1/19/20151 George Mason University General Chemistry 212 Chapter 23 Transition Elements...

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George Mason UniversityGeneral Chemistry 212

Chapter 23Transition Elements

Acknowledgements

Course Text:Course Text: Chemistry: the Molecular Nature of Matter and Chemistry: the Molecular Nature of Matter and

Change, 7Change, 7thth edition, 2011, McGraw-Hill edition, 2011, McGraw-Hill Martin S. Silberberg & Patricia AmateisMartin S. Silberberg & Patricia Amateis

The Chemistry 211/212 General Chemistry courses taught at George Mason are intended for those students enrolled in a science /engineering oriented curricula, with particular emphasis on chemistry, biochemistry, and biology The material on these slides is taken primarily from the course text but the instructor has modified, condensed, or otherwise reorganized selected material.Additional material from other sources may also be included. Interpretation of course material to clarify concepts and solutions to problems is the sole responsibility of this instructor.

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Transition Elements Properties of the Transition Elements

The Inner Transition Elements

Highlights of Selected Transition Elements

Coordination Compounds

Theoretical Basis for the Bonding and Properties of Complexes

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Transition Elements Main-Group vs Transition Elements

Most important uses of Main-Group elements involve the compounds made up of these elements

Transition Elements are highly useful in their elemental or uncombined form

Main –Group Transition Elements

Main-group elements change from metal to non-metal across a period

All transition elements are metals

Most main-group ionic compounds are colorless and diamagnetic (non-magnetic)

Many transition metal compounds are highly colored and paramagnetic

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Transition Elements Properties of Transition Elements

Recall: The “A” (Main Group) elements make up the “s” and “p” blocks

Transition Elements make up the “d” block (B group) “f” block elements (Inner Transition Elements)

As ions, transition metals (elements) provide fascinating insights into chemical bonding and structure

Transition metals play an important role in living organisms

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Transition Elements

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Transition Elements Electron Configurations of the Transition Metals

In the Periodic Table, the Transition metals, designated “d-block (B-Group)” elements, are located in:

40 elements in 4 series within Periods 4 -7

Lie between the last ns-block elements in group [2A(2)] (Ca – Ra) and the first np-block elements in group [(3A(13)] (Ga & element 113 (unnamed)

Each series represents the filling of the 5 d orbitals

l = 2 [ml = -2 -1 0 +1 +2]

(5 orbitals per period x 2 electrons per orbital x 4 Periods

= 40 Elements

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Transition Elements Condensed d-block ground-state electron

configuration:

[noble gas] ns2(n-1)dx, with n = 4 -7; x= 1-10

(several aufbau build-up exceptions)

Partial (valence shell) electron configuration

ns2(n-1)dx

Recall: Chromium (Cr) and Copper (Cu) are exceptions to the above aufbau configuration setup

Expected: Cr [Ar] 4s23d4 Cu [Ar] 4s23d9

Actual: Cr [Ar] 4s13d5 Cu [Ar] 4s13d10

Reasons: change in relative energies of 4s & 3d orbitals and the unusual stability of ½ filled and

filled sublevels (level 4 relative to level 3)

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Transition Elements

Note Aufbau build up exceptions for “Cr” & “Cu”

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Transition Elements The “Inner Transition” elements

Lie between the 1st and 2nd members of the “d-block” elements in Periods 6 & 7 (n=6 & n=7)

Condensed f-block ground-state electron configuration (Periods 6 & 7):

[noble gas] ns2 (n-2)f14(n-1)dx, with n = 6 -7

The 28 “f” orbitals are filled as follows:

l = 3 [ml = -3 -2 -1 0 +1 +2 +3]

7 orbitals per period x 2 electrons per orbital x 2 periods

= 28 Elements

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Transition Elements Transition Metal Ions

Form through the loss of the “ns” electronsbefore the (n-1)d electrons

Ex. Ti2+ [Ar] 3d2 4s2 → [Ar] 3d2 + 2e- (not [Ar] 4s2)

(Ti2+ also called d2 ion) Ions of different transition metals with the same

electron configuration often have similar properties

Ex. Mn2+ and Fe3+ are both d5 ions

Mn2+ [Ar] 3d54s2 → [Ar] 3d5 + 2e-

Fe3+ [Ar] 3d64s2 → [Ar] 3d5 + 3e-

Both Ions have pale colors in aqueous solutions

Both form complex ions with similar magnetic properties

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Practice ProblemWrite condensed electron configurations for the following ions:

