Transition Metales and Coordination Compound

Name: Transition Metales and Coordination Compound
Uploaded: 2017-06-29T01:24:55+00:00
Duration: PTM35S21
Channel: Ashlee Rogers
Description: Transition Metales and Coordination Compound

Transition Metales and Coordination Compound

coordination compounds
coordination bond ionic bond [ Cu ( NH3 )4 ] SO42- Complex ion Counter ion （inner sphere）（outer sphere） coordination compounds ★ coordination molecules have no outer sphere but inner sphere

The molecules or ions that surround a metal ion in a complex are called ligands.
Ligands may be anions or polar molecules. Some common ligands are water, ammonia, hydroxide, cyanide, and chloride.

Each of these has at least one unshared pair of electrons that can be used to coordinate (form a coordinate covalent bond to the metal ion). Metal ions can act as Lewis acids, accepting electrons from ligands, which act as Lewis bases. The atom that is bound directly to the metal ion is called the donor atom. It's the atom that donates a lone pair of electrons to the metal ion.

[ Cu（NH3）4 ] 2+ [ Fe(H2O) 6]3+ [ SiF6 ] 2- Cu(Ⅱ) Fe(Ⅲ) Si(Ⅳ) NH3 H2O F- N O F

： cations or atoms offering empty orbitals
： the molecules or ions that surround the metal ion in a complex ： the atoms of ligand bound directly to the metal （central atom）（Ligand）（donor atom） ★ Ligand is from the Latin word ligare, meaning “to bind”

 Coordination sphere  Coordination number
Metal and ligands bound to it  Coordination number number of donor atoms bonded to the central metal atom or ion in the complex Most common = 4, 6 Determined by ligands Larger ligands and those that transfer substantial negative charge to metal favor lower coordination numbers

 Ligands classified according to the number of donor atoms Examples
monodentate = 1 bidentate = 2 tetradentate = 4 hexadentate = 6 polydentate = 2 or more donor atoms

Coordination Compounds
 Common ligands: water H - O - H ammonia NH3 chloride ion Cl - cyanide ion C N -

Monodentate （monodentate ligand )：
— the ligands possess a single donor atom and are able to occupy only one site in a coordination sphere Monodentate Examples: H2O, CN-, NH3, NO2-, SCN-, OH-, X- (halides), CO, O2- Example Complexes [Co(NH3)6]3+ [Fe(SCN)6]3-

— the ligands possess two donor atoms
（bidentate ligand )： — the ligands possess two donor atoms Examples oxalate ion = C2O42- ethylenediamine (en) = NH2CH2CH2NH2 ortho-phenanthroline (o-phen) Example Complexes [Co(en)3]3+ [Cr(C2O4)3]3- [Fe(NH3)4(o-phen)]3+ N - CH2- CH2 - N H Ethylenediamine (en)

oxalate ion Structures of other bidentate ligands. The coordinating atoms are shown in blue. Ortho- phenanthroline carbonate ion bipyridine

Ethylenediamine (en) N - CH2- CH2 - N H

（hexadentate ligand )： — the ligands possess six donor atoms
ethylenediaminetetraacetate (EDTA) = (O2CCH2)2N(CH2)2N(CH2CO2)24- Example Complexes [Fe(EDTA)]-1 [Co(EDTA)]-1

 Uses for EDTA Used in food products to complex metals that promote decomposition reactions Used to remove toxic heavy metals from the body Used to complex metal ions that inhibit or promote polymerization reactions

Porphine, an important chelating agent found in nature
Metalloporphyrin

Porphyrins are naturally occurring complexes derived from porphin. Heme (Fe2+) Chlorophyll (Mg2+)

Hymoglobin, a protein that stores O2 in cells

Coordination Environment of Fe2+ in Oxymyoglobin, and Oxyhemoglobin

Ferrichrome (Involved in Fe transport in bacteria)
Microorganisms secrete siderophore which forms very stable, water-soluble complexes with Fe3+. The resulting complex, ferrichrome, has a net charge of 0 which allows it to pass through the cell wall of the microorganism. In the cell, an enzyme reduces Fe3+ to Fe2+, which forms a weak complex with siderophore so it is released and utilized by the cell. Bacteria use siderophore to obtain iron from human blood, allowing them to reproduce.

