Name: color and spin in coordination complexes
Uploaded: 2020-09-24T15:06:53.000Z
Duration: 3 min 2 s
Channel: JoVE

过渡金属络合物表现出 各种不同的颜色，归因于这些化合物对特定波长 可见光的吸收。当光具有将电子从低能级 激发到高能级所需的能量时，它就会被吸收。因此，过渡金属络合物 通常吸收与络合物的晶体场分裂能 德耳塔 匹配的光，通常在可见光范围内。例如，六氟钴酸盐(III)强烈吸收红光，但吸收绿光最少，导致它呈现绿色。德耳塔 较高的六氨合钴(III)强烈吸收高能蓝光，但吸收黄光最少。因此，六氨合钴(III)显示为黄色。较小的德耳塔 对六氟钴酸盐(III)的影响 不仅限于它的颜色。当德耳塔 足够低时，例如在六氟钴酸盐(III)中，电子会先单独占据较高能量轨道，然后再与较低能量轨道配对。这里，与自旋配对能相比，德耳塔 较小 自旋配对能量是同一轨道中电子之间的 静电排斥能量。因此，电子克服德耳塔 并占据高能轨道，比克服自旋配对能 在低能轨道中配对 在能量上更可行。相反，在六氨合钴(III)中，德耳塔 大于自旋配对能。因此，正如洪德规则所预期的那样，电子在较低能量的轨道上配对，留下较高能量的轨道空着。由于这种电子分布的差异，尽管 Co(III)离子在六氟钴酸盐(III)中有四个未配对电子，但在六氨合钴(III)中没有未配对电子。据此，前者被归类为 高自旋络合物，后者被标记为低自旋络合物。一般来说，弱场配体 与小德耳塔 值相关，形成高自旋络合物，强场配体 促进高德耳塔 值，形成低自旋络合物。高自旋和低自旋络合物 可以表现出非常不同的磁性。例如，高自旋六氟钴酸盐(III)由于它的未配对电子而被磁体吸引，称为顺磁性。同时，低自旋的六氨合钴(III)被 磁铁排斥，并标记为抗磁性。

颜色和磁性

配位化合物中的颜色 
 当原子或分子以适当的频率吸收光线时，它们的电子会被更高能量的轨道激发。 对于许多主要的组原子和分子，吸收光子处于电磁光谱的紫外线范围内，这是人眼无法检测到的。 对于配位化合物， d 轨道之间的能量差通常允许吸收和发射可见范围内的光子，这被人眼视为颜色。 
 <img alt="Image1" src="/files/ftp_upload/11463/11463_Image_1.png" style="text-align: center; height: 300px;" />
 图 1. 可见光和吸光的电磁频谱。 
 电子在轨道之间转换的相对能量的微小变化可能导致吸收光的颜色发生剧烈变化。 因此，配位化合物的颜色取决于许多因素，例如： 
 • 不同的水性金属离子具有不同的颜色。 
 •  一种金属的不同氧化状态会产生不同的颜色。 
 •  与金属中心协调的特定配体会影响配位化合物的颜色。 例如，铁 (II) 化合物 [Fe (H2O) 6]SO4 呈蓝色 - 绿色，因为高旋转化合物吸收红色波长的光子。 相比之下，低转子铁 (II) 化合物 K4[Fe (CN) 6] 由于吸收了能量较高的紫色光子，因此呈淡黄色。 
 通常，强磁场配体会导致中心金属原子的 d 轨道能量 (大 Δ) 发生大分裂。 带有这些配体的过渡金属配位化合物为黄色，橙色或红色，因为它们吸收了能量更高的紫色或蓝色光。 
 另一方面，带有弱射野配体的过渡金属 of 配位化合物通常为蓝绿色，蓝色或靛蓝，因为它们吸收了低能量的黄色，橙色或红色光。 用于分割 d 轨道的配体的强度列于光谱化学系列中。 此处的配体以晶体场分割能量 (Δ) 的值增加而书写。 
 <img alt="Image2" src="/files/ftp_upload/11463/11463_Image_2.png" style="text-align: center; height: 100px;" />
 图 2. 光谱化学系列。 
 例如， Cu+ 离子的配位化合物具有 d10 配置，所有轨道均已填充。 要使电子达到更高水平 (如 4p 轨道) ，必须使用极高能量光子。 该能量对应于光谱紫外区域极短的波长。 没有可见光被吸收，所以眼睛看不到变化，化合物显示为白色或无色。 例如，包含 [Cu (CN) 2]− 的溶液无色。 另一方面，八面体 Cu2+ 化合物在轨道等方面有一个空缺，电子可以达到这一水平。 吸收的光线的波长 (能量) 与频谱的可见部分相对应， Cu2+ 化合物几乎总是彩色的—蓝色，蓝色 - 绿色紫色或黄色。 
 配位化合物中的磁性 
 磁性测量的实验证据支持高和低自旋化合物理论。 含未配对电子的分子 (如含氧量) 为准磁性分子。 准磁性物质被磁场吸引。 许多过渡金属化合物都有未配对的电子，因此是准磁性的。 氮气等分子以及不含未配对电子的 Na+ 和 [Fe (CN)6]4− 等离子是双核。 Diamagnic 物质有被磁场排斥的轻微倾向。 
 <img alt="Image3" src="/files/ftp_upload/11463/11463_Image_3.png" style="text-align: center; height: 300px;" />
 图 3. D4 ， D5 ， d6 和 D7 系统的八面体络合物高和低自旋状态的轨道图。 这种区别不能用于 D1 ， D2 ， D3 ， D5 ， D8 ， D9 和 d10 系统。 
 当原子或离子中的电子不成对时，其旋转产生的磁力矩会使整个原子或离子成为准磁性。
 包含未配对电子的系统的磁力矩大小与这些电子的数量直接相关：未配对电子的数量越大，磁力矩越大。 因此，观察到的磁力矩用于确定存在的未配对电子数。 测得的低自旋 d6 [Fe (CN) 6]4− 的磁力矩确认铁是直径，而高自旋 d6 [Fe (H2O) 6]2+ 有四个未配对电子，其磁力矩确认了这种排列 (图 2)。 
 本文改编自<a href="https://openstax.org/books/chemistry-2e/pages/19-3-spectroscopic-and-magnetic-properties-of-coordination-compounds"> Openstax, 化学 2e, 第 19.3 节：配位化合物的光谱和磁性。 </a>