Zr V3+ Mo3+

Vanadium (V) – Period 4

Zirconium (Zr) & Molybdenum (Mo) – Period 5

General Configuration: ns2(n-1)dx

a. Zr is 2nd element in the 4d series: [Kr] 5s24d2 (d2 ion)

b. V is the 3rd element in the 3d series: [Ar] 4s23d3

“ns” electrons lost first

In forming V3+, 3 electrons lost – two 4s & one 3d

V3+ = [Ar] 4s23d3 → [Ar] 3d2 (d2 ion) + 3e-

c. Mo lies below Cr in Period 5, Group 6B(6): [kr] 5s1 4d5

Note: Same electron configuration exception as Cr

Mo3+ = [Kr] 5s1 4d5 → [Kr] 4d3 (d3 ion) + 3 e-

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Transition Elements Trends of Transition Elements Across a Period

Transition elements exhibit smaller, less regular changes in Size Electronegativity First Ionization Energy

than the Main Group Elements in the same group

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Transition Elements Atomic Size

General overall decrease across a period for both Main group and Transition group elements

As the “d” orbitals are filled across a period, the change in atomic size within the transition elements evens out because the “d” orbitals are less effective in shielding the outer electrons from the increased nuclear charge Transition Metals Main groupMain group

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Transition Elements Electronegativity

Electronegativity generally increases across period

Change in electronegativity within a series (period) is relatively small in keeping with the relatively small change in size

Small electronegativity change in Transition Elements is in contrast with the steeper increase between the Main Group elements across a period

Magnitude of Electronegativity in Transition elements is similar to the larger main-group metals

Transition Metals

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Transition Elements Ionization Energy

Ionization Energy of Period 4 Main-group elements rise steeply from left to right as the electrons become more difficult to remove from the poorly shielded increasing nuclear charge, i.e., no “d” electrons; thus, electrons held tighter to nucleus

In the Transition metals, however, the first ionization energies increase relatively little because of the combined effects of less effective shielding by the inner “d” electrons and the increasing nuclear chargeTransition Metals

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Transition Elements Trends Within (down) a Group (relative to main-group

elements) Vertical trends differ from those of the Main Group

elements Atomic Size

Increases, as expected, from Period 4 to 5 where electron repulsion dominates the increasing nuclear charge

No increase from Period 5 to 6 The Lanthanide Contraction describes the atomic

radius trend that the Lanthanide series exhibit The Lanthanide Contraction refers to the fact that

the 5s and 5p orbitals penetrate the 4f sub-shell so the 4f orbital is not shielded from the increasing nuclear change, which causes the atomic radius of the atom to decrease

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Transition Elements

n=5

ml = 0

l=0(5s)

-1 0 +1

l=1(5p)

-2 -1 0 +1 +2

l=2 (5d)

-3 -2 -1 0 +1 +2 +3

l=3 (5f)

ml =

0

l=0(1s)

n=1

n=4

ml = 0

l=0(4s)

-1 0 +1

l=1(4p)

-2 -1 0 +1 +2

l=2 (4d)

-3 -2 -1 0 +1 +2 +3

l=3 (4f)

n=3

0

l=0(3s)

l=2 (3d)

-2 -1 0 +1 +2

l=1(3p)

-1 0 +1

n=2

0

l=0(2s)

l=1(2p)

-1 0 +1

n=6,7

ml = 0

l=0(6s,7s)

-1 0 +1

l=1(6p,7p)

-2 -1 0 +1 +2

l=2 (6d)

-3 -2 -1 0 +1 +2 +3

l=3 (6f)

Note: n > 7 &

l > 3 Sublevels

not utilized for

any element in

the current Period Table

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Transition Elements

Order of Sublevel Orbital Filling

Inner Transition Metals

Transition MetalsMain Group Metals

Main Group Non-metals

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Transition Elements Trends Within a Group (relative to main-group elements)

Electronegativity (EN) – Relative ability of an atom in a covalent bond to attract shared electrons EN of Main-group elements decreases down group

greater size means less attraction by nucleus Greater Reactivity

EN in Transition elements is opposite the trend in Main-group elements because of less effective shielding of “d” orbitals