Complex charge ☛ The charge on a central metal ion can be
determined in much the same way that oxidation numbers are determined in molecular compounds. It is important to recognize common molecules, to remember that they are neutral, and to know the charges on both monoatomic and polyatomic ions. The sum of charges on the metal ion and the ligands must equal the net charge on the complex.

neutral compound [Co(NH3)6]Cl2
Complex charge = sum of charges on the metal and the ligands [Fe(CN)6]3- +3 6(-1) Neutral charge of coordination compound = sum of charges on metal, ligands, and counterbalancing ions neutral compound [Co(NH3)6]Cl2 +2 6(0) 2(-1)

Tetrahedral(more common) or square planar
☛ Coordination numbers of 2, 4, and 6 are common. The geometry of a complex is determined in part by its coordination number. (Size of the ligands can also impact the geometry.) Coordination Number Geometry 2 Linear 4 Tetrahedral(more common) or square planar 6 Octahedral

Structures of geometries,
(a) [Zn(NH3)4] tetrahedral (b) [Pt(NH3)4]2+, square-planar These are the two common geometries for complexes in which the metal ion has a coordination number of 4.

Attentions: e.g：H2O、 F- e.g：SCN-、 NCS-、 NO2- 、ONO-
 In mono-core coordination compounds, each donor atom can only bind directly to one central atom to form coordination bond. e.g：H2O、 F-  Though some ligands have two or more donor atoms, they are still monodentate ligand because only one donor atom can form bond with central atom. e.g：SCN-、 NCS-、 NO2- 、ONO-

— The one formed by central metal atom and monodentate ligand.
1. Simple complex — The one formed by central metal atom and monodentate ligand. e.g: [Cu(NH3)4]SO4、K[Ag(CN)2]、 K2[PtCl4]、etc.

2. (Chelate compound) — the ringed coordination compounds formed by metal atom and polydentate Polydentate ligands tend to coordinate more readily and form more stable complexes than monodentate ligands. This phenomenon is known as the Chelate effect .

(chelating agents) — polydentate ligands are also known as chelating agents which are usually organic ligands possessing donor atoms like N、P、O、S. Cd2+

The [CoEDTA]– ion, showing how the Ethylene diamine tetraacetate ion is able to wrap around a metal ion, occupying six positions in the coordination sphere.

Nomenclature of CC (IUPAC Rules)
 The cation is named before the anion  When naming a complex: Ligands are named first alphabetical order Metal atom/ion is named last oxidation state given in Roman numerals follows in parentheses Use no spaces in complex name Pentaamminechlorocobalt (III) [Co(NH3)5Cl]2+

Ligand Name bromide, Br- bromo chloride, Cl- chloro cyanide, CN- cyano hydroxide, OH- hydroxo oxide, O2- oxo fluoride, F- fluoro  The names of anionic ligands end with the suffix -o -ide suffix changed to -o -ite suffix changed to -ito -ate suffix changed to -ato

Ligand Name carbonate, CO32- carbonato oxalate, C2O42- oxalato sulfate, SO42- sulfato thiocyanate, SCN- thiocyanato thiosulfate, S2O32- thiosulfato Sulfite, SO32- sulfito

 Neutral ligands are referred to by the usual name for the molecule Example ethylenediamine Exceptions water, H2O = aqua ammonia, NH3 = ammine carbon monoxide, CO = carbonyl

 Greek prefixes are used to indicate the number of each type of ligand when more than one is present in the complex di-, 2; tri-, 3; tetra-, 4; penta-, 5; hexa-, 6  If the ligand name already contains a Greek prefix, use alternate prefixes: bis-, 2; tris-, 3; tetrakis-,4; pentakis-, 5; hexakis-, 6 The name of the ligand is placed in parentheses e.g: [Co(en)3]Cl3－tris(ethylenediamine)cobalt(III) chloride

 If a complex is an anion, its name ends with the -ate appended to name of the metal e.g: K4[Fe(CN)6] potassium hexacyanoferrate(II) [CoCl4] tetrachlorocobaltate(II) ion

Transition Metal Name if in Cationic Complex Name if in Anionic Complex Sc Scandium Scandate Ti titanium titanate V vanadium vanadate Cr chromium chromate Mn manganese manganate Fe iron ferrate Co cobalt cobaltate Ni nickel nickelate Cu Copper cuprate Zn Zinc zincate

