Topic 23: Transition Metals Flashcards
Electron configuration of transition metals in periods n = 4 - 5
[NG] ns2 (n - 1)dx
Electron configuration of transition metals in periods n = 6 - 7
[NG] ns2 (n - 2)f14 (n - 1)dx
Partial electron configuration of transition metals
ns2 (n - 1)dx
Exceptions of e- configuration
Cu - 4s1 3d10
Cr - 4s1 3d5
Ions formation in transition metals
Loss of ns e- before (n-1)d e-
Atomic size of transition metals across a period
Less steady decrease than across the main group
d e- fill inner orbitals, shielding outer e- efficiently
Electronegativity of transition metals across a period
Relatively constant EN
Consistent with their relatively constant size
Ionization energy of transition metals across a period
a) Values increase relatively little
b) Inner 3d e- shield more effectively
Density of transition metals across a period
Increase, level off, and dip a bit
Atomic size of transition metals within a group
a) Increase from Period 4 to 5
b) No size increase from period 5 to 6
Explanation of lanthanide contraction
a) 14 4f e- shield outer e- poorly from the increase in nuclear charge
b) Size decrease = Normal increase | Same atomic size
Electronegativity of transition metals within a group
a) Increase from Period 4 to 5
b) No increase in period 6
Explanation of bonding in transition metals
Heavier transition metals form bonds with more covalent character as they attract e- more strongly than main group metals do
Ionization energy of transition metals within a group
a) Small increase in size with a large increase in nuclear charge
b) IE values generally increase down a transition group
Density of transition metals within a group
a) Increase dramatically as atomic volumes change little, but atomic masses increase
b) Period 6 has the densest elements known: tungsten/rhenium/osmium/iridium/platinum/gold
Patten of multiple oxidation states in transition metals
a) The highest oxidation state of elements in Group 3B (3) through 7B (7) equals the group #
b) Elements in Groups 8B (8), 8B (9), and 8B (10) exhibit fewer oxidation states
=> +2 and +3 states are common for Fe / Co
=> +2 state is common for Ni, Cu, and Zn
Explanation of multiple oxidation states in transition metals
As ns and (n-1)d e- are close in energy, transition elements can use them in bonding
Metallic behavior of transition metals based on their oxidation state
For lower oxidation states, metals behave chemically more like metals
For higher oxidation states, covalent bonding is more prevalent
Oxide acidity and oxidation state in transition metals
As the oxidation state increases, the oxide becomes less basic (more acidic)
Definition of valence-state electronegativity
A metal atom with a positive oxidation state has a greater attraction for the bonded e- than it does when it has a zero oxidation state
Reducing strength in transition metals
All Period 4 transition metals (except Cu) are active enough to reduce H+ to form H2(g)
Explanation of color of transition metals
e- in a partially filled d orbitals absorb visible wavelengths and move to slightly higher energy orbitals
Which transition metals might not show color? Why?
Exception: Sc / Ti / Zn
Reason: Empty or fully filled orbitals
Magnetic properties of transition metals
Many transition metal compounds are paramagnetic due their unpaired d e-
Reactivity within a group
IE increases down a group, decreasing reactivity
Why leads to similar chemical behavior of elements of Periods 5/6?
Similar atomic size
Ores often occur together in nature
Which are inner transition elements?
14 lanthanides in Period 6
=> [Ce (58) – Lu (71)] | 7 inner 4f orbitals
14 actinides in Period 7
=> [Th (90) – Lr (103)] | 7 inner 5f orbitals
Abundance of lanthanides
Rare earth elements
Exception: Ce ranks 26 in abundance
e- configuration of lanthanides
Ground-state e- configuration [Xe] 6s2 4fx 5d0
Exception of e- configuration of lanthanides
Ce ([Xe] 6s2 4f1 5d1) forming stable 4+ ion
Gd3+ and Lu3+ (Stable half-filled or filled f sublevel)
Physical properties of lanthanides
Silvery, high-melting metals
Chemical properties of lanthanides
a) Elements exist as M3+ ions of very similar radii
b) Ores of lanthanides (Ce / Y) are mixtures of compounds of all 14
Applications of lanthanides
a) Tinted glass / Electronic devices
b) Gasoline refining
c) Steelmaking
Properties of actinides
All actinides are radioactive
Uranium has +6 state as the most prevalent
Components of coordination compounds
Complex ion - Central metal ion bonded to molecules/anions
Counterion - Cation/Anion/Complex ion
Factors that affect geometry of coordination compounds
a) Coordination number
b) Ligands
Coordination number definition
Number of ligands atoms bonded directly to the central metal ion in a complex ion
Geometry based on coordination number
2 - linear
4 - Square planar (d8)
4 - Tetrahedral (d10)
6 - Octahedral
Definition of ligands
Molecules or anions with one or more donor atoms
Monodentate ligands
H2O
NH3
Halides
Cyanide ion
Thiocyanate ion -
Hydroxide ion -
Nitrite ion -
Chelate definition
A compound containing a ligand bonded to a central metal atom at two or more points.
