.
3.
4.
Al
1s² 2s² 2p⁶ 3s²
1s² 2s² 2p⁶ 3s² 3p³
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁵
1s² 2s² 2p⁶ 3s² 3p¹
Practice Problem
Write the orbital diagram for sulfur and determine the number of unpaired electrons.
1s² 2s² 2p⁶ 3s² 3p⁴
Two unpaired electrons