Transcript Chapter 1 Structure and Bonding

Chapter 20 The Transition Metals
I.
Survey of Transition Metals
A.
General Properties
1) Transition metals have similar reactivity across the period
a) Representative Elements: filling valence shells as you go across a period
b) Transition Metals: filling core subshells (d or f) as you go across a period
i. d-block elements (Sc—Zn)
ii. f-block elements (Ce—Lu)
c) Core electrons don’t effect reactivity as much as valence electrons
2)
All Transition elements are metals: luster, conductors, malleable, ductile
3)
Transition metals do vary in their properties
a. Melting points: W = 3400 oC, Hg = 25 oC
b. Hardness: Titanium is hard, Gold is soft
c. Chemical Reactivity: Iron is very reactive (rust), Platinum is very inert
4)
Formation of Ionic compounds with nonmetals
a. More than one oxidation number is possible: FeCl2 vs. FeCl3
b. Different than Main Group metals: NaCl, CaCl2, AlCl3
5)
Formation of Complex Ions
a. Transition metal cation at the center
b. Multiple Lewis Bases (Ligands) donating lone pairs
6)
Compounds are generally colored
a. Absorption of light by d-electrons
b. Most main group compounds are colorless (or black)
7)
Many transition metal complexes are paramagnetic: unpaired d-electrons
B.
Electron Configuration of First Row Transition Metals (Sc—Zn)
1) Expected configuration is usual: Sc = [Ar]4s23d1, Ti = [Ar]4s23d2, V = [Ar]4s23d3
2) Chromium has an unexpected configuration. Why? Cr = [Ar]4s13d5
a) 3d and 4s subshells have almost same energies for first row metals
b) Cr occupies essentially degenerate orbitals equally: 4s__ 3d__ __ __ __ __
3)
4)
Copper is the only other exception having a single 4s electron: Cu = [Ar]4s13d10
For cations, 3d energy < 4s energy
a) 4s electrons are lost first because remaining electrons are more stable
Cr2+ = d4
4s__
3d__ __ __ __ __
b)
c)
C.
Mn2+
Ti3+
Mn0 = [Ar]4s23d5
Ti0 = [Ar]4s23d2
Mn2+ = [Ar]3d5
Ti3+ = [Ar]3d1
Oxidation state and ionization energy (Table 21.2)
1) For the first five metals: maximum oxidation state = loss of all 4s and 3d electrons
Cr0 -----> Cr6+ (loss of 4s13d5)
Ti0 -----> Ti4+ (loss of 4s23d2)
2) Last five metals: usually +2 or +3
a) 3d energy lowers as the nuclear charge increases
b) Harder to remove more electrons
3) Ionization Energy increases across the period: M0 -----> Mn+ + n e-
First row transition metal Ionization Energies
Removal of first d-electron
Removal of first s-electron
D.
1)
Standard Reduction Potentials
Generally listed with metal as
reducing agent (metal is oxidized)
2)
This is the reverse of how reduction
potentials are usually written
2H+ + 2e- -----> H2 E = 0.0 V becomes
H2 -----> 2H+ + 2e- E = 0.0 V
3)
“Oxidation Potentials” as written
would have the most easily oxidized
metal with the most positive potential
4) All but copper can reduce H+ to H2
Zn + 2H+ -----> Zn2+ + H2 E0 = +0.76
5)
Reducing Ability decreases from left
to right (Zn and Cr are exceptions)
Reaction
Potential (V)
Sc -----> Sc3+ + 3e-
2.08
Ti -----> Ti2+ + 2e-
1.63
V -----> V2+ + 2e-
1.20
Mn -----> Mn2+ + 2e-
1.18
Cr -----> Cr2+ + 2e-
0.91
Zn -----> Zn2+ + 2e-
0.76
Fe -----> Fe2+ + 2e-
0.44
Co -----> Co2+ + 2e-
0.28
Ni -----> Ni2+ + 2e-
0.23
Cu -----> Cu2+ + 2e-
-0.34
E.
4d and 5d Transition Metals
1. Atomic Radii
a. Decreases from left to right: more +/- attraction in same shell
b. Increases as you go from 3d to 4d to 5d: filling larger shell
c. 4d and 5d metals are very similar in size = Lanthanide Contraction
i. Fifth period fills up 4f electrons before beginning to fill 5d subshell
ii. Each 4f electron added decreases the radius, so the atoms are smaller
than expected by the time we get to 5d
iii. By the time we get to La, size is back down to near that of Y
2. Other Properties
a. Generally less reactive than first row metals (poor orbital size match)
b. Important for specialized materials
c. Range from rare to very rare in abundance on Earth’s crust
II.
Chemistry of the First Row Transition Metals
A.
Scandium (Sc)
1) Rare element found mostly as Sc3+ in ionic compounds like Sc2O3
2) Chemistry is most similar to Lanthanides = colorless, no d-electrons, diamagnetic
3) Sc0 is used in some high intensity lamps
B.
Titanium (Ti)
1) Common (0.6% of crust), strong, low density metal; structural uses (jet engines)
2) Fairly unreactive, so used for pipes, pumps, reaction vessels
3) Usually found in Ti4+ state, no d-electrons, colorless/white
4) TiO2 (titanium(IV) oxide) used as white pigment in paper, paint, etc…
5) Ti3+ is stable: Ti(H2O)63+ complex ion is purple (1 d-electron)
6) Ti2+ is not very stable, but TiO and TiCl2 do exist
C.
