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Chemistry LibreTexts

24.5: Color and Magnetism

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  • Skills to Develop

    • To get a simple overview of the origin of color and magnetism in complex ions.

    Electromagnetic radiation is a form of energy that is produced by oscillating electric and magnetic disturbance, or by the movement of electrically charged particles traveling through a vacuum or matter. Electron radiation is released as photons, which are bundles of light energy that travel at the speed of light as quantized harmonic waves. This energy is then grouped into categories based on its wavelength into the electromagnetic spectrum and have certain characteristics, including amplitude, wavelength, and frequency (Figure \(\PageIndex{1}\)).

    Figure \(\PageIndex{1}\): Electromagnetic spectrum with light highlighted. Image used with permission (CC -SA-BY 2.0 unported; Philip Ronan).

    General properties of all electromagnetic radiation include:

    1. Electromagnetic radiation can travel through empty space, while most other types of waves must travel through some sort of substance. For example, sound waves need either a gas, solid, or liquid to pass through to be heard.
    2. The speed of light (\(c\)) is always a constant (2.99792458 x 108 m s-1).
    3. Wavelengths (\(\lambda\)) are measured between the distances of either crests or troughs.

    The energy of a photon is expressed by Planck's law in terms of the frequency (\(\nu\)) of the photon 

    \[E=h\nu \label{24.5.1}\]

    since \(\lambda \nu =c\) for all light Plancks law can be also expressed in terms of the wavelength of the photon

    \[E = h\nu = \dfrac{hc}{\lambda} \label{24.5.2}\]

    If white light is passed through a prism, it splits into all the colors of the rainbow (Figure \(\PageIndex{2}\)). Visible light is simply a small part of an electromagnetic spectrum most of which we cannot see - gamma rays, X-rays, infra-red, radio waves and so on. Each of these has a particular wavelength, ranging from 10-16 meters for gamma rays to several hundred meters for radio waves. Visible light has wavelengths from about 400 to 750 nm (1 nanometer = 10-9 meters).


    Figure \(\PageIndex{2}\): A triangular prism, dispersing light; waves shown to illustrate the differing wavelengths of light. Image used with permission (Public Domain;     Lucas V. Barbosa).

    Example \(\PageIndex{1}\): Blue Color of Copper (II) Sulfate in Solution

    If white light (ordinary sunlight, for example) passes through copper(II) sulfate solution, some wavelengths in the light are absorbed by the solution. Copper(II) ions in solution absorb light in the red region of the spectrum. The light which passes through the solution and out the other side will have all the colors in it except for the red. We see this mixture of wavelengths as pale blue (cyan). The diagram gives an impression of what happens if you pass white light through copper(II) sulfate solution.


    Working out what color you will see is not easy if you try to do it by imagining "mixing up" the remaining colors. You would not have thought that all the other colors apart from some red would look cyan, for example. Sometimes what you actually see is quite unexpected. Mixing different wavelengths of light doesn't give you the same result as mixing paints or other pigments. You can, however, sometimes get some estimate of the color you would see using the idea of complementary colors.

    Origin of Colors

    The process of absorption involves the excitation of the valence electrons in the molecule typically from the low lying level called the Highest Occupied Molecular Orbital (HOMO) into a higher lying state called the the Lowest Unoccupied Molecular Orbital (LUMO). When this HOMO and LUMO transition (Figure \(\PageIndex{3}\)) involves the absorption of visible light, the sample is colored.

    Figure \(\PageIndex{4}\): The general electron excitation process from a HOMO to a LUMO.

    The HOMO-LUMO energy difference

    \[\Delta E = E_{HOMO} - E_{LUMO} \label{24.5.3A}\]

    depends on the nature of the molecule and can be connected to the wavelength of the light absorbed

    \[\Delta E = h\nu = \dfrac{hc}{\lambda} \label{24.5.3B}\]

    Equation \(\ref{24.5.3B}\) is the most important equation in the field of light-matter interactions (spectroscopy).