★ Some names of coordination compounds
[Co(NH3)4Cl2]Cl: tetraamminedichlorocobalt(Ⅲ) chloride [Pt(NH3)2Cl2]: diamminedichloroplatinum(Ⅱ) K[Ni(C2O4)2]: potassium bisoxalatonicketate(Ⅱ) oxalic acid [Cr (NH3)5H2O](NO3)3: pentaammineaquachromium (Ⅲ) nitrate

Which name is correct for the coordination compound [Fe(en)2Cl2]Cl?
Questions Which name is correct for the coordination compound [Fe(en)2Cl2]Cl? A dichlorobisethylenediamineiron(III) chloride B diethylenediaminechlorineiron(I) chloride C iron(III) bisethylenediaminedichloro chloride D iron(III) trichloridebisethylenediamine

Classes of isomers Dr.Monther F.Salem

✽ Isomerism of CC  isomer ：Two or more compounds with the same elemental composition but different arrangements of atoms. — Major Types  structural isomers  stereoisomers

Structural Isomers  structural isomers ： — isomers that have different bonds. 1. Coordination-sphere isomers differ in a ligand bonded to the metal in the complex, as opposed to being outside the coordination-sphere Example [Co(NH3)5Cl]Br vs. [Co(NH3)5Br]Cl

[Co(NH3)5Cl]Br vs. [Co(NH3)5Br]Cl
— Consider ionization in water [Co(NH3)5Cl]Br  [Co(NH3)5Cl]+ + Br- [Co(NH3)5Br]Cl  [Co(NH3)5Br]+ + Cl- — Consider precipitation [Co(NH3)5Cl]Br(aq) + AgNO3(aq)  [Co(NH3)5Cl]NO3(aq) + AgBr(s) [Co(NH3)5Br]Cl(aq) + AgNO3(aq)  [Co(NH3)5Br]NO3(aq) + AgCl(aq)

Linkage Isomers

— differ in the atom of a ligand bonded to the metal in the complex
2. Linkage isomers — differ in the atom of a ligand bonded to the metal in the complex Example: [Co(NH3)5(ONO)]2+ vs. [Co(NH3)5(NO2)]2+ Example: [Co(NH3)5(SCN)]2+ vs. [Co(NH3)5(NCS)] Co-SCN vs. Co-NCS

Stereoisomers Geometric isomers  stereoisomers
— isomers with the same bonds but different spatial arrangements of those bonds. Geometric isomers －Differ in the spatial arrangements of the ligands －Have different chemical/physical properties different colors, melting points, polarities, solubilities, reactivities, etc.

Geometric Isomers Pt(NH3)2Cl2
cis isomer, trans isomer， Pt(NH3)2Cl2

[PtCl2(NH3)2] (cis-platin),
[PtCl2(NH3)2] (trans-platin),

Geometric Isomers cis isomer trans isomer [Co(H2O)4Cl2]+

2. Optical isomers isomers that are nonsuperimposable mirror images said to be “chiral” (handed) referred to as enantiomers A substance is “chiral” if it does not have a “plane of symmetry”

Stereoisomers Other stereoisomers, called optical isomers or enantiomers, are mirror images of each other. Just as a right hand will not fit into a left glove, two enantiomers cannot be superimposed on each other.

Enantiomers A molecule or ion that exists as a pair of enantiomers is said to be chiral.

Example 1 mirror plane cis-[Co(en)2Cl2]+

Example 1 rotate mirror image 180° 180 °

Example 1 Nonsuperimposable, enantiomers cis-[Co(en)2Cl2]+

Example 2 mirror plane trans-[Co(en)2Cl2]+

Example 2 rotate mirror image 180° 180 ° trans-[Co(en)2Cl2]+

Superimposable-not enantiomers
Example 2 Superimposable-not enantiomers trans-[Co(en)2Cl2]+

No Optical activity Does have Optical activity
(a) Trans isomer of Co(en)2Cl2+ and its mirror image a are identical(superimposable) (b) cis isomer of Co(en)2Cl2+ No Optical activity Does have Optical activity

Properties of Optical Isomers
 Enantiomers possess many identical properties solubility, melting point, boiling point, color, chemical reactivity (with nonchiral reagents) different in: interactions with plane polarized light reactivity with “chiral” reagents Example

(vibrates in one plane)
Optical Isomers polarizing filter plane polarized light light source unpolarized light (random vibrations) (vibrates in one plane)