Bidentate ligands
ethylenediamine (en)
oxalate ion 2-
Polydentate ligands
diethylenetriamine (3)
Triphosphae ion 5- (3)
EDTA 4- (6)
Rules for writing formulas in coordination compounds
a) Cation is written before the anion
b) Charge of cation is balanced by the charge of anion
c) Neutral ligands (in alphabetical order) are written before anionic ligands (in alphabetical order)
d) Formula of the complex ion is placed in brackets
e) Counter ions are written outside the brackets
Rules for naming coordination compounds
a) Cation is named before the anion
b) Within the complex ion, ligands are named in alphabetical order before the metal ion (Numerical prefixes)
c) Oxidation state of the metal ion has a roman numeral (in parantheses)
d) Typographical space comes between cation and anion
e) If the complex ion is an anion, drop the ending of the metal name and add -ate
Prefix for ligands
a) H2O
b) NH3
c) CO
d) NO
e) OH-
f) CN-
g) Halide
aqua
ammine
carbonyl
nitrosyl
hydroxo
cyano
fluoro/chloro/bromo
Latin root used for metal ions
Iron (Fe) - Ferrate
Copper (Cu) - Cuprate
Lead (Pb) - Plumbate
Silver (Ag) - Argentate
Gold (Au) - Aurate
Tin (Sn) - Stannate
Constitutional isomers
Compounds with the same formula but atoms connected differently
Types of constitutional isomers
a) Coordination isomers
b) Linkage isomers
Coordination isomers definition
Composition of complex ion is different but not of the compound
a) Exchange of ligand and counter ion
b) Exchange of ligand between two complex ions
Linkage isomer definition
Composition of complex ion is the same, but ligand donor atom is different
Naming based on linkage isomer
a) Nitrite ion
b) Cyanate ion
c) Thiocyanate ion
a) Nitrito / Nitro
b) Cyanato / Isocyanato
c) Thiocyanato / Isothiocyanato
Stereoisomers definition
Compounds with the same atomic connections but different atom spatial arrangement
Types of stereoisomers
a) Geometric
b) Optical
Geometric isomers
Atoms arranged differently in space relative to the central metal ion
a) Cis (Next to each other)
b) Trans (Across)
Optical isomers (Enantiomers)
Molecule with its mirror image that cannot be superimposed
a) Identical except for the direction in which they rotate the plane of polarized light
Valence bond theory principle
a) Filled ligand orbital overlaps an empty metal ion orbital, forming a coordinate covalent bond (LA / LB)
b) Mixing particular combinations of s, p, and d orbitals to obtain sets of hybrid orbitals
Orbital hybridization determines the geometry of the complex ion
a) Octahedral
b) Square planar (d8)
c) Tetrahedral (d10)
a) 3 d orbitals + 6 d2sp3 orbitals
b) 4 d orbitals + 4 dsp2 orbitals
c) 5 d orbitals + 4 sp3 orbitals
Strengths / Limitations of valence bond theory
Strength
a) Explanation for bonding and shape
Limitation
a) No insight into the colors of complex ions
b) No explanation for magnetic properties
Crystal field theory in octahedral complex
a) Electrostatic attraction between metal cation and negative charge of the ligands
b) e- pairs from ligands repel e- in the 5 d orbitals of the metal ion
c) d e- are repelled unequally due to their different orbital orientation
Repulsion of e- in octahedral complex
a) Stronger repulsion in d(x2 - y2) and d(z2) than the other 3 orbitals
b) Ligands moving along x, y, and z approach…
=> Directly towards lobes of d(x2 - y2) and d(z2)
=> Between lobes of dxy, dxz, and dyz
Name of groups of d-orbitals in octahedral complex
eg orbitals: d(x2 - y2) / d(z2)
t2g orbitals: dxy / dxz / dyz
Comparison of field split in tetrahedral and octahedral
Δtetrahedral < Δoctahedral
Only high spin tetrahedral complexes are known
Repulsion of e- in tetrahedral complex
d(x2 - y2) + d(z2) low energy
dxy + dxz + dyz high energy
Spinning in tetrahedral complexes
Only high spin tetrahedral complexes are known as Δ is always smaller than Δ_pairing
Arrangement of orbitals based on energy in square planar complexes
1) d(x2 - y2)
2) d(xy)
3) d(z2)
4) dxz + dyz
Spinning in square planar complexes
Low spin and diamagnetic as lower energy orbitals are filled first
Weak field ligands
a) High spin = Max # unpaired e-
b) Smaller splitting of energy
Strong field ligands (Above / equal to NH3)
Low spin = Max # unpaired e-
Larger splitting of energy
Color of transition metal complexes according to crystal field theory
a) e- uses absorbed radiant energy from photon to move to higher energy level
b) Δ corresponds to wavelength of visible spectrum
c) Complementary color of the wavelength absorbed is reflected
Factors that affect the color of transition metal complexes
a) Oxidation state of metal ion
b) Ligands (Spectrochemical series)
Formula to calculate the wavelength of a photon
∆Ephoton = hc/λ
The relative sizes of Epairing and crystal field determine the occupnacy of d orbitals
a) Weak field ligands: # unpaired e- in the complex ion is the same as in the free ion
b) Strong field ligands: # unpaired e- in the complex ion is less than in the free ion
What defines the magnetic properties of a complex ion?
Number of unpaired e- in d orbitals