    As Example \(\PageIndex{1}\) demonstrated, when white light passes through or is reflected by a colored substance, a characteristic portion of the mixed wavelengths is absorbed. The remaining light will then assume the complementary color to the wavelength(s) absorbed. This relationship is demonstrated by the color wheel shown below. Here, complementary colors are diametrically opposite each other (Figure \(\PageIndex{5}\)). Thus, absorption of 420-430 nm light renders a substance yellow, and absorption of 500-520 nm light makes it red. Green is unique in that it can be created by absorption close to 400 nm as well as absorption near 800 nm.

    Figure \(\PageIndex{5}\): The color wheel used to identify the color of species.

    Colors directly opposite each other on the color wheel are said to be complementary colors. Blue and yellow are complementary colors; red and cyan are complementary; and so are green and magenta. Mixing together two complementary colors of light will give you white light. What this all means is that if a particular color is absorbed from white light, what your eye detects by mixing up all the other wavelengths of light is its complementary color. Copper(II) sulfate solution is pale blue (cyan) because it absorbs light in the red region of the spectrum and cyan is the complementary color of red (Table \(\PageIndex{1}\)).

    Table \(\PageIndex{1}\): The Visible Spectrum
    Color Wavelength (nm) ΔE HOMO - LUMO gap (eV)
    UV 100 - 400 12.4 - 3.10
    Violet 400 - 425 3.10 - 2.92
    Blue 425 - 492 2.92 - 2.52
    Green 492 - 575 2.52 - 2.15
    Yellow 575 - 585 2.15 - 2.12
    Orange 585 - 647 2.12 - 1.92
    Red 647 - 700 1.92 - 1.77
    Near IR 700 - 10,000 1.77 - 0.12
    If the compound absorbs in one region of the spectra, it appears with the opposite (complementary) color, since all of the absorbed color has been removed. For example:
    • the material absorbs violet light ⇒ color is yellow
    • the material absorbs blue light ⇒ color is orange
    • the material absorbs yellow-green light ⇒ color is red-violet

    The Origin of Color in Complex Ions

    We often casually talk about the transition metals as being those in the middle of the Periodic Table where d orbitals are being filled, but these should really be called d block elements rather than transition elements (or metals). The definition of a transition metal is one which forms one or more stable ions which have incompletely filled d orbitals. Zinc with the electronic structure [Ar] 3d104s2 does not count as a transition metal whichever definition you use. In the metal, it has a full 3d level. When it forms an ion, the 4s electrons are lost - again leaving a completely full 3d level. At the other end of the row, scandium ([Ar] 3d14s2) does not really counts as a transition metal either. Although there is a partially filled d level in the metal, when it forms its ion, it loses all three outer electrons. The Sc3+ ion does not count as a transition metal ion because its 3d level is empty.

    Example \(\PageIndex{3}\): Hexaaqua Metal Ions

    The diagrams show the approximate colors of some typical hexaaqua metal ions, with the formula [ M(H2O)6 ]n+. The charge on these ions is typically 2+ or 3+.

    • Non-transition metal ions


    These ions are all colorless.

    • Transition metal ions


    The corresponding transition metal ions are colored. Some, like the hexaaquamanganese(II) ion (not shown) and the hexaaquairon(II) ion, are quite faintly colored - but they are colored.

    So, what causes transition metal ions to absorb wavelengths from visible light (causing color) whereas non-transition metal ions do not? And why does the color vary so much from ion to ion? This is discussed in the next sections.


    The magnetic moment of a system measures the strength and the direction of its magnetism. The term itself usually refers to the magnetic dipole moment. Anything that is magnetic, like a bar magnet or a loop of electric current, has a magnetic moment. A magnetic moment is a vector quantity, with a magnitude and a direction. An electron has an electron magnetic dipole moment, generated by the electron's intrinsic spin property, making it an electric charge in motion. There are many different magnetic forms: including paramagnetism, and diamagnetism, ferromagnetism, and anti-ferromagnetism. Only the first two are introduced below.