Optical Isomers (+) dextrorotatory (right rotation )
(-) levorotatory (left rotation )

optically active sample in solution rotated polarized light
polarizing filter plane polarized light optically active sample in solution rotated polarized light

optically active sample in solution rotated polarized light
polarizing filter plane polarized light optically active sample in solution Dextrorotatory (d) = right rotation Levorotatory (l) = left rotation Racemic mixture = equal amounts of two enantiomers; no net rotation rotated polarized light

Geometry of NH3

2+ NH3 NH3 Zn2+ NH3 NH3

Geometry of [Zn(NH3)4] 2+ 30Zn2+ [Ar]18 4s 4p 3d10 sp3 [Zn(NH3)4] 2+
Tetrahedron

Crystal-Field Theory When ligands coordinate to a metal ion, they increase the energies of the metal‘s d orbitals by ligand d-orbital repulsion. Because of their spatial arrangement, the d orbitals do not all experience the same increase in energy. If we consider the formation of an octahedral complex in which the ligands approach the metal ion along the x, y, and z axes, we see that only orbitals that lie along the axes (the dz2 and dx2-y2) are approached directly.

(红宝石)

Crystal-Field Theory  Crystal field theory describes bonding in transition metal complexes.  The formation of a complex is a Lewis acid-base reaction.  Both electrons in the bond come from the ligand and are donated into an empty, hybridized orbital on the metal.  Charge is donated from the ligand to the metal.  Assumption in crystal field theory: the interaction between ligand and metal is electrostatic.  The more directly the ligand attacks the metal orbital, the higher the energy of the d orbital.

- Crystal Field Theory + Octahedral Crystal Field
(-) Ligands attracted to (+) metal ion; provides stability d orbital e-’s repulsed by (–) ligands; increases d orbital potential energy ligands approach along x, y, z axes

less electrostatic repulsion = lower potential energy
greater electrostatic repulsion = higher potential energy less electrostatic repulsion = lower potential energy

Splitting of d orbital of metal atoms
 The complex metal ion has a lower energy than the separated metal and ligands.  In an octahedral field, the five d orbitals do not have the same energy: three degenerate orbitals are higher energy than two degenerate orbitals.  The energy gap between them is called , the crystal field splitting energy.

Splitting of d-orbital Energies by a Tetrahedral Field and a Square Planar Field of Ligands

Factors affecting the magnitude of ∆
effect of ligands

2. Effect of metal atoms [Co(H2O)6]3+ △o=18600 cm－1
[Co(NH3)6]3+ △o= cm－1 [Co(NH3)6]2+ △o= cm－1 [Co(NH3)6]3+ △o=23000 cm－1 [Rh(NH3)6]3+ △o= cm－1 [Ir(NH3)6] △o= cm－1

Electronic Configurations of Transition Metal Complexes
Expected orbital filling tendencies for e-’s: occupy a set of equal energy orbitals one at a time with spins parallel (Hund’s rule) minimizes repulsions occupy lowest energy vacant orbitals first These are not always followed by transition metal complexes.

 d orbital occupancy depends on  and pairing energy, P
e-’s assume the electron configuration with the lowest possible energy cost If  > P ( large; strong field ligand) e-’s pair up in lower energy d subshell first If  < P ( small; weak field ligand) e-’s spread out among all d orbitals before any pair up

d-orbital energy level diagrams octahedral complex

low spin  > P high spin  < P

Electronic Configurations of Transition Metal Complexes
 Determining d-orbital energy level diagrams: determine oxidation # of the metal determine # of d e-’s determine if ligand is weak field or strong field draw energy level diagram

High-spin and low-spin complex ions of Mn2+

High spin weak-field ligand Low spin strong-field ligand
Orbital Occupancy for High- and Low- Spin Complexes of d4 Through d7 Metal Ions

Hemoglobin and the Octahedral Complex in Heme

Application of crystal field theory
Crystal field theory can offer a good explanation of many properties of transition metal, such as stability、 spectrum、 and magnetism.

