Ch109_LectureNotes

*Humboldt State University* ® *Department of Chemistry*

Richard A. Paselk
Chem 109	General Chemistry	Spring 2009
Lecture Notes:: 23 March	© R. Paselk 2002

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Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 109

General Chemistry

Spring 2009

Lecture Notes:: 23 March

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Electronic Configurations & Periodicity , cont.

Ions: When an atom loses electrons we would expect it to lose its outermost electrons first. But which are outermost? Remember the "last added" electrons in the transition elements are in the d orbitals of the next outermost shell. The the d orbital electrons should not be the outermost electrons in an atom. Thus we will lose the s & p electrons first then the d electrons if any are present. If additional electrons are lost then we can go into the d shell. Examples:

Na⁺ = 1s² 2s² 2p⁶ 3s⁰ or [Ne] 3s⁰ (In both cases the 3s⁰ is usually not shown, I've shown it here for clarity.)
Cu²⁺ = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s⁰ 3d⁹ or [Ar] 4s⁰ 3d⁹ (In both cases the 4s⁰ is usually not shown, I've shown it here for clarity.)
Fe³⁺ = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s⁰ 3d⁵ or [Ar] 4s⁰ 3d⁵ (In both cases the 4s⁰ is usually not shown, I've shown it here for clarity.)

Symmetry considerations: It turns out that symmetry is a strong driving force in nature and symmetry considerations are a powerful tool for predicting how nature operates. This is important in predicting electronic configurations because when two electronic energy levels are close to each other, as in the 3d orbitals (highest energy in the 3 shell) and the 4s orbitals (lowest energy in the 4 shell), symmetry considerations can result in an electron preferring to "fill" the 3d orbital set, making it symmetrical, instead of going to the already symmetrical 4s orbital. This can be done in two ways: we can put one electron in each of the five d orbitals giving a spherical half-filled d orbital set, or we can put 2 electrons in each orbital. Examples:

Cr = [Ar] 4s¹3d⁵ instead of [Ar] 4s²3d⁴. This occurs because the s orbitals are already spherically symmetrical, whereas the d orbital set only becomes fully spherically symmetrical when all of the d orbitals are filled in the same way, in this case having one unpaired electron each. or when all of the d orbitals have two electrons each.
Cu = [Ar] 4s¹3d¹⁰ instead of [Ar] 4s²3d⁹. This occurs because the s orbitals are already spherically symmetrical, whereas the d orbital set only becomes fully spherically symmetrical when all of the d orbitals are filled in the same way, in this case all of the d orbitals have two electrons each.
Cu¹⁺ = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s⁰ 3d¹⁰ or [Ar]4s⁰ 3d¹⁰ This occurs because we are removing the outermost electron, the single electron in the 4s orbital. Note that without the symmetry filling of the 4d orbitals there would be two 4s electrons and we would then expect only Cu²⁺ as we see in the alkaline earths (Mg, Ca, etc.).
Zn²⁺ = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s⁰ 3d¹⁰ or [Ar]4s⁰ 3d¹⁰. Note that the electronic structure for the Zn(II) ion is the same as the Cu(I) ion. In this case however, Zn behaves like the alkaline earths, that is it only exhibits a 2+ ion.

Chemical Bonds

Chemical bonds are the strongest forces that exist between atoms. They are the forces that hold atoms together in molecules and atoms or ions together in solids. We will look at other weak bonds and forces later.

The two most important and common strong bond types in chemistry are ionic bonds and covalent bonds, a third bond type, found in metallic solids, will be discussed later.

Electronegativity

So how do we determine whether two atoms will form an ionic or a covalent bond? Use a new property - electronegativity (EN). Electronegativity is a periodic measure of how electrons are shared by atoms with the highest value for F and the lowest for Cs. There are a couple of ways of determining EN's:

Bond Type: So how do we use this to predict whether a bond is covalent or ionic?

Ionic Bonds

An ionic bond is the result of the electrostatic force of attraction between ions that carry opposite electrical charges, as described by Coulomb's Law:

Formation of ionic bonds. We can visualize the formation of ionic bonds as the transfer of an electron from a metal atom to a non-metal atom to form an ion pair. in vacuo:

Lewis Structures for Atoms & Ions

Lewis Dot Structures are a very simple way of modeling atoms, ions, and molecules involving the representative elements (IUPAC groups 1, 2 & 13 - 18). In a Lewis Structure the nucleus and "core" electrons (all but the outermost shell) are represented by the symbol of the element, now referred to as a "kernel." (Note that kernals are not the same as Noble gas cores for atoms in Period 4 and up because of the d electrons which are included in the kernal.) Examples:

Name	Lewis Structure	Kernel electrons	Valence electrons
Sodium	Na^.	1s² 2s² 2p⁶	3s¹
Phosphorus		1s² 2s² 2p⁶	3s² 3p³
Bromine		1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ (≠ [Ar] = 1s² 2s² 2p⁶ 3s² 3p⁶ )	4s²4p⁵

Sodium Demonstration

© R A Paselk

Last modified 26 March 2009