Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 109	General Chemistry	Summer 2002
Lecture Notes:: 2 July	© R. Paselk 2002
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Atomic Structure & Chemical Periodicity, cont.

Electronic Configurations & Periodicity, cont.

• 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.

A Quantum Picture of the Atom

We've taken a brief look at the physics underlying atomic structure, focusing on Schrödinger's Equation and the wave picture of electron distribution in atoms. Let's flesh this out a bit.

What we need to explain is the energy distribution of electrons in atoms and how this correlates with atomic properties. First recall the line spectrum of hydrogen and the Bohr model. We are going to keep the concepts of ground state and quantized energy levels from Bohr, after all they worked very well for Hydrogen. But we will need to build a new structure which will give these same predictions but with other factors which explain the details of hydrogen's spectra as well as other atoms. We'll again start by modelling hydrogen.

Electronic Energy Levels:

• We will designate the primary energy level, corresponding to the average radial distance of the electron from the nucleus as a shell, and give it the symbol n. The lowest possible energy level is then the ground state with n = 1.
• The value of n also gives the number of nodes in each of the orbitals in that shell, with each shell having one node at infinity, where:
  • A node is a region of zero probability of finding an electron.
  • Nodes can have two general geometries:
    • radial (or spherical, since they describe a spherical shell at a specific radial distance from the nucleus), with each atom having at least one radial (spherical) node at infinity;
    • angular (either planar, e.g. as in the planar p-node and diagonal d-nodes, or cone shaped, e.g. as in the cone-shaped nodes of the d_z² orbitals which results in the donut shaped orbitals).
• Shells with n > 1 may have subshells which are different geometrical patterns of electron distribution. Thus:
  • The lowest energy pattern is spherical and given the designation s.
  • The next lowest energy distribution is bi-lobed with a planar symmetry. It is given the designation p.
  • The third lowest energy distribution has diagonal planes of symmetry and is designated d.
  • The fourth lowest energy distribution is designated f. This is the highest subshell type occupied by ground state electrons in any atom, so we will not look any further (an infinite number of subshells exist in theory for excited states, but they are not important to our understanding).
• The average energies of the different subshells are the energy of the shell, thus when subshells are present the energy of the shell is split. For example, in the n=2 shell the 2s orbital becomes lower in energy than the shell, while the 2p orbitals become higher in energy.
• The regions of electron occupancy in subshells are called orbitals.
  • For each shell there is one s orbital.
  • For each shell with n = 2 or greater there are three p orbitals: p_x, p_y, and p_z.
  • For each shell with n = 3 or greater there are five d orbitals: d_xz, d_yz, d_xy, d_{x²- y²}, and d_z²

Atomic Orbitals Supplement

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.

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:

• Example: 2 Na + Cl₂ Æ NaCl.
  • Na + energy Æ Na⁺ + e^-
  • Cl + e^- Æ Cl^- + energy
  • Na⁺ + Cl^- Æ NaCl

• Crystal Structure (model)

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." Examples:

Name  Lewis Structure Core electrons  Valence electrons Sodium  Na.   1s2 2s2 2p6  3s1  Phosphorus    1s2 2s2 2p6  3s2 3p3  Bromine    1s2 2s2 2p6 3s2 3p6 3d10 4s24p5

For ions the charge is always shown. Thus for metal ions such as calcium the Lewis Structure simply becomes the symbol for the ion. For negative ions such as we see for oxygen (2-) we enclose the ion and its electrons in brackets to indicate that the electrons are all "owned" by the oxygen - it does not share. Notice that the Lewis Structures of monoatomic ions are isoelectronic with the nearest Noble gas. Thus Li loses an electron to leave a kernel isoelectronic with helium, whereas bromine gains an electron to become isoelectronic with Kr. Examples:

• Sodium ion: Na⁺
• Bromide ion: We need the brackets to show that the bromide ion "owns" all of the electrons rather than sharing them.

Brackets are particularly important when we make an ionic compounds:

• Sodium chloride
• Potassium bromide
• Aluminum chloride

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© R A Paselk

Last modified 11 July 2002