EN increases from period 4 to period 5 No change from period 5 to period 6, since the change

in volume is small and Zeff increases ( weak shielding from f orbital electrons)

Transition metals exhibit more covalent bonding and attract electrons more strongly than main-group metals

The EN values in the heavy metals exceed those of most metalloids, forming salt-like compounds, such as CsAu and the Au- ion

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Transition Elements Trends Within a Group (relative to Main-group

elements)

Ionization Energy – Energy required to remove an electron from a gaseous atom or ion

Main-group elements increase in size down a group, decreasing the Zeff , making it relatively easier to remove the outer electrons

The relatively small increase in the size of transition metals because of ineffective shielding from the increasing nuclear charge (Zeff) by “d” orbital electrons makes it more difficult to remove a valence electron, resulting in a general increase in the first ionization energy down a group

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Transition Elements Trends Within a Group (relative to Main-group

elements) Density

Atomic size (volume) is inversely related to density

(As size increases density decreases) Transition element density across a period initially

increases, then levels off, finally dips at end of series

From Period 5 to Period 6 the density increases dramatically because atomic volumes change little while nuclear mass increases significantly

Period 6 series contains some of the densest elements known:

Tungsten, Rhenium, Osmium, Iridium, Platinum, Gold

(Density 20 times greater than water,

2 times more dense than lead)

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Transition Elements

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Trends are unlike those for the Main-group elements in several ways 2nd & 3rd members of a transition group are nearly same size Electronegativity increases down a transition group 1st ionization energies are highest at the bottom of transition group Densities increase down a transition group (mass increases faster

than density

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Transition Elements Chemical Properties of the Transition Elements

Atomic & physical properties of Transitions elements are similar to Main group elements

Chemical properties of transition elements are very different from main group elements

Oxidation States Main-group elements display one, or at most

two, oxidation states The ns & (n-1)d electrons in transition

elements are very close in energy

All or most can be used as valence electrons in bonding – Transition metals can have multiple oxidation states

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Transition Elements

Oxidation State (Number)Magnitude of charge an atom in a covalent compound would have if its shared electrons were held completely by the atom that attracts them more strongly

Oxidation State

Manganese (Mn)

dx

ElectronicConfiguratio

n

0 d5 [Ar] 4s2 3d5

+1 d5 [Ar] 4s1 3d5

+2 d5 [Ar] 3d5

+3 d4 [Ar] 3d4

+4 d3 [Ar] 3d3

+5 d2 [Ar] 3d2

+6 d1 [Ar] 3d1

+7 d0 [Ar]

4s 3d 4p

Note: All 3 d5

Ex. MnO4- ; O.N. Mn

+7

Ex. MnO2 ; O.N. Mn

+4

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Transition Elements Metallic Behavior

Atomic size and oxidation state have a major effect on the nature of bonding in transition metal compounds

Transition elements in their lower oxidation states behave more like metals – Oxides more basic

Transition elements in their higher oxidation states exhibit more covalent bonding – Oxides more acidic

Ex. TiCl2 (Ti2+) is an ionic solid

TiCl4 (Ti4+) is a molecular liquid

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In the higher oxidation states: The atoms have fewer electrons The nuclear charge pulls remaining electrons closer,

decreasing the volume and increasing the density The charge density (ratio of the ion’s charge to its

volume) increases The increase in charge density leads to more

polarization of the electron clouds in non-metals The bonding becomes more covalent The stronger the covalent bond, the less metallic The oxides, therefore, become less basic

Ex. TiO (Ti2+) is weakly basic in waterTiO2 (Ti4+) is amphoteric, reacting with

both acid and base

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Transition Elements Electronegativity, Oxidation State, Acidity/Basicity

Why does oxide acidity increase with oxidation state?