    Paramagnetism refers to the magnetic state of an atom with one or more unpaired electrons. The unpaired electrons are attracted by a magnetic field due to the electrons' magnetic dipole moments. Hund's Rule states that electrons must occupy every orbital singly before any orbital is doubly occupied. This may leave the atom with many unpaired electrons. Because unpaired electrons can spin in either direction, they display magnetic moments in any direction. This capability allows paramagnetic atoms to be attracted to magnetic fields. Diatomic oxygen, \(O_2\) is a good example of paramagnetism (described via molecular orbital theory). The following video shows liquid oxygen attracted into a magnetic field created by a strong magnet:


    A chemical demonstration of the paramagnetism of oxygen, as shown by the attraction of liquid oxygen to a magnet. Carleton University, Ottawa, Canada.

    As shown in the video, molecular oxygen (\(O_2\) is paramagnetic and is attracted to the magnet. Incontrast, Molecular nitrogen, \(N_2\), however, has no unpaired electrons and it is diamagnetic (this concept is discussed below); it is therefore unaffected by the magnet.

    There are some exceptions to the paramagnetism rule; these concern some transition metals, in which the unpaired electron is not in a d-orbital. Examples of these metals include \(Sc^{3+}\), \(Ti^{4+}\), \(Zn^{2+}\), and \(Cu^+\). These metals are the not defined as paramagnetic: they are considered diamagnetic because all d-electrons are paired. Paramagnetic compounds sometimes display bulk magnetic properties due to the clustering of the metal atoms. This phenomenon is known as ferromagnetism, but this property is not discussed here.


    Diamagnetic substances are characterized by paired electrons—except in the previously-discussed case of transition metals, there are no unpaired electrons. According to the Pauli Exclusion Principle which states that no two identical electrons may take up the same quantum state at the same time, the electron spins are oriented in opposite directions. This causes the magnetic fields of the electrons to cancel out; thus there is no net magnetic moment, and the atom cannot be attracted into a magnetic field. In fact, diamagnetic substances are weakly repelled by a magnetic field. In fact, diamagnetic substances are weakly repelled by a magnetic field as demonstrated with the pyrolytic carbon sheet in Figure \(\PageIndex{6}\).

    Figure \(\PageIndex{6}\): Levitating pyrolytic carbon: A small (~6 mm) piece of pyrolytic graphite levitating over a permanent neodymium magnet array (5 mm cubes on a piece of steel). Note that the poles of the magnets are aligned vertically and alternate (two with north facing up, and two with south facing up, diagonally). Image used with permission from Wikipedia.

    How to Tell if a Substance is Paramagnetic or Diamagnetic

    The magnetic form of a substance can be determined by examining its electron configuration: if it shows unpaired electrons, then the substance is paramagnetic; if all electrons are paired, the substance is diamagnetic. This process can be broken into four steps:

    1. Find the electron configuration
    2. Draw the valence orbitals
    3. Look for unpaired electrons
    4. Determine whether the substance is paramagnetic (one or more unpaired electrons) or diamagnetic (all electrons paired)

    Example \(\PageIndex{4}\): Chlorine atoms

    Are chlorine atoms paramagnetic or diamagnetic?


    Step 1: Find the electron configuration

    For Cl atoms, the electron configuration is 3s23p5

    Step 2: Draw the valence orbitals

    Ignore the core electrons and focus on the valence electrons only.


    Step 3: Look for unpaired electrons

    There is one unpaired electron.

    Step 4: Determine whether the substance is paramagnetic or diamagnetic

    Since there is an unpaired electron, Cl atoms are paramagnetic (but is quite weak).

    Example 2: Zinc Atoms

    Step 1: Find the electron configuration

    For Zn atoms, the electron configuration is 4s23d10

    Step 2: Draw the valence orbitals


    Step 3: Look for unpaired electrons

    There are no unpaired electrons.

    Step 4: Determine whether the substance is paramagnetic or diamagnetic

    Because there are no unpaired electrons, Zn atoms are diamagnetic.


    1. Pettrucci, Ralph H. General Chemistry: Principles and Modern Applications. 9th. Upper Saddle River: Pearson Prentice Hall, 2007
    2. Sherman, Alan, Sharon J. Sherman, and Leonard Russikoff. Basic Concepts of Chemistry Fifth Edition. Boston, MA: Houghton Mifflin Company, 1992. Print.