19.3 Spectroscopic and Magnetic Properties of Coordination Compounds

• The oxygen and sulfur atoms have single pairs of electrons that can be used to coordinate to a metal center, so there are six possible donor atoms.
• Only two of these atoms can be coordinated to a metal.
• A five-member ring is formed by the coordination of one sulfur atom and one oxygen atom.

• Many scientific organizations don't like the idea of alternative medicine practitioners giving treatments for ailments that aren't related to heavy metals.

• The electroplating industry uses Ca, Fe, Zn, and Cu Ligands.
• When metal ion levels are reduced, metals can clump together to form clusters.
• The metal atoms are isolated from each other when metal coordination complexes are used.
• When plated from a bath containing the metal as a complex ion, many metals plate out as a smooth, uniform, better-looking, and more adherent surface.
• The electroplating industry uses complexes such as [Ag(CN)2]- and [Au(CN)2]- a lot.

• Scientists at Michigan State University discovered in the 1960's that there was a platinum complex that prevented cell division.
• The inhibition of cell division indicated that this compound could be an anti-cancer agent.
• The FDA approved cisplatin for use in the treatment of certain types of cancer in 1978.
• Platinum compounds have been developed for the treatment of cancer.
• The diammine is retained with other groups.

• Carboplatin, oxaliplatin, and satraplatin are newer drugs.

• The same theories used for main group element chemistry can't explain the behavior of coordination compounds.
• The observed geometries of coordination complexes are not consistent with hybridized orbitals on the central metal.
• A bonding model has been developed to explain the properties of transition metal complexes.
• Crystal field theory can be used to understand and predict the behavior of transition metal complexes.

• It allows us to understand, interpret, and predict the colors, magnetic behavior, and some structures of coordination compounds of transition metals.

• CFT focuses on the nonbonding electrons on the central metal ion in coordination complexes.
• The story of the behavior of complexes is only told by CFT.
• It tells the part that the theory does not work.
• CFT ignores any bonding between metal and ligands.
• The metal and the ligand are small point charges.

• The central metal's electrons will be repelled by the electrons donated from the ligands.
• The six ligands coordinate along the axes.

• The shaded portions show the phases of the orbitals.
• The axis labels could be shown.

• Different crystal field splittings are produced by different ligands.

• The Doct value for a metal with I- is smaller than the Doct value for a metal with cyanide.

• The like charges repel each other when two electrons occupy the same orbital.

• The electrons will always occupy each orbital.
• P is similar to Doct.

• A large Doct is produced by the strong field of six cyanide ligands.

• There is no crystal field that causes the orbitals to be degenerate.

• The weak field of the water molecule produces only a small crystal field splitting.

• The [Fe(H2O)6]3+ and [FeF6]3- ion are high-spin complexes with five unpaired electrons, but the [Fe(CN)6]3- ion is a low-spin complex with one unpaired electron.

• The complexes are made of two parts.

• The electrons will not bepaired.

• There will be one unpaired electron.

• The crystal field splitting will determine how many arepaired.

• The arrangement of electrons is dependent on the size of the crystal field splitting.

• CFT is applicable to other geometries.

• The axis system for the orbitals is shown in the diagram.

• Since CFT is based on repulsion, the orbitals closer to the ligands will be raised in energy.

• There will be 4 ion.

• The other geometry is square.
• axes become less stable.
• The theory of high- and low-spin complexes is supported by experimental evidence.
• O2 contains unpaired electrons and they are paramagnetic.
• Magnetic fields attract paramagnetic substances.
• Paramagnetic transition metal complexes have unpaired electrons.
• Molecules with no unpaired electrons are diamagnetic.
• Magnetic fields can be repelled by diamagnetic substances.

• The entire atom or ion is paramagnetic when an electron is unpaired.
• The larger the number of unpaired electrons, the larger the magnetic moment.
• The number of unpaired electrons is determined by the observed magnetic moment.
• A magnetic moment confirms the arrangement of 6 [Fe(H2O)6]2+ with four unpaired electrons.

• The electrons of the atoms are excited when they absorb light.

• The human eye can't detect the absorbed photons in the ultraviolet range of the spectrum.

• The human eye sees a mixture of all the colors as white light.

• In color vision, the colors located across from each other on a color wheel are used.
• The eye can see white light from a mixture of two contrasting colors.
• When a color is missing from white light, the eye sees its complement.
• The eyes see green when red light is absorbed.
• Lemon yellow can be seen in the eyes when violet photons are removed.