Vanadium (V)
1) Common (0.02%) metal most often used as alloy with Fe or Ti
2) Usually found in V5+ state, colorless, no d-electrons
3) V2O5 is an industrial catalyst used in the manufacture of H2SO4
4) V4+, V3+, V2+ can all exist; usually strong reducing agents (form V5+)
D.
Chromium (Cr)
1) Rare, but important industrially
2) Chromite Ore reduced to Ferrochrome = mixture of Fe and Cr used in steelmaking
FeCr2O4 + 4C -----> Fe + 2Cr + 4CO
3)
4)
E.
Pure Cr metal used to plate steel: shiny surface due to “invisible” Cr oxide
Cr2+, Cr3+, and Cr6+ oxidation states are common (Table 20.5)
a. Cr2+ is a powerful reducing agent (used to remove O2)
b. Cr6+ is a powerful oxidizing agent: Dichromate = Cr2O72- -----> Cr3+
Cr2O72- + 14H+ + 6e- -----> 2Cr3+ + 7H2O Eo = 1.33 V
c. Dichromate equilibrates in basic water to chromate = CrO42d. CrO3 in H2SO4 very strong oxidizer used to clean glassware
Manganese (Mn)
1) Abundant (0.1%) but no sources in the United States
2) With iron, it makes an especially hard steel (crush rock, bank vaults)
3) Nodules of almost pure Mn made on sea floor by marine organisms
4) Mn2+ up to Mn7+ are all possible, Mn2+ and Mn7+ are most common
5) Permanganate = MnO4- deep purple strong oxidizing agent
MnO4- + 8H+ + 5e- -----> Mn2+ + 4H2O Eo = 1.5 V
E.
Iron (Fe)
1) Most abundant transition metal (4.7%)
2) Solid form rapidly reacts with O2 to form rust = Fe2O3
3) Usually found as Fe2+, Fe3+ oxidation states
4) Makes steel in combination with other elements
5) Important biologically (later)
F.
Cobalt (Co)
1) Rare, found in ores = smaltite = CoAs2
2) Hard metal with a blue/white appearance
3) Mostly used in alloys with Fe, Cu, W
4) Co2+ and Co3+ are most common oxidation states
5) Historically important in the development of “Coordination Chemistry”
G.
Nickel (Ni)
1) Abundant in ores with As, Sb, and S
2) Corrosion resistant, white metal used to plate other metals
3) Usually in the Ni2+ oxidation state, such as in the green compound Ni(H2O)62+
4) Well characterized in many coordination compounds
H.
Copper (Cu)
1) Abundant in ores with S, As, Cl, and CO322) High electrical conductance, low corrosion make it suitable for wires and pipes
3) Corrodes in air to green “patina”: Cu + 2H2O + SO2 + O2 -----> Cu3(OH)4SO4
4) Usually found as Cu+ or Cu2+ as in blue Cu(H2O)62+
5) Biologically necessary, but toxic in large amounts
I.
Zinc (Zn)
1) Abundant in ore = sphalerite = ZnS
2) Good reducing agent: Zn -----> Zn2+ + 2e- E = 0.76 V
3) Colorless, full d-orbitals
4) Zn2+ is the only oxidation state commonly encountered
III. Coordination Compounds
A.
Definitions
1) Complex Ion = metal cation with Lewis Bases covalently bonded to it
2) Ligands = Lewis bases covalently bonded to a metal ion
3) Counter Ions = anions present in the formula to balance the charge of complex ion
a) [Co(NH3)5Cl]Cl2 = Co(NH3)5Cl2+ (complex ion) + 2 Cl- (counter ions)
b) Solid = packed cations/anion; Solution = dissociated like NaCl
4)
B.
Alfred Werner developed theory of coordination chemistry in 1890’s
a) Primary Valence = total positive charge = number of anions total
b) Secondary Valence = Lewis Acid ability = how many Lewis Bases bound
c) [Co(NH3)5Cl]Cl2 = CoCl3 • 5 NH3
i. Primary Valence: +3 is satisfied by 3 Cl- total
ii. Secondary Valence: 6 Lewis bases bound to Co (5 NH3, 1 Cl)
iii. Primary Valence = Oxidation State
iv. Secondary Valence = Coordination Number
Coordination Number
1) Metals can bind between 2 and 8 ligands:
Coordination Number = CN = 2-8
2) Size, charge, d-electron count decide
how many ligands bind
3) Most prevalent CN = 6 > 4 > 2
4) A given metal ion can have any CN,
depending on what the ligands are
5) Common coordination numbers: Table 21.12
C.
Ligands
1) Coordinate Covalent Bond = Lewis base donates both electrons (lone pair)
2) Monodentate = one-toothed = only able to donate one lone pair to the metal ion
H2O, OH-, NH3, CN-, SCN-, Cl-
H3N
NH3
+2
Ni
NH3
H3N
3)
4)
Bidentate = 2 bonds from the same ligand
H2C H2N
a) Chelate = claw
b) Ethylenediamine, oxalate
H2C
H2N
NH2
Ni
EDTA = ethylenediaminetetraacetate
i. Very stable complexes
ii. Used to scavenge toxic metals
CH2
NH2
Polydentate
a) Diethylenetriamine: H2NCH2CH2NHCH2CH2NH2
b)
CH2
+2
O
O
O
N
O
+2
Ni
O
N
O O
O
CH2
CH2