Metal with a higher oxidation state is more positively charged

Attraction of electrons is increased, i.e., electronegativity increases

Effective Electronegativity = Valence State Electronegativity

EN Cr – 1.6 Al – 1.5 (basic oxide) Cr3+ – 1.7 Cr6+ – 2.3 P – 2.1 (acidic oxides)

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Reduction Strength (Redox) Standard Electrode Potential,

Eo , generally decreases across a period

As the value of Eo becomes more negative, i.e., at the beginning of the series, the ability of the species to act as a reducing agent increases

Thus, Ti2+, Eo = -01.63V, is a stronger reducing agent than Ni2+, Eo = -0.25V

All species with a negative value of Eo can reduce H+

2H+(aq) + 2e- H2(g) Eo = 0.0V) Note: Cu2+ (Eo = +0.34 V) cannot reduce H+

The magnitude of the Eo values between two species, and the relative degree of surface oxidation, determines the level of reactivity of the oxidation/reduction reaction in water, steam, or acid solution

Standard Electrode Potentialsof Period 4 M2+ Ions

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Transition Elements Color in Transition Elements

Most Main-Group Ionic Compounds are colorless

Metal ions have a filled outer shell

With only much higher energy orbitals available to receive an “excited” electron, the ion does not absorb visible light

The partially filled “d” orbitals of the transition metals can absorb visible wavelengths and move to slightly higher energy “d” levels

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Transition Elements Magnetism in Transition Elements

Magnetic properties are related to electron sublevel occupancy

A “Paramagnetic” substance has atoms or ions with “unpaired” electrons

A “Diamagnetic” substance has atoms or ions with only “paired” electrons

Most Main-Group metal ions are diamagnetic (filled outer shells)

Many Transition metal compounds are paramagnetic because of unpaired electron in the “d” subshells

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Transition Elements Chemical Behavior Within a Group

Main_Group

The decrease in Ionization Energy (IE) going down a group results in “increased reactivity”

Transition metals

Ionization Energy increases down group

The Standard Electrode Potential (Eo) also increases (becomes more positive)

Chromium is stronger reducing agent

Some Properties of Group 6B(6) Elements

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Transition Elements The Inner Transition Elements

Lanthanides (Rare Earth Elements)

(Cerium (Ce); Z = 58 – Lutetium (Lu); Z = 71) Silvery, high melting point (800 – 1600oC)

metals Small variations in chemical properties makes

them difficult to separate Occur naturally in the +3 oxidation state as M3+

ions of very similar radii Most lanthanides have the ground-state

electron configuration filling the “f” subshell level

[Xe] 6s2 4fx 5d0 x varies across series (Period) Exceptions – Ce, Gd, Lu have single e- in 5d

orbital

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Sample ProblemFinding the Number of Unpaired Electrons

The alloy SmCo5 forms a permanent magnet because both Samarium and Cobalt have unpaired electrons

How many unpaired electrons are in the Sm atom (Z=62)?

Ans:

Samarium is the eighth element after Xe (Noble Shell)

[Xe] 6s2 4f6

Two (2) electrons go in the 6s sublevel

In general, the 4f sublevel fills before the 5d sublevel (slide 17)

Recall previous slide - only Ce, Gd, Lu have 5d electrons

Remaining 6 electrons go into the 4f orbitals

6s 4f 5d 6pSix unpaired electrons

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Transition Elements The Actinides:

(Thorium (Th); Z=90 - Lawrencium; Z=103)

All Actinides are Radioactive (Alpha (4He2) Decay

Only Thorium & Uranium occur in nature

Share very similar chemical & physical properties

Silvery and chemically reactive

Principal oxidation state is +3, similar to lanthanides

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Transition Elements Highlights of Selected Transition Metals

Period 4 – Chromium & Manganese Chromium

Silvery, shiny metal with many colorful compounds Cr2O3 acts as protective coating on easily corroded

(oxidized) metals, such as iron “Stainless” steels contain as much as 18 % Cr,

making them highly resistant to corrosion Electron Configuration ([Ar] 4s1 3d5) with 6 valence

electrons occurs in all possible positive oxidation states

Important ions Cr2+, Cr3+, Cr6+

Non-metallic character and oxide acidity increase with metal oxidation state

Cr2+ potential reducing agent (Cr loses additional electrons)

Cr6+ potential oxidizing agent (Cr gains electrons)

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Chromium Chromium (II) – Cr2+

CrO is basic and largely ionic Forms insoluble hydroxide in neutral or basic

solution Dissolves in acid to yield Cr2+ ion and water

CrO(s) + 2H+ → Cr2+ (aq) + H2O(l) Chromium(III) – Cr3+

Cr2O3 is amphoteric, similar properties as Aluminum

Dissolves in acid to yield violet Cr3+ ionCr2O3(s) + 6H+(aq) → 2Cr3+(aq) + 3H2O(l)

