Chemistry, Canadian EditionChapter 18: Student Study Guide

Chapter 18: The Transition Metals

Learning Objectives

Upon completion of this chapter you should be able to

• predict periodic properties of transition metals

• recognize and name transition metal coordination complexes

• use crystal field theory to explain the colour and magnetic properties of complexes

• explain the importance of transition metal complexes in biological processes

• explain the chemistry of essential steps in the production of pure metals from ores

• recognize the importance of transition metals in everyday life

Practical Aspects

This chapter uses concepts from Chapter 7 (relation of light energy to wavelength and color), Chapter 8, (electron configurations and energy level diagrams), and Chapter 9 (shapes of molecules). In this chapter, you will learn why transition metal compounds are usually vividly coloured. You will also learn about the many important societal uses of transition metals as well as the essential roles they play within biological systems.

18.1overview of the transition metals

Skills to Master:

  • Naming groups of transition metals.
  • Knowing trends in melting point, density, and oxidation states.
  • Determining oxidation states in compounds of transition metals.

Key Concepts:

  • Transition metals are good conductors of heat and electricity.
  • Most transition metals can exist in several different oxidation states.
  • Most transition metals are found in nature as oxides or sulphides, rather than in their elemental form.
  • In transition metal solids, electrons fill in bonding orbital bands up through Group 6. After Group 6, electrons are added to delocalized antibonding orbital bands (a destabilizing factor).
  • The physical properties of a transition element depend upon the element’s electron configuration. For example, some periodic trends are:
  • Melting point – increases across a row as electrons are added to bonding orbital bands, but decreases sharply as soon as electrons are added to antibonding orbital bands.
  • Density – increases within a given column (as elements become more massive), and increases along a given row, while electrons are added to bonding orbital bands. When electrons start adding to antibonding orbital bands, density decreases.
  • Ionization energy – increases across a given row because Zeff increases. (Recall that the definition of ionization energy is the energy required to remove a valence electron from an atom in its gaseous state, so delocalized orbitals are not an option here.) Ionization energy decreases down a column as electrons are held less tightly.

* Remember that there are exceptions to every trend.

Exercise 1: Predict and explain. Which will have the greater: a) melting point, Mo or Tc? b) density, Zr or Hf? c) ionization energy, Sc or V?

Strategy: Use periodic trends to compare the elements.

Solution:

a) Tc is to the immediate right of Mo on the periodic table, so it contains one more proton and electron than Mo. However, that extra electron is the first added to a destabilizing antibonding orbital band. The melting point of Mo is therefore greater than Tc, because it will require more thermal energy to disrupt Mo’s structure than Tc’s.

b) Hf is below Zr on the periodic table, making Hf more massive. Hf will therefore be more dense than Zr.

c) Sc and V are in the same row, but V has two more protons than Sc. The ionization energy depends upon the attraction of the electrons to the nucleus within a gaseous atom. V will therefore have a higher ionization energy than Sc.

Exercise 2: Ag is an excellent electrical conductor. Use the electron configuration of Ag to explain why it is such a good conductor.

Strategy: Electrical conductivity in a metal depends upon the ability of its electrons to delocalize.

Solution: Ag is a d9 transition element. A transition element with over 6 d electrons has electrons in delocalized antibonding orbitals. The electrons in Ag are highly delocalized, making Ag an excellent electrical conductor.

Try It #1: Use periodic trends to explain why Hg is a liquid at room temperature.

18.2coordination complexes

Skills to Master:

  • Recognizing and naming some common ligands.
  • Drawing cis and trans as well as fac and mer isomers.
  • Naming coordination compounds.
Key Terms:
  • Geometric isomers – two complexes that have the same formula, but different 3-D orientations of the ligands. Two main categories:
  • Cis vs. trans – applies to complexes that contain two ligands of one type, ML4X2:

“cis” indicates “same side” / “trans” indicates “across”
  • Facial vs. meridianal – applies to complexes that contain three ligands of one type, ML3X3:

“fac” indicates facial and means that the three ligands lie on one face of the octahedron. / “mer” indicates meridianal and means that the three ligands lie in one plane that goes through the center of the complex.
  • Linkage isomers – two complexes that have the same formula, but show different connectivities of the ligands to the metal center. This occurs only if a given ligand can attach to the metal using either of two donor atoms.

