Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 109 - General Chemistry - Spring 2011

Lecture Notes 26: 25 March

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Paramagnetism

Atoms and compounds are paramagnetic when they have unpaired electrons. Recall that electron spin can be thought of as electrons behaving as tiny magnets, and the arrows we use in orbital filling diagrams corresponding to the direction of the magnetic poles.

As a result, atoms such as oxygen and iron will be paramagnetic and be attracted to a magnetic field. For example, oxygen has two unpaired electrons as seen in the orbital filling diagram: , so it is paramagnetic. Similarly, iron is paramagnetic due to its unpaired d electrons:

Fe
	1s	2s	2p	3s	3p	4s	3d

Whereas zinc, with the d shell filled, is not paramagentic since it has no unpaired electrons:

Zn
	1s	2s	2p	3s	3p	4s	3d

Chemical Bonds, cont.

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:

E = 2.31 x 10^-19J*nm (Q₁Q₂/r)

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:

M_(g) + energy M_(g)⁺ + e^-

X_(g) + e^- X_(g)^- + energy

M_(g)+ X_(g) MX_(g)

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

• 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." (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

Lewis Structures for Ions and Ionic Compounds

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 sodium loses an electron to leave a kernel isoelectronic with neon, whereas bromine gains an electron to become isoelectronic with krypton. 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, and that the charge is distributed over the entire structure - it is not associated with any particular electron or locale.

Brackets are particularly important when we make ionic compounds:

• Sodium chloride
• Potassium bromide
• Aluminum chloride
  or Al³⁺ 3
• Barium iodide

Covalent Bonds

Covalent Bonds occur with the sharing of electrons by two atoms with similar tendencies to gain and loose electrons.

Let's look at the formation of HCl as an example of the creation of a covalent bond:

H₂ + Cl₂ 2 HCl

In this case can consider that we get two equations each involving a homo dissociation to give radicals, that is atoms with unpaired electrons:

• H₂ 2 H^.
• Cl₂ 2 Cl^.These radical then combine to form a bond with these two electrons shared between the two atoms.
• H^. + Cl^. H:Cl (This is not a proper Lewis structure, I have only shown the bonding pair of electrons.)
• Lewis Structure:

Note Lewis Structure homework set in Discussion Manual.

Lewis Structures for Covalent Molecules

Covalent Bonds occur with the sharing of electrons by two atoms with similar tendencies to gain and loose electrons.

• Covalent Compound Lewis Structure Examples:
  • Water: first write formula = H₂O, then determine central atom, O, then add electrons to get octet and draw L.S.
  • Ammonia = NH₃
  • Ammonium ion = NH₄⁺
  • Extra "practice" structures below:
    • Methane (CH₄)
    • Hydrogen sulfide
    • Arsenic triiodide
    • Hydrogen selenide
    • Oxygen difluoride

Formal Charge

This is another mode of "electron bookkeeping." Like oxidation numbers it uses a very simple set of rules to enable us to make realistic guesses about how atoms behave in molecules without having to have a pocket supercomputer to do that "quick" quantum mechanics calculation.

Formal charge helps us tp determine how charges are distributed on atoms in a molecule or molecular ion. It is not always terribly accurate, but is very useful for approximating how molecules will behave in some situations. It is particularly useful in choosing among resonance structures in organic chemistry to determine which are likely to make the greatest contribution to the "real" structure.

Formal Charge (FC) = the charge an atom would have if all bonding pairs were shared equally (polar bonds don't exist in this model).

To assign Formal Charges:

1. Draw a correct Lewis Structure.
2. Assign both electrons of a lone pair to its associated atom.
3. Divide all bonding pairs, giving one electron of each pair to each atom in the bond.
4. Calculate FC = # electrons on the unbonded (elemental) atom - # electrons assigned to the bonded atom.

Examples:

• Phosphoric acid (H₃PO₄)
  • LS:
    • FC_P: 5 - 4 = +1
    • FC_H: 1 - 1 = 0
    • FC_O: 6 - 7 = -1
    • FC_{3 O's}: 6 - 6 = 0
  • FC = 1 + 3 (0) + (-1) = 0. Notice that the Formal charges add up to give zero, the charge on the molecule.

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© R A Paselk

Last modified 28 March 2011