Transition Metals and Coordination Compounds Flashcards

The Chemistry of Transition Metals and Coordination Compounds

• Transition metals are situated in the d block of the periodic table and share several physical properties.

• Most possess high densities, high melting points, extreme hardness, and excellent electrical conductivity due to electrons occupying d orbitals participating in metallic bonding.

• Coordination compounds contain complex ions consisting of a central metal ion bonded to one or more ligands.

• These compounds serve critical roles as therapeutic drugs, sensors, coloring agents, and components in biological systems like hemoglobin.

Case Study: Gemstone Colorations and Crystal Field Splitting

• The color variations in gemstones like rubies and emeralds are due to the presence of the $Cr^{3+}$ ion in different host crystals.

• Ruby: Composed of aluminum oxide ($Al_2O_3$) where approximately $1\%$ of $Al^{3+}$ ions are replaced by $Cr^{3+}$. The crystal field is stronger, leading to larger d-orbital splitting and a red appearance.

• Emerald: Composed of beryllium aluminum silicate ($Be_3Al_2(SiO_3)_6$) with $Cr^{3+}$ replacement. The crystal field is weaker, resulting in green light reflection.

• Garnet and Peridot: Both contain $Fe^{2+}$ ions. In garnet ($Mg_3Al_2(SiO_4)_3$) it appears red; in peridot ($Mg_2SiO_4$) it appears yellow-green.

• Turquoise: The blue color in $[Al_6(PO_4)_4(OH)_8 \cdot 4H_2O]^{2-}$ is caused by transitions in $Cu^{2+}$ d orbitals.

Electron Configurations and Ion Formation

• Ground state configurations follow the general pattern: $[\text{Noble Gas}]\,ns^2(n-1)d^x$

• Notable Exceptions:

  • Chromium (Cr): $[Ar]\,4s^1 3d^5$ (rather than $4s^2 3d^4$).

  • Copper (Cu): $[Ar]\,4s^1 3d^{10}$ (rather than $4s^2 3d^9$).

  • These occur due to the stability of half-filled or fully filled d sublevels.

• Transition Metal Ionization: Electrons are always removed from the outermost $ns$ orbital before the $(n-1)d$ orbitals.

  • Example: $Fe$ is $[Ar]\,4s^2 3d^6$; $Fe^{2+}$ is $[Ar]\,3d^6$.

  • Example: $Co$ is $[Ar]\,4s^2 3d^7$; $Co^{3+}$ is $[Ar]\,4s^0 3d^6$ or $[Ar]\,3d^6$.

Periodic Trends in the d Block

• Atomic Radius:

  • Horizontal variation is minimal because electrons are added to $(n-1)d$ orbitals, providing consistent shielding for the outermost electrons.

  • Vertical variation: Sizes increase from the first to second row, but the third row remains similar to the second due to the Lanthanide Contraction ($14$ electrons in $(n-2)f$ orbitals shield poorly, allowing the nucleus to pull outer electrons tighter).

• Ionization Energy:

  • Slowly increases across a row.

  • Vertically, the third row has higher ionization energies than the first two rows because outer electrons are held more tightly due to the lack of size increase between the second and third periods.

• Electronegativity:

  • Generally increases across a row.

  • Increases from the first to second row, with little change between the second and third rows.

  • Gold ($Au$) is the most electronegative metal ($EN = 2.4$).

• Oxidation States:

  • Transition metals exhibit multiple oxidation states, unlike main-group metals.

  • Highest state is $+7$ (Manganese), corresponding to the loss of all $4s$ and $3d$ electrons.

  • Metals in high oxidation states typically exist bound to electronegative elements like oxygen rather than as bare ions.

Components of Coordination Compounds

• Complex Ion: A central metal atom or ion bonded to ligands.

• Ligand: A Lewis base (electron pair donor) that forms a coordinate covalent bond with the metal.

• Counterions: Ions of opposite charge that maintain electrical neutrality in the coordination compound (e.g., $Cl^-$ in $[Co(NH_3)_6]Cl_3$).

• Primary Valence: The oxidation state of the metal atom.

• Secondary Valence: The coordination number, or the number of sites bonded directly to the metal.

Classification of Ligands

• Monodentate: Donate one electron pair (e.g., $H_2O$, $NH_3$, $Cl^-$, $CN^-$).

• Bidentate: Donate two electron pairs from different atoms (e.g., ethylenediamine ($en$), oxalate ($ox$)).

• Polydentate: Donate multiple electron pairs. Ethylenediaminetetraacetate ($EDTA^{4-}$) is hexadentate, wrapping around the metal and donating six pairs.

• Chelates: Complex ions containing bidentate or polydentate ligands.

Geometries of Complex Ions

• Coordination Number 2: Linear geometry (e.g., $[Ag(NH_3)_2]^+$).

• Coordination Number 4:

  • Square Planar: Common for $d^8$ configurations (e.g., $[PdCl_4]^{2-}$).

  • Tetrahedral: Common for $d^{10}$ configurations (e.g., $[Zn(NH_3)_4]^{2+}$).

• Coordination Number 6: Octahedral geometry (e.g., $[Co(NH_3)_6]^{3+}$).

Nomenclature Rules for Coordination Compounds

1. Name the cation before the anion.

2. In the complex ion, name ligands in alphabetical order before the metal.

3. Ligand Names:

   • Neutral: $H_2O$ (aqua), $NH_3$ (ammine), $CO$ (carbonyl).

   • Anionic: End in -o (e.g., chloride becomes chloro, cyanide becomes cyano).