Reacts with base to form the green Cr(OH)4- ionCr2O3(s) + 3H2O + OH- → 2Cr(OH)4

-(aq)

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Chromium (con’t)

Chromium (VI) - Cr6+ (Deep Red)

CrO3 is covalent and acidic

Dissolves in water to form Chromic Acid (H2CrO4)

CrO3(s) + H2O(l) → H2CrO4(aq)

H2CrO4 yields yellow Chromate ion (CrO42-) in

base

H2CrO4(aq) + 2OH-(l) → CrO42-(aq) + 2H2O(l)

Chromate ion forms orange dichromate (Cr2O7

2-) ion in acid

2CrO42-(aq) + 2H+(aq) ⇆ Cr2O7

2-(aq) H2O(l)

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Manganese Hard and Shiny Like Vanadium & Chromium used to make steel

alloys Chemistry of Manganese is similar to Chromium Metal reduces H+ from acids to form Mn2+ ionMn(s) + 2H+(aq) → Mn2+(aq) + H2(g) Eo = 1.18 V

Manganese can use all its valence electrons (several oxidation states) to form compounds Mn2+ Mn4+ Mn7+ most important

As oxidation state rises from +2 to +7, the valence state electronegativity increases and the oxides of Mn change from basic to acidic Mn(II)O (basic) Mn(III)2O3 (amphoteric) Mn(IV)O2 (insoluble) Mn(VII)2O7 (acidic)

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Transition Elements All Manganese species with oxidation states greater

than +2 act as oxidizing agents (gaining the electrons lost by the atoms being oxidized)

Mn7+O4-(aq) + 4H+ + 3e- → Mn4+O2(s) + 2H2O(l) Eo =

1.68

Mn7+O4-(aq) + 2H2O + 3e- → Mn4+O2(s) + 4OH- Eo =

0.59

(Mn7+O4- is a much stronger oxidizing agent in acid

solution than in basic solution – note difference in Eo values)

Oxidation State

Manganese (Mn)

dx

ElectronicConfigurati

on

0 d5 [Ar] 4s2 3d5

+1 d5 [Ar] 4s1 3d5

+2 d5 [Ar] 3d5

+3 d4 [Ar] 3d4

+4 d3 [Ar] 3d3

+5 d2 [Ar] 3d2

+6 d1 [Ar] 3d1

+7 d0 [Ar]

4s 3d 4p

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Transition Elements Manganese

Unlike Cr2+ & Fe2+, the Mn2+ (3d5) ion resists oxidation in air Recall: half-filled (-1/2 spin electrons missing)

& filled sublevels are more stable than partially filled sublevels

Cr2+ is a d4 species and readily loses a 3d electron to form the d3 ion Cr3+, which is more stable

Fe2+ is a d6 species and removing a 3d electron yields the stable, half-filled d5 configuration of Fe3+

Removing an electron from Mn2+ disrupts the more stable d5 configuration

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Transition Elements & TheirCoordination Compounds

Coordination Compounds (Complexes) Most distinctive aspect of transition metal chemistry Complex – Substances that contain at least one

complex ion Complex ion – Species consisting of a “central metal

cation” (either a main-group or transition metal) that is bonded to molecules and/or anions called “Ligands”

The Complex ion is typically associated with other (counter) ions to maintain neutrality

A coordination compound behaves like an electrolyte in water Complex ion and counter ion separate Complex ion behaves like a polyatomic ion – the

ligands and central atom remain attached

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Components of Coordination Compound When solid complex dissolves in water, the complex

ion and the counter ions separate, but ligands remain bound to central atom

CentralAtom

Ligands CounterIons

[Co(NH3)6]Cl3(s)

OctahedralGeometry

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Complex ions A complex ion is described by the metal ion

and the number and types of ligands attached to it The bonding between metal and ligand

generally involves formal donation of one or more of the ligand's electron pairs

The metal-ligand bonding can range from covalent to more ionic

Furthermore, the metal-ligand bond order can range from one to three (single, double, triple bonds)

