Ferromagnetic, Dimagnetic and Paramagnetic

These are easily confused, here are some definitions:


Diamagnetic materials have a weak, negative susceptibility to magnetic fields. Diamagnetic materials are slightly repelled by a magnetic field and the material does not retain the magnetic properties when the external field is removed. In diamagnetic materials all the electron are paired so there is no permanent net magnetic moment per atom. Diamagnetic properties arise from the realignment of the electron paths under the influence of an external magnetic field. Most elements in the periodic table, including copper, silver, and gold, are diamagnetic.

Paramagnetic materials have a small, positive susceptibility to magnetic fields. These materials are slightly attracted by a magnetic field and the material does not retain the magnetic properties when the external field is removed. Paramagnetic properties are due to the presence of some unpaired electrons, and from the realignment of the electron paths caused by the external magnetic field. Paramagnetic materials include magnesium, molybdenum, lithium, and tantalum.

Ferromagnetic materials have a large, positive susceptibility to an external magnetic field. They exhibit a strong attraction to magnetic fields and are able to retain their magnetic properties after the external field has been removed. Ferromagnetic materials have some unpaired electrons so their atoms have a net magnetic moment. They get their strong magnetic properties due to the presence of magnetic domains. In these domains, large numbers of atom's moments (10¹² to 10¹⁵) are aligned parallel so that the magnetic force within the domain is strong. When a ferromagnetic material is in the unmagnitized state, the domains are nearly randomly organized and the net magnetic field for the part as a whole is zero. When a magnetizing force is applied, the domains become aligned to produce a strong magnetic field within the part. Iron, nickel, and cobalt are examples of ferromagnetic materials. Components with these materials are commonly inspected using the magnetic particle method.

http://www.ndt-ed.org/EducationResources/CommunityCollege/MagParticle/Physics/MagneticMatls.htm

Electronic Configuration and Quantum Numbers

The following gives a summary:

Traditional nomenclature

Many different models have been proposed throughout the history of quantum mechanics, but the most prominent system of nomenclature spawned from the Hund-Mulliken molecular orbital theory of Friedrich Hund, Robert S. Mulliken, and contributions from Schrödinger, Slater and John Lennard-Jones. This system of nomenclature incorporated Bohr energy levels, Hund-Mulliken orbital theory, and observations on electron spin based on spectroscopyand Hund's rules.

This model describes electrons using four quantum numbers, n, ℓ, m_ℓ, m_s. It is also the common nomenclature in the classical description of nuclear particle states (e.g. protons and neutrons).

• The first, n, describes the electron shell, or energy level.
  • The value of n ranges from 1 to "n", where "n" is the shell containing the outermost electron of that atom. For example, in cesium (Cs), the outermost valence electron is in the shell with energy level 6, so an electron in cesium can have an n value from 1 to 6. This is known as the principal quantum number.
• The second, ℓ, describes the subshell (0 = s orbital, 1 = p orbital, 2 = d orbital, 3 = f orbital, etc.).
  • The value of ℓ ranges from 0 to n − 1. This is because the first p orbital (ℓ = 1) appears in the second electron shell (n = 2), the first d orbital (ℓ = 2) appears in the third shell (n = 3), and so on. A quantum number beginning in 3, 0, … describes an electron in the s orbital of the third electron shell of an atom.
• The third, m_ℓ, describes the specific orbital (or "cloud") within that subshell.*
  • The values of m_ℓ range from −ℓ to ℓ. The s subshell (ℓ = 0) contains only one orbital, and therefore the m_ℓ of an electron in an s subshell will always be 0. The p subshell (ℓ = 1) contains three orbitals (in some systems, depicted as three "dumbbell-shaped" clouds), so the m_ℓ of an electron in a p subshell will be −1, 0, or 1. The d subshell (ℓ = 2) contains five orbitals, with m_ℓ values of −2, −1, 0, 1, and 2.
• The fourth, m_s, describes the spin of the electron within that orbital.*
  • An electron can have a spin of ±½, m_s will be either, corresponding with "spin" and "opposite spin." Each electron in any individual orbital must have different spins, therefore, an orbital never contains more than two electrons.

* Note that, since atoms and electrons are in a state of constant motion, there is no universal fixed value for m_ℓ and m_s values. Therefore, the m_ℓ and m_s values are defined somewhat arbitrarily. The only requirement is that the naming schematic used within a particular set of calculations or descriptions must be consistent (e.g. the orbital occupied by the first electron in a p subshell could be described as m_ℓ = −1 or m_ℓ = 0, or m_ℓ = 1, but the m_ℓ value of the other electron in that orbital must be the same, and the m_ℓ assigned to electrons in other orbitals must be different).

These rules are summarized as follows:

name	symbol	orbital meaning	range of values	value example
principal quantum number	n	shell	1 ≤ n	n = 1, 2, 3, …
azimuthal quantum number (angular momentum)	ℓ	subshell (s orbital is listed as 0, p orbital as 1 etc.)	0 ≤ ℓ ≤ n − 1	for n = 3: ℓ = 0, 1, 2 (s, p, d)
magnetic quantum number, (projection of angular momentum)	mℓ	energy shift (orientation of the subshell's shape)	−ℓ ≤ mℓ ≤ ℓ	for ℓ = 2: mℓ = −2, −1, 0, 1, 2
spin projection quantum number	ms	spin of the electron (−½ = counter-clockwise, ½ = clockwise)	−½, ½	for an electron, either: −½, ½

Example: The quantum numbers used to refer to the outermost valence electrons of the Carbon (C) atom, which are located in the 2p atomic orbital, are; n = 2 (2nd electron shell), ℓ = 1 (p orbital subshell), m_ℓ = 1, 0 or −1, m_s = ½ (parallel spins).

Source: Wikipedia

Introduction to the wave nature of electrons and orbitals

Orbitals

Wave nature of electrons

Atomic Orbitals

Atomic Orbitals (chemguide)

Electronic Structure and Atomic Orbitals

Electrons in Atoms

The Hydrogen Spectrum

The Hydrogen Spectrum

The mass spectrometer

The Mass Spectrometer

Mass Spectra Data

Periodic Table

link to Periodic Table

Periodic Table

Homework

Write up the Charle's law experiment following the IB critirea for DCP and CE.