配位化合物中的颜色

当原子或分子吸收特定频率的光时，其电子会被激发至更高能级的轨道。对于……

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Transition Metals and Coordination Complexes

过渡金属与配位化合物

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KEY TERMS AND CONCEPTS

关键术语与概念

Colors and Magnetism

color and spin in coordination complexes

Color in coordination complexes comes from how metal electrons absorb light. When atoms or molecules absorb light at the right frequency, electrons move to higher-energy orbitals. For many main-group atoms and molecules, the absorbed light is ultraviolet and cannot be seen. In coordination compounds, the energy gap between d orbitals often matches visible light, so the compound can absorb and emit colors that the human eye detects.

Small changes in orbital energy can cause large shifts in the color that is absorbed. That is why coordination compounds can change color depending on the metal ion in water, the oxidation state of the metal, and the ligands attached to the metal center. Ligands are molecules or ions that bind to the metal. The iron(II) complex [Fe(H2O)6]SO4 is blue-green because the high-spin complex absorbs red light. The low-spin iron(II) complex K4[Fe(CN)6] is pale yellow because it absorbs higher-energy violet light.

Ligand strength also affects color through crystal field splitting, the size of the energy gap between d orbitals. Strong-field ligands create a large split, and these transition metal complexes often look yellow, orange, or red because they absorb violet or blue light. Weak-field ligands create a smaller split, and these complexes often look blue-green, blue, or indigo because they absorb yellow, orange, or red light. The spectrochemical series lists ligands in order of increasing crystal field splitting energy.

Some coordination compounds do not absorb visible light at all. A Cu+ complex has a d10 electron configuration, and the eg orbitals are filled. To move an electron to a higher level such as a 4p orbital, the compound would need very high-energy ultraviolet light. Since no visible light is absorbed, the compound appears white or colorless. A solution of [Cu(CN)2]- is one example. In contrast, octahedral Cu2+ complexes have a vacancy in the eg orbitals, so electrons can be excited there. These complexes are usually colored, often blue, blue-green, violet, or yellow.

Magnetism gives another way to study coordination complexes. Molecules such as O2 with unpaired electrons are paramagnetic, which means they are attracted to magnetic fields. Many transition metal complexes are also paramagnetic because they contain unpaired electrons. Molecules such as N2 and ions such as Na+ and [Fe(CN)6]4- have no unpaired electrons, so they are diamagnetic and show a slight repulsion by magnetic fields.

Magnetic measurements can support the idea of high-spin and low-spin complexes. This comparison is useful for d4, d5, d6, and d7 octahedral systems, but not for d1, d2, d3, d5, d8, d9, and d10 systems. When an electron is unpaired, its spin creates a magnetic moment, and the size of that moment depends on how many unpaired electrons are present. The measured magnetic moment can therefore help determine electron pairing. For example, low-spin d6 [Fe(CN)6]4- is diamagnetic, while high-spin d6 [Fe(H2O)6]2+ has four unpaired electrons and a magnetic moment that matches that arrangement.