• It is white if it reflects all the light.
• An object has a color if it absorbs all colors.
• If the strip absorbs the color from white light, it becomes yellow.

• The indigo, violet, and red wavelength will be transmitted, and the complex will appear purple.

• Changes in the relative energies of the electrons can lead to changes in the color of light.
• Many factors affect the colors of coordination compounds.
• The ion can have different colors.
• Different oxidation states of one metal can produce different colors, as shown in the link below.

• The color of coordination complexes is influenced by the specific ligands coordinated to the metal center.
• The iron(II) complex [Fe(H2O)6]SO4 appears blue-green because the high-spin complex absorbs photons in the red wavelength.
• The low-spin iron(II) complex K4[Fe(CN)6]) is pale yellow because it absorbs higher-energy violet photons.

• There are 6 iron(II) octahedral metal centers.

• The colorful effect of changing oxidation states can be observed by watching this vanadium complexes.

• Transition metal coordination compounds absorb higherenergy violet or blue light.
• The coordination compounds of transition metals absorb lower-energy yellow, orange, or red light and are often blue-green, blue, or indigo.

• This energy can be seen in the ultraviolet region of the spectrum.
• No visible light can be seen, so the compound is white or odorless.
• A solution containing [Cu(CN)2]- is odorless.

• Cu2+ complexes are usually blue, green, or yellow, and the wavelength of the light absorbed corresponds to the visible part of the spectrum.

• Cu(NO3)2*5H2O is one of the brightly colored copper complexes.

• The reactivity of the transition elements varies from very active metals such as iron to almost inert elements, such as the Platinum metals.
• The type of chemistry used in the isolation of the elements from their ores depends on the concentration of the element in the ores.
• It is more difficult to reduce metals that are more active.

• Transition metals have similar chemical behavior to metals.
• They oxidize in air after heating and form halides.
• The elements that lie above hydrogen react with acids to produce salts and hydrogen gas.
• Transition metal compounds in low oxidation states are basic.
• Halides and other salts are stable in the water.
• Stable oxidation states allow transition metals to demonstrate a wide range of reactivity.

• The transition elements and main group elements can form coordination compounds, in which a central metal atom or ion is bonding to one or more ligands.
• conjugates with more than one donor atom are called polydentate ligands.
• There are two common geometries found in complexes, one with a coordination number of four and the other with a coordination number of six.

• There are also optical isomers that are mirror images but not superimposable.
• Oxygen transport in blood, water purification, and pharmaceutical use are some of the uses of coordination complexes.

• The magnetic properties of a complex can be attributed to the crystal field splitting.
• The magnitude of the splitting depends on the nature of the metal.
• Formation of high-spin complexes is favored by weak-field ligands.

• The reactions happen in a blast furnace.

• Justify your answer.

• Iron(III) can be converted to iron(II) by dichromate ion, which is reduced to chromium(III) in acid solution.
• A 2.5000-g sample of iron is dissolved and turned into iron.
• There is a requirement for Na2Cr2O7 in the titration.

• Assume air is 18% oxygen by volume.

• A 2.5624-g sample of a pure solid alkali metal chloride is dissolved in water and treated with excess silver nitrate.

• Balance the equations by predicting the products of the reactions.

• 2 is added to a solution containing hydrochloric acid.

• Balance the chemical equations by predicting the products of the reactions.

• The silver nitrate solution is slowly stirred.
• As the addition of sodium cyanide continues, a white precipitate forms.
• Chemical equations can be used to explain this observation.
• There is a similarity between silver cyanide and silver chloride.

• Predict which will be more stable.

• Give the oxidation state of the metal for each of the oxides.

• Take a look at the structures of the complexes.

Do you want to know if the following complexes have isomers?

• The isomers for [cocl5cn][cn] are drawn.

• Determine the number of unpaired electrons for [Fe(NO2)6]3- and for [FeF6]3- in terms of crystal field theory.

• Draw the crystal field diagrams.

• The solid anhydrous solid CoCl2 is blue.
• Because it absorbs water from the air, it is used as a humidity indicator to monitor if equipment has been exposed to excessive levels of water.

• Predict how many unpaired electrons this complex will have.

• The lone pair of electrons on the phosphorus atom can be donated by Trimethylphosphine.
• If trimethylphosphine is added to a solution of nickel(II) chloride in acetone, a blue compound that has amolecular mass of approximately 270 g can be isolated.
• There are no isomeric forms in this blue compound.