4. Prefixes: Use di-, tri-, tetra-, penta-, hexa- for count. Use bis-, tris-, tetrakis- if the ligand name already has a prefix.

5. Metal Name:

   • If the complex is a cation: Use the metal name and Roman numeral (e.g., platinum(II)).

   • If the complex is an anion: Add -ate to the metal name (e.g., platinate(II)). Use Latin roots for specific metals: Iron (ferrate), Copper (cuprate), Gold (aurate), Silver (argentate), Lead (plumbate), Tin (stannate).

Isomerism in Coordination Compounds

Structural Isomers

• Coordination Isomers: Ligands and counterions swap positions (e.g., $[Co(NH_3)_5Br]Cl$ vs. $[Co(NH_3)_5Cl]Br$).

• Linkage Isomers: A ligand bonds to the metal through different atoms.

  • Nitrite can bond via $N$ (nitro) or $O$ (nitrito).

  • Thiocyanate ($SCN^-$) can bond via $S$ (thiocyanato) or $N$ (isothiocyanato).

Stereoisomers

• Geometric Isomers:

  • Cis-trans: Occurs in square planar ($MA_2B_2$) and octahedral ($MA_4B_2$) complexes. "Cis" ligands are adjacent; "trans" ligands are opposite. Note: Tetrahedral complexes do not show cis-trans isomerism.

  • Fac-mer: Occurs in octahedral $MA_3B_3$ complexes. "Facial" (fac) isomers have three identical ligands on one face; "meridional" (mer) isomers have them in an arc/meridian.

• Optical Isomers: Nonsuperimposable mirror images (enantiomers) that exhibit optical activity. Often found in octahedral complexes with chelating ligands (e.g., $[Co(en)_3]^{3+}$).

Bonding Models: Valence Bond and Crystal Field Theory

• Valence Bond Theory: Coordinate covalent bonds result from ligands donating electron pairs into hybridized metal orbitals. Hybridization sets include:

  • $sp$ (Linear)

  • $sp^3$ (Tetrahedral)

  • $dsp^2$ (Square Planar)

  • $d^2sp^3$ (Octahedral)

• Crystal Field Theory (CFT): Explains color and magnetism via the repulsion between ligand electrons and metal d-orbital electrons.

  • Octahedral Splitting: Ligands on the axes repel $d_{x^2-y^2}$ and $d_{z^2}$ orbitals more, raising their energy. The $d_{xy}$, $d_{xz}$, and $d_{yz}$ orbitals remain at lower energy.

  • Crystal Field Splitting Energy ($\Delta$): The energy gap between split orbital levels.

Magnetism and Spin Configurations

• Strong-Field Complexes: $\Delta$ is large (\Delta > \text{Pairing Energy}). Electrons fill the lower energy levels completely before occupying higher levels, leading to low-spin (fewer unpaired electrons).

• Weak-Field Complexes: $\Delta$ is small (\Delta < \text{Pairing Energy}). Electrons follow Hund's rule and occupy all d orbitals singly before pairing, leading to high-spin (more unpaired electrons).

• Spectrochemical Series: Arranges ligands by field strength:

  • CN^- > NO_2^- > en > NH_3 > H_2O > OH^- > F^- > Cl^- > Br^- > I^-

• Metals with higher charges ($+3$ vs $+2$) generally have larger $\Delta$, as they draw ligands closer.

Calculating Splitting Energy from Spectra

• The color observed is the complement of the color absorbed.

• $\Delta = E_{\text{photon}} = \frac{hc}{\lambda}$

• Example calculation for $[Ti(H_2O)_6]^{3+}$ with maximum absorbance at $498\,nm$:

  • $\Delta = \frac{(6.626 \times 10^{-34}\,J \cdot s) \times (3.00 \times 10^8\,m/s)}{(498\,nm) \times (1 \times 10^{-9}\,m/nm)} = 3.99 \times 10^{-19}\,J$

  • In molar terms: $\Delta = (3.99 \times 10^{-19}\,J/ion) \times (6.02 \times 10^{23}\,ions/mol) = 240\,kJ/mol$

Splitting in Alternative Geometries

• Tetrahedral: Pattern is the reverse of octahedral ($d_{xy}, d_{xz}, d_{yz}$ are higher). Most are high-spin because $\Delta$ is smaller (only 4 ligands).

• Square Planar: Most complex splitting pattern; primarily observed in low-spin $d^8$ metal ions.

Applications of Coordination Compounds

• Medical Treatments: $[Ca(EDTA)]^{2-}$ is used for lead poisoning treatment; lead displaces calcium because its complex is more stable ($K_f = 2 \times 10^{18}$ vs $4 \times 10^{10}$).

• Cisplatin: The cis isomer of $[Pt(NH_3)_2Cl_2]$ is an effective anticancer drug, whereas the trans isomer is ineffective due to its inability to bind correctly to DNA.

• Biomolecules:

  • Hemoglobin: Contains an iron heme complex in an octahedral arrangement for $O_2$ transport.

  • Chlorophyll: Magnesium coordinated to a porphyrin ring for photosynthesis.

  • Carbonic Anhydrase: Zinc in a tetrahedral complex that catalyzes $H_2O + CO_2 \rightleftharpoons H^+ + HCO_3^-$.

• Chemical Analysis: $SCN^-$ identifies $Co^{2+}$ (blue) or $Fe^{3+}$ (red); dimethylglyoxime ($dmg$) identifies $Ni^{2+}$ (red) or $Pd^{2+}$ (yellow).