Ligands are viewed as Lewis Bases (donate electron pairs), although rare cases are known involving Lewis acidic ligands

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Complex ions The complex ion structure is related to

three characteristics: Coordination Numbers

The number of ligand atoms that are bonded directly to the central metal ion

Coordination number is specific for a given metal ion in a particular oxidation state and compound

Coordination number in [Co(NH3)6]3+ is 6

The most common coordination number in complex ions is 6, but 2 and 4 are common, with a few higher

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Complex ions Geometry – Depends on Coordination No. & Nature of

Metal IonMetal ion

CN Shape dx

Cu+ 2 Linear d10

Ag+ 2 Linear d10

Au+ 2 Linear d10

Ni2+ 4Octahedral Sq

Planard8

Pd2+ 4Octahedral Sq

Planard8

Pt2+ 4Octahedral Sq

Planard8

Cu2+ 4Octahedral Sq

Planard9

Cu3+ 4 Tetrahedral d8

Zn2+ 4 Tetrahedral d10

Cd2+ 4 Tetrahedral d10

Mn2+ 4 Tetrahedral d5

Ti3+ 6 Octahedral d1

V2+ 6 Octahedral d3

Cr3+ 6 Octahedral d3

Mn2+ 6 Octahedral d5

Fe3+ 6 Octahedral d5

Co3+ 6 Octahedral d6

d10

d8

d9

d1

d3

d5

d6

Coordination Numbers and Shapes of Some Complex Ions

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Complex Ions

Donor Atoms per Ligand

The Ligands of complex ions are “molecules” or “anions” with one or more donor atoms that each donate a lone pair of electrons to the metal ion to form a covalent bond

Atoms with lone pairs of electrons often come from Groups 5A, 6A, or 7A (main-group elements)

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Complex Ions

Ligands are classified in terms of the number of donor atoms (teeth) that each uses to bond to the central metal ion

Monodentate Ligands use a “single” donor atom

Bidentate Ligands have two donor atoms

Polydentate Ligands have more than two donor atoms

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The Ligands contains one or more Donor atoms that have electron pairs to donate to

the Central Atom

Donor Atom

Some Common Ligands in Coordination Compounds

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Complex Ions Chelates (Greek “chela” – crab’s claw)

Bidentate and Polydentate ligands give rise to “rings” in the complex ion

Ex: Ethylene Diamine (abbreviated (en) in formulas)(:N – C – C – N:)

forms a 5-member ring, with the two electron donatingN atoms bonding to the metal atomSuch ligands seem to grab the metal ion like claws

Ethylenediaminetetraacetate (EDTA)

Used in treating heavy-metal poisoning, by acting as a scavenger of lead and other heavy-metal ions, removing them from blood and other body fluids

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Formulas and Names of Coordination Compounds Important rules for writing formulas of

coordinate compounds The cation is written before the anion The charge of the cation(s) is balanced

by the charge of the anions In the complex ion, neutral ligands are

written before anionic ligands The entire ion is placed in brackets, i.e.,

[ ]

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Formulas and Names of Coordination Compounds

Coordination Compound Formulas

Example # 1

Two compound cations (K+) – Total Charge +2Ion Central Metal Cation (Co2+) – Total Charge +2Neutral Ligands (2 NH3) – Total Charge 0

Counter Ions (4 Cl-) – Total Charge -4Net Charge on Complex Ion – - 2 [Co(NH3)2Cl4]-2

-2+ 2+ -2 3 2 4K [Co (NH ) Cl ]

2 3 2 4K [Co(NH ) Cl ]

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Coordination Compound Formulas

Example # 2 – Complex Ion and Counter Ion

[Co(NH3)4Cl2]Cl

Counter Ion (Cl-) (not part of complex ion) – Total charge -1

Complex Ion - Neutral Ligands (4 NH3) – Total Charge 0

Complex Ion - Anion Ligands (2 Cl-) – Total Charge -2

Complex Ion - [Co(NH3)4Cl2]+ – Total Charge +1

Complex Ion - Central Metal Atom (Co) – Total Charge +3

[Co3+(NH3)4Cl-2]+Cl-

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Formulas and Names of Coordination Compounds Example #3 – Complex Cation and Complex Anion

