Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 107

Fundamentals of Chemistry

Fall 2009

Lecture Notes: 14 October

© R. Paselk 2005


Return Exam I

Electrons in Atoms, cont.

For chemistry on Earth we can assume that all atoms are in their ground states. That is, each of the electrons in an atom will be at the lowest energy it can attain.

What is the basis of the periodicity of properties?

As we noted above, electrons are arranged in an atom into specific energy levels. These energy levels are called shells.

A shell indicates the average distance of its electrons from the nucleus, since higher energy electrons are more loosely held (much like planets in the Solar system, where higher energy = faster speed moves planets away from the Sun).

At energy levels greater than n = 1 get subshells, with the theoretical number of subshells = n. So for n = 1, one subshell, for n = 2 two subshells, etc.

Within the shells and subshells electrons occupy geometrical regions of space called orbitals.

Note that for atoms in the grond state all filled shells and subshells have spherical symmetry.

So let's look at the elements in the Periodic Table in light of this model.

The second Period is special, having greater tendency to form covalent bonds (share electrons). In particular the second period elements tend to from the strongest covalent bonds - C is the only element to form strong, stable, multiple covalent bonds to itself, making C based polymers possible.

A Quantum Picture of Atomic Orbitals

A brief introduction to orbitals is illustrated with QuickTime movies based on quantum calculations in the Atomic orbitals supplement.


Electronic Configurations

How do we figure out where electrons go in atoms? Two common methods:
  1. Follow along the Periodic Table, noting the "s-block", "p-block", and "d-block" elements in the table.
  2. Use the Aufbau (filling) diagram:
    7s 7p
    6s 6p 6d
    5s 5p 5d 5f
    4s 4p 4d 4f
    3s 3p 3d
    2s 2p
    and draw diagonal arrows at an angle such that they parallel np - (n+1)s. The order at which the arrows cross the orbitals gives the filling order.

There are a number of different notation conventions for electronic configurations.

Note that when we get to the d electrons they are added to the next inner orbital - they are added inside the atom and are not outermost! Note also that you may write them in the order they show up on the Periodic Table, or, if you prefer, you may group them by shell

The savings here is really obvious! Note that if you use the Periodic chart that comes on exams you will fill the f's before the d's of a given period (this is not the case for the wall chart).

So how do we decide which electrons are lost in ionization?

We also find a number of "special cases where filling diagrams and tools such as the Aufbau diagram fail in their predictions. What's going on? We will leave that as a mystery for other course such as Chem 109.

Periodic Table and Periodic Properties

Read Chapter 7 on the Periodic Properties of Elements

Let's look again at trends for: (use Clickers to discuss vertical, horizontal, and diagonal trends)

plot of atomic radii versus atomic number

Can be represented by the trend shown on the Periodic Table:

Periodic Table with Atomic Radii Trend arrow

plot of ionization energy versus atomic number

Can be represented by the trend shown on the Periodic Table:

Periodic Table with ionization energies Trend arrow

plot of electronegativity versus atomic number

Can be represented by the trend shown on the Periodic Table:

Periodic Table with Electronegativity Trend arrow


Two additional properties:

periodic table with highest density elements indicated by period


Plots of periodic properties ©1994 Hanson, Harper, Paselk, & Russell

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Last modified 15 October 2009