Properties of Transition Metal Complexes
usually have color dependent upon ligand(s) and metal ion many are paramagnetic due to unpaired d electrons degree of paramagnetism dependent on ligand(s) [Fe(CN)6]3- has 1 unpaired d electron [FeF6]3- has 5 unpaired d electrons

 Crystal Field Theory Can be used to account for
Colors of transition metal complexes A complex must have partially filled d subshell on metal to exhibit color A complex with 0 or 10 d e-s is colorless Magnetic properties of transition metal complexes Many are paramagnetic # of unpaired electrons depends on the ligand

1. Account for the color of complexs
 Compounds/complexes that have color: absorb specific wavelengths of visible light (400 –700 nm) wavelengths not absorbed are transmitted color observed = complementary color of color absorbed

Visible Spectrum wavelength, nm 400 nm 700 nm higher energy
(Each wavelength corresponds to a different color) 400 nm 700 nm higher energy lower energy White = all the colors (wavelengths)

absorbed color observed color

Relation Between Absorbed and Observed Colors
Absorbed Observed Color  (nm) Color  (nm) Violet Green-yellow Blue Yellow Blue-green Red Yellow-green Violet Yellow Dark blue Orange Blue Red Green

Color & Visible Spectra

Color& Spectra  The plot of absorbance versus wavelength is the absorption spectrum. —For example, the absorption spectrum for [Ti(H2O)6]3+ has a maximum absorption occurs at 510 nm (green and yellow). —So, the complex transmits all light except green and yellow. —Therefore, the complex appears to be purple.

 Absorption of UV-visible radiation by atom, ion, or molecule:
Occurs only if radiation has the energy needed to raise an e- from its ground state to an excited state i.e., from lower to higher energy orbital light energy absorbed = energy difference between the ground state and excited state “electron jumping”

green light observed red light absorbed white light Absorption raises an electron from the lower d subshell to the higher d subshell. For transition metal complexes,  corresponds to energies of visible light.

 Different complexes exhibit different colors because:
color of light absorbed depends on  larger  = higher energy light (Shorter wavelengths) absorbed smaller  = lower energy light(Longer wavelengths) absorbed magnitude of  depends on: ligand(s) metal

red light absorbed (lower energy light)
green light observed red light absorbed (lower energy light) white light [M(H2O)6]3+

Colors of Transition Metal Complexes
white light orange light observed blue light absorbed (higher energy light) [M(en)3]3+

octahedral metal centre
Colour of transition metal complexes Ruby Corundum Al2O3 with Cr3+ impurities Sapphire Corundum Al2O3 with Fe2+ and Ti4+ impurities octahedral metal centre coordination number 6 Emerald Beryl AlSiO3 containing Be with Cr3+ impurities

∆4 ∆3 ∆2 ∆1 green violet yellow yellow red yellow violet violet
The wavelength of light absorbed depends on the size of the energy gap, ∆ , between lower- and higher-energy d orbitals. The size of ∆ depends in part on the identity of the ligands coordinating to the metal ion. Figure below shows a series of four chromium(III) complexes, each with a different ∆ value. ∆4 ∆3 ∆2 ∆1 24Cr3+ 3d3 green violet yellow yellow red yellow violet violet

The energy gap between the three lower-energy d orbitals and the two higher-energy d orbitals is such that visible light causes this electron promotion. Thus, a particular wavelength of visible light is absorbed, giving a coordination complex its characteristic color. The gap between lower- and upper-level d orbitals in the [Ti(H2O)6]3+ ion is of the magnitude that light of 510 nm wavelength promotes an electron. Absorption of this wavelength of light (510 nm is a green wavelength) gives the ion its characteristic red color.

22Ti [Ar]3d24s2 Ti: titanium 22Ti3+ [Ar]3d14s0 [Ti(H2O)6]3

Color of a complex depends on: (i) the metal and (ii) its oxidation state.
Pale blue [Cu(H2O)6]2+ can be converted into dark blue [Cu(NH3)6]2+ by adding NH3(aq). A partially filled d orbital is usually required for a complex to be colored. So, d 0 metal ions are usually colorless. Exceptions: MnO4- and CrO42-. Colored compounds absorb visible light. The color our eye perceives is the sum of the light not absorbed by the complex.

Transition Metales and Coordination Compound

Similar presentations

Presentation on theme: "Transition Metales and Coordination Compound"— Presentation transcript:

Similar presentations

About project

Feedback

Log in

Auth with social network:

Transition Metales and Coordination Compound

Similar presentations

Presentation on theme: "Transition Metales and Coordination Compound"— Presentation transcript:

Similar presentations

About project

Feedback