[Co(NH3)5Br]2[Fe(CN)6] Complex Cation - [Co(NH3)5Br]2+

Complex Cation Central Atom (Co+3) – Total charge +3 Complex Cation Neutral Ligands (5 NH3) – Total

Charge 0

Complex Cation Anionic Ligand (Br-) – Total Charge -1 Complex Anion ([Fe(CN)6]4-) – Total Charge -4

Complex Anion Central Cation (Fe2+) – Total Charge +2

Complex Anion Ligand (6 CN-1) – Total Charge -6 [Co3+(NH3)5Br-]2 [Fe2+(CN-)6]

2 x (3 - 1) = 4 2 - 6 = -4

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Formulas and Names of Coordination Compounds Naming Coordination Compounds

RulesThe Cation is named before the AnionWithin the Complex Ion, the Ligands are named, in

alphabetical order, before the metal ionNeutral Ligands generally have the molecule

name, with exceptions Ex NH3 (ammine), H2O (aqua), C=O (carbonyl)

Anionic Ligands drop the –ide and add –o after the root name Ex. Cl- becomes “chloro”

A numerical prefix indicates the number of ligands of a particular type Ex di (2), tri (3), tetra (4)

[Co(NH3)4Cl2]Cl

Tetra ammine di chloro cobalt(III)chloride

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Symbol

FeCuPbAgAuSn

Di Bis II

Tri Tris III

Tetra Tetrakis IV

Penta pentakis V

Hexa Hexakis VI

Septa Septakis VII

Names of Some Neutraland Anionic Ligands

Names of Some Metals Ionsin Complex Anions

Numerical Prefixes usedIn Complex Anions

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RulesSome ligand names already contain a

numerical prefix

Ethylenediamine

In these cases the number of ligands is indicated by such terms as:

bis (2) tris(3) tetrakis(4)

A compound with two ethylene ligands would contain the following ligand name

bis(ethylenediamine)

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Rules The oxidation state of the central metal ion is

given by a Roman numeral (in parentheses) only if the metal ion can have more than one state, as in the compound

[Co(NH3)4Cl2]Cl [Co3+(NH3)4Cl-2]Cl-

Tetra ammine di chloro cobalt(III)chloride If the complex ion is an anion, drop the ending

of the Central metal name and add “–ate”K[Pt(NH3)Cl5] K+[Pt4+(NH3)Cl-5]

-

Potassium ammine penta chloro platinate(IV)

Na4[FeBr6] Na+4[Fe2+Br-

6]

Sodium hexa bromo ferrate(II)

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Practice ProblemWhat is the systematic name of Na3[AlF6]?

Ans: Complex ion – [AlF6]3-

Ligands 6 (hexa) F- ions (Fluoro)

Complex ion is an “anion” (net charge -3)

End of metal ion Aluminum must be changed to –ate

Complex ion name – hexafluoroaluminate

Aluminum has only the +3 oxidation state so Roman numerals are not required

Na3+ is the positive counter ion; it is separated from the complex anion by a space

Na3[AlF6] Sodium Hexfluoroaluminate

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Practice ProblemWhat is the systematic name of [Co(en)2Cl2]NO3?

Ans: Listed alphabetically, there are two Cl- (dichloro) and two “en” [bis(ethylenediamine)] ligands

Note: Alphabetically refers to the root chemical names:

Chloro & Ethylenediamine

The “Complex ion” is a “Cation,” with a charge of +1

[Co3+(en)2Cl-2]+

The metal name in a complex ion is unchanged - Cobalt

Because Cobalt can have several oxidation states,its charge must be specified - Cobalt (III)

One Nitrate ion (NO-3) balances the +1 complex cation

Dichloro bis (ethylene diamine)cobalt(III) nitrate

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Practice ProblemWhat is the formula of:

Tetra ammine bromo chlroro platinum(IV) chloride

Ans: The central atom of the complex cation is written first

Platinate(IV) Pt4+

The ligands follow in alphabetical order of root chemical name

Tetraammine (NH3) Bromo (Br-) Chloro (Cl-)

Complex ion formula - [Pt(NH3)4BrCl]2+ [Pt4+(NH3)4Br-Cl-]2+

To balance the +2 charge of the complex cation,2 Cl- counter ions are required

[Pt(NH3)4BrCl]Cl2

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Practice ProblemWhat is the formula of