Key Concepts:

  • Ligands are classified by the number of coordination sites they can bind to: monodentate (1), bidentate (2), tridentate (3), etc.
  • In aqueous solution, water molecules will bind to any open coordination sites on a transition metal. For example, Co2+ (aq) is really [Co(H2O)6]2+ (aq).
  • The geometry of a coordination complex depends upon its coordination number.
  • The most common coordination number is 6, which typically corresponds to an octahedral geometry.
  • A coordination number of 4 will result in either a tetrahedral or square planar geometry.

Rules for Naming Transition Metal Complexes

Detailed rules for naming coordination compounds are provided in Section 18.2 of the text. Here is a summary:

All coordination compounds follow this pattern for the root of the name:

  1. Name the cation first and the anion second.
  2. Write the complex as one long word: name the ligands in alphabetical order, then the metal, then indicate the metal’s charge with Roman numerals. Note that these neutral ligands have special names: H2O = aqua, NH3 = ammine (two “m”s), CO = carbonyl.

These suffixes and prefixes must be added accordingly:

  • If the ligand is an anion, add the suffix “–o” to it.
  • Use Greek prefixes to indicate how many ligands of a given type are present. If the ligand name contains a Greek prefix, then use these alternate prefixes to describe the number of ligands: “bis-“ for two, “tris-”for three, “tetrakis-” for four. When using this alternate format, enclose the ligand name in parentheses.
  • If the complex is an anion, add the suffix “–ate” to the metal name.
  • Some complex anions take on the metal’s Greek name. For example, Ag  argentate.

Exercise 3: Determine the oxidation state, coordination number, and number of d valence electrons for the metal in each: a) [Co(NH3)3(NO2)3]; b) Na2[MnCl4]; c) [Ni(en)(H2O)4]SO4.

Strategy: Use guidelines for assigning oxidation states (numbers). The coordination number equals the steric number. The number of d valence electrons can be determined from the electron configuration.

Solution:

a) [Co(NH3)3(NO2)3]: The overall charge is zero, each ammonia is neutral, each of the three nitrites is –1, so the oxidation state of cobalt is +3. There are six monodentate ligands present, so the coordination number is six. The Co3+ ion has the valence electron configuration d6.

b) Na2[MnCl4]: The overall charge is zero, each sodium is +1, each chloride is –1, so the oxidation state of manganese is +2. There are four monodentate ligands present, so the coordination number is four. The Mn2+ ion has the valence electron configuration d5.

c) [Ni(en)(H2O)4]SO4: The overall charge is zero, each ligand is neutral, and sulphate is –2, so the oxidation state of nickel is +2. There is one bidentate ligand (en) and there are four monodentate ligands, so the coordination number is six. The Ni2+ ion has an electron configuration of d8.

Exercise 4: Name each: a) K4[Fe(CN)6]; b) [Ni(en)(H2O)4]SO4; c) [Ni(en)3]Cl2

Strategy: Apply the rules for naming coordination complexes.

Solution:

a) K4[Fe(CN)6]

Cation and anion: The cation is potassium and the anion is the complex ion. The suffix “-ate” will need to be added.

Ligands: Six CN- ligands = “hexacyano.”

Metal charge/name: 4 K+ and 6 CN- = -2, so the charge on iron must be +2. (Fe2+) = ferrate(II).

Name: Potassium hexacyanoferrate(II)

b) [Ni(en)(H2O)4]SO4

Cation and anion: The cation is the complex ion and the anion is sulphate.

Ligands: One en and four waters. “aqua” is before “ethylenediamine” in the alphabet, so “tetraaquaethylenediamine” is the collective name for the ligands.

Metal charge/name: All of the ligands are neutral, so SO42- is the only counter ion: nickel(II).

Name: Tetraaquaethylenediaminenickel(II) sulphate.

c) [Ni(en)3]Cl2

Cation and anion: The cation is the complex ion and the anion is chloride.

Ligands: Three ethylenediamine ligands. The Greek prefix “di-” is in the ligand name, so we must denote “three” with the alternate prefix, tris. The ligand name is tris(ethylenediamine).

Metal charge/name: en is neutral, and there are 2 Cl- ions, so Ni2+ = nickel(II)

Name: Tris(ethylenediamine)nickel(II) chloride.

Try It #2: Name each compound: a) [Co(NH3)3(NO2)3]; b) Na2[MnCl4]

Exercise 5: Write the formula for each: a) Hexamminemolybdenum(III) chloride;

b) Triaquatrifluorocobalt(III); c) potassium dicarbonyltetracyanochromate(II).