Color and Spin in Coordination Complexes

properties of transition metals

Transition Metals in the Periodic Table

Transition metals are the d-block elements and the inner transition metals in the f-block. They are defined by partially filled d orbitals. In the periodic table, the transition elements are in groups 3 through 11, and the inner transition metals appear in the two rows below the main body of the table. The d-block is often grouped into transition series. The first transition series runs from Sc through Cu. The second series runs from Y through Ag. The third series includes La and Hf through...

coordination compounds and nomenclature

How Transition Metal Complexes Are Named

Transition metal complexes are named by following rules for ligands, metal ions, and oxidation state. These compounds form when a central metal ion or atom, often a transition metal or inner transition metal, accepts electron pairs from ligands. A ligand is a Lewis base that can be an atom, molecule, or ion with one or more lone pairs to donate. In these complexes, the bond is a coordinate covalent bond. That means both electrons in the bond come from the donor, such as ammonia, and are...

metal-ligand bonds

Coordination Complexes and Chelation

Coordination complexes form when a metal ion bonds to ligands. These compounds appear in familiar examples such as hemoglobin in blood, chlorophyll in green plants, vitamin B-12, and the catalyst used to make polyethylene. Transition metal ions are especially likely to form these complexes. In a coordination complex, the metal acts as a Lewis acid. The ligand acts as a Lewis base and donates both electrons in the coordinate covalent bond. A ligand can be an atom, molecule, or ion, as long as...

coordination number and geometry

Transition Metal Shapes from Coordination Number

Transition metal complexes are often identified by their coordination number and the shape around the metal ion. The coordination number is the number of donor atoms bonded to the central metal. In many complexes, the most common shapes are octahedral, tetrahedral, and square planar. Coordination number and molecular geometry are linked in a clear pattern. A coordination number of 2 gives a linear shape, as in [Ag(NH3)2]+. A coordination number of 3 gives a trigonal planar shape, as in...

structural isomerism

Linkage and Ionization Isomers

Isomers are different chemical species that share the same chemical formula. In coordination compounds, structural isomerism includes linkage isomers and coordination-sphere isomers. These two forms show how the same atoms can be arranged in different ways around a metal center. Linkage isomers form when a ligand can attach to a transition metal through two different atoms. The cyanide ligand, CN−, can bind through carbon or nitrogen. The thiocyanate ligand, SCN−, can bind through sulfur or...

stereoisomerism

Isomer Shapes in Coordination Complexes

Isomer shapes in coordination complexes can change a compound’s properties even when the chemical formula stays the same. Isomers are different chemical species with the same formula. In transition metal complexes, the atoms stay connected through the same types of bonds, but their arrangement in space can differ. One common form is geometric isomerism. In these complexes, ligands (the ions or molecules attached to the metal) can sit in cis or trans positions. The octahedral [Co(NH3)4Cl2]+ ion...

valence bond theory

Metal Ligand Bonds in Coordination Complexes

Coordination complexes form metal-ligand bonds through valence bond theory. In this model, Linus Pauling explained bonding with hybridization and overlap of atomic orbitals. The central metal atom or ion acts as a Lewis acid. It uses empty orbitals of suitable energy to accept electron pairs from ligands, which act as Lewis bases. These coordinate covalent bonds help explain the color, geometry, and magnetic behavior seen in coordination compounds. The type of hybridization helps determine the...

crystal field theory - octahedral complexes

crystal-field-theory-octahedral-complexes

Octahedral d-Orbital Splitting in Crystal Fields

Crystal field theory explains how ligands affect the d orbitals in transition metal complexes. This electrostatic model helps predict the colors, magnetic behavior, and some structures of coordination compounds. It focuses on the nonbonding electrons on the central metal ion, not on the metal-ligand bonds. In crystal field theory, the metal ion and ligands are treated as very small point charges. Because electrons are negative, electrons donated by the ligands repel the electrons in the...

crystal field theory - tetrahedral and square planar complexes

crystal-field-theory-tetrahedral-and-square-planar-complexes

Crystal Field Splitting in Non-Octahedral Shapes

Crystal field theory also explains d orbital splitting in tetrahedral and square planar complexes. These shapes do not follow the same pattern as an octahedral complex, so the d orbitals respond to ligand placement in different ways. In tetrahedral complexes, the four ligands sit between the axes. None of the d orbitals points directly at the ligands, but the d x2−y2 and d z2 orbitals lie along the Cartesian axes and overlap less with the ligands than the d xy, d xz, and d yz orbitals. The d...

colors and magnetism

Color in coordination complexes comes from how metal electrons absorb light. When atoms or molecules absorb light at the right frequency, electrons move to higher-energy orbitals. For many main-group atoms and molecules, the absorbed light is ultraviolet and cannot be seen. In coordination compounds, the energy gap between d orbitals often matches visible light, so the compound can absorb and emit colors that the human eye detects. Small changes in orbital energy can cause large shifts in the...

key terms and concepts

Learn how coordination complexes get their colors, why ligands change light absorption, and how magnetic measurements reveal unpaired electrons in transition metal complexes.

Color and Spin in Coordination Complexes | JoVE K12

chemistry

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Color and Spin in Coordination Complexes

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