Strategy: Follow the rules for naming complex ions.

Solution:

a) Hexamminemolybdenum(III) chloride: This contains 6 ammonia ligands attached to Mo3+, and the counter ion is Cl-. Ammonia is neutral, so 3 Cl- ions are needed for the compound to be electrically neutral: [Mo(NH3)6]Cl3.

b) Triaquatrifluorocobalt (III): Three waters and three fluorides are attached to a Co3+ ion. The three F- ions counterbalance the Co3+, making the complex neutral overall: [Co(H2O)3F3].

c) Potassium dicarbonyltetracyanochromate(II): two CO ligands and four CN- ligands are attached to a Cr2+ metal ion. The charge on the complex is –2, so two K+ ions are needed to counterbalance the charge: K2[Cr(CO)2(CN)4].

Try It#3: Write the formulas for: a) Triaquatrichlororuthenium(III); b) Tetraamminetitanium(II) chloride

Exercise 6: Name the two isomers of [Pt(NH3)4(Cl)2]Cl2, then draw structures for their complex ions.

Strategy: This complex contains two chloride ligands and four ammonia ligands (ML4X2 pattern), so the overall shape is octahedral. The two chloride ligands can be arranged next to each other (cis) or across from each other (trans). The complex has an overall charge of +2.

Solution:

cis-tetraamminedichloroplatinum (IV) chloride / trans-tetraamminedichloroplatinum (IV) chloride

Exercise 7: Name and draw structures for the two isomers that fit this formula: [Co(NH3)3(NO2)3].

Strategy: There are six ligands, so the geometry will be octahedral. When two sets of three ligands are attached to a metal center (ML3X3 pattern), all ligands of the same type can be clustered together on one face of the octahedron, or they can be arranged so that two of the same type are opposite each other. These are termed “facial” and “meridianal,” respectively. The cobalt must have a +3 charge to counter the three –1 charges of the nitrites. Notice that within the complex “NO2-” is named “nitro.”

Solution:

fac-triamminetrinitrocobalt(III) / mer-triamminetrinitrocobalt(III)

Try It #4: Draw the structure for the dicarbonyltetracyanochromate(II) ion. If more than one isomer exists, draw both isomers.

18.3bonding in coordination complexes

Skills to Master:

  • Constructing a crystal field energy-level diagram for an octahedral complex.
  • Determining the electron configuration of a complex from magnetic properties.
  • Determining the value of ?from the absorption spectrum of a complex.
Key Terms:
  • Crystal field splitting energy () – difference in energy between non-degenerate d orbitals.
  • Pairing energy (P) – destabilizing energy resulting from repulsive forces between two electrons that occupy the same orbital.
  • High-spin – term used to describe electron arrangement in which the maximum unpaired spins are present. This occurs when P > , and it is easier for the electrons to occupy the higher energy d orbitals than to pair up in the lower energy orbitals.
  • Low-spin – term used to describe electron arrangement in which the maximum paired spins are present. This occurs when P < .
  • Spectrochemical series – list of ligands in order of increasing energy level splitting ability:

I- < Br- < Cl- < F- < OH- < H2O < NH3 < en < NO2- < CN- < CO

Key Concepts:

  • All d orbitals are not degenerate. The set of d orbitals is split into two or more degenerate groups, based on slight differences in energy arising from electron-electron repulsions.
  • The magnitude of the crystal field splitting energy depends upon these factors:
  • the charge on the metal ion – the greater the charge, the greater the splitting.
  • the size of the metal ion – the greater the n, the greater the splitting.
  • the coordination number – the higher the coordination number, the greater the splitting.
  • the identity of the ligand – the more tightly the ligand binds, the greater the splitting.
  • The pattern of orbital splitting depends upon the geometry of the metal cation:

Geometry:

/ Octahedral / Tetrahedral / Square Planar
Crystal Field Energy Level Diagram: / / /
Notes: / The t2g set includes the dxy, dxz, and dyz orbitals. The eg set includes the orbitals. / The two sets are the reverse of octahedral. Splitting is smaller than in octahedral, so tetrahedral complexes are almost always
high-spin. / The orbital is greatly destabilized relative to the other orbitals because of its orientation relative to the ligands in this geometry.
  • The same rules apply for drawing energy level diagrams for transition metal cations as did for atoms - Hund’s Rule, Aufbau Principle, Pauli Exclusion Principle.
Helpful Hints
  • One way to remember the spectrochemical series is to notice that it is broken down into groups of what’s attaching to the metal: “halogens < oxygen < nitrogen < carbon.”
  • The names “t2g” and “eg” are used to describe the two categories of d orbitals in an octahedral geometry. The names “e” and “t” are used to describe d orbitals in a tetrahedral geometry.
  • An energy-level diagram for a transition metal ion is sometimes called a crystal field energy diagram.

Exercise 8: Draw a crystal field energy diagram for each complex ion. Then, write the electron configuration for each. a) [Co(NH3)3(NO2)3]; b) Na2[MnCl4], tetrahedral; c) K4[CrF6].

Strategy: Determine the splitting pattern from the orbital geometry. Then determine the relative magnitude of the  by assessing the factors that affect it: the size and charge of the metal ion, the coordination number, and the identity of the ligand.

Solution:

a) [Co(NH3)3(NO2)3]: Co3+, 3d6 arrangement, octahedral geometry. Relatively high oxidation state, fairly strong ligands. These factors indicate a high splitting energy. The complex will be low-spin.

b) Na2[MnCl4], tetrahedral: Mn2+, 3d5 arrangement. Tetrahedral complexes are almost always high spin, because their splitting energies are small. The weak Cl- ligands shouldn’t increase the splitting energy.

c) K4[CrF6]: Cr2+, 3d4 arrangement, octahedral geometry. Lower oxidation state, weak ligands. These factors favour a low splitting energy. This complex will be high-spin.

Complex: / [Co(NH3)3(NO2)3] / Na2[MnCl4] / c) K4[CrF6]
Diagram: / / /
Electron configuration: / (t2g)6(eg)0 / (e)2(t)3 / (t2g)3(eg)1

Try It #5: Draw the crystal field splitting diagram and electron configuration for [Pt(NH3)6]Cl2.

Magnetism and Color

Key Term:

  • Complementary colors – colors related to each other by their absorption and transmission characteristics. Table 18-5 in the text lists complementary colors.

Key Concepts:

  • Recall that paramagnetic substances can behave like magnets because they contain unpaired electron spins, while diamagnetic substances don’t behave like magnets because all of their spins are paired.
  • The crystal field splitting energy of a metal complex is related to its color.
  • We can see reflected light (light that bounces off a substance) and transmitted light (light that passes through a substance). We see the complement of the color that the substance absorbs.

Useful Relationship:

  • Emolecule =  = h = hc/
/ This relationship can be used to relate  to the wavelength of light that the complex absorbs.

Exercise 9: Which of the complexes listed in Exercise 8 will respond most strongly to a magnetic field? Which will not respond at all?

Strategy: A paramagnetic substance will respond to a magnetic field, and a diamagnetic substance will not. The substance with the highest overall spin will respond the most intensely.

Solution: [Co(NH3)3(NO2)3] has all spins paired, so it is diamagnetic. It will not respond to a magnetic field. Both Na2[MnCl4] and K4[CrF6] have unpaired spins, so they are both paramagnetic. Na2[MnCl4], with five unpaired spins, will respond slightly more to a magnetic field than K4[CrF6], with four unpaired spins.

Exercise 10: [Ni(en)3]SO4 is lavender, [Ni(en)(H2O)4]SO4 is blue, and [Ni(H2O)6]SO4 is teal. Explain the color differences.

Strategy: The color we can see is the complement of the color that the substance absorbs. Use Table 18-5 in the text to determine the color that each complex absorbed, then compare the relative energies of the absorbed colors.

Solution: Table 18-5 shows that:

Complex / Color observed / Color absorbed / Wavelength absorbed
[Ni(en)3]SO4 / Lavender (light purple) / Yellow-green / 560 nm
[Ni(en)(H2O)4]SO4 / blue / Orange / 610 nm
[Ni(H2O)6]SO4 / Teal (blue-green) / Red-orange / 680 nm

The only difference between each complex is the ligands. H2O is lower in the spectrochemical series than en, so the splitting energy for the all-water complex should be the lowest. The all-en complex should have the highest splitting energy. This is the trend observed. The complex with the shortest wavelength of light absorbed corresponds to the complex with the greatest splitting energy.

Exercise 11: The splitting energy of [Rh(Cl)6]3- is 243 kJ/mol. a) Determine the wavelength of maximum absorption for this complex. b) What color is it?

Strategy: Use the relationship E = hc/ to determine the wavelength. Use Table 18-5 in the text to determine the color.