Hexa ammine cobalt(III) tetra chloro ferrate(III)

Ans: Compound consists of two complex ions

Complex Cation – Six hexammine (NH3) & cobalt(III) (Co3+)

Complex Cation – [Co(NH3)6]3+ [Co3+(NH3)6]3+

Complex Anion – tetrachloro - 4 Cl-

Complex Anion – ferrate(III) - Fe3+

Complex Anion – [FeCl-4]-

Complex cation – balanced by 3 complex anions

Coordinate Compound – [Co(NH3)6][FeCl4]-3

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Isomerism in Coordination Compounds Isomers are compounds with the same chemical

formula but different properties Constitutional (Structural) Isomers

Two compounds with the same formula, but with atoms connected differently Two Types

Coordination Isomers – Composition of the complex ion changes but not the compoundEx. Ligand and counter ion exchange positions

[Pt(NH3)4Cl2](NO2)2 [Pt(NH3)4(NO2)2]Cl2Ex. Two sets of ligands reversed

[Cr(NH3)6][Co(CN)6] [Co(NH3)6][Cr(CN)6]

(NH3 is ligand of Cr3+ in one compound and of Co3+ in the other)

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Constitutional (Structural) Isomers Linkage Isomers

Composition of the complex ion remains the same, but the attachment of the ligand donor atom changes

Some ligands can bind to the metal ion through either of two donor atomsEx. pentaamminenitrocobalt(III) chloride

[Co(NH3)5(NO2]Cl2

pentaamminenitritocobalt(III) chloride[Co(NH3)5(ONO]Cl2

Ex. Cyanate ion can attach via lone pair of electrons on

the Oxygen atom (NCO:)or the Nitrogen atom (isocyanato (OCN:)

Other examples of alternate electrondonor pairs for Linkage IsomerS

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Constitutional (Structural) Isomers Stereo Isomers

Compounds that have the same atomic connections but different spatial arrangements of the atoms

Geometric Isomers (cis-trans isomers [diastereomers])

Atoms or groups of atoms arranged differently in space relative to the “Central” metal

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Constitutional (Structural) Isomers Stereo Isomers Optical Isomers (enantiomers)

Occur when a molecule and its mirror image can not be superimposed

Optical isomers have distinct physical properties like other types of isomers, with one exception – the direction in which they rotate the plane of polarized lightOptical isomerism in an octahedral complex ion

Rotating structure I in the cis compound gives structure III, which is not the same as structure II, its mirror image, Image I & Image III are optical isomers

Rotating structure I in the trans compound gives structure III,which is the same as structure II, its mirror image, The trans compound does not have any mirror images

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Practice ProblemDraw all stereo isomers for the following [Pt(NH3)2Br2] Cr(en)3]3+ (en =

H2NCH2CH2NH2)

Pt

NH3Br

H3N Br

Pt

H3N Br

H3N Brtrans

Pt(II) complex is Square Planar GeometryTwo different monodentate ligandsGeometric Isomers Each isomer is superimposable on the mirror image – no optical isomerism

Ethylenediamine is a bidentate ligand

The Cr3+ has a coordination number of 6 and an octahedral geometry, similar to Co3+

The three bidendate ions are identical

No geometric isomerism

This complex ion has a nonsuperimposable mirror image

Optical Isomerism does occur

cis

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Theoretical Basis for the Bonding and Properties of Complexes

Questions

How do Metal Ligand bonds form

Why certain geometries are preferred

Why are complexes often brightly colored

Why are complexes often paramagnetic – attracted to a magnetic field as a result of their electron pairs being unpaired

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Theoretical Basis for the Bonding and Properties of Complexes Application of Valence Bond Theory to Complex Ions

In the formation of a complex ion, the filled ligand orbital overlaps the empty metal-ion orbital

The Ligand (Lewis Base) donates the electron pair and the metal-ion (Lewis Acid) accepts it to form one of the covalent bonds of the complex ion (Lewis adduct)

When one atom in a bond donates both electrons the bond is referred to as a ”coordinate covalent bond”

The number and type of metal-ion hybrid orbitals occupied by ligand lone pairs determine the geometry of the complex ion

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Application of Valence Bond Theory to Complex Ions Octahedral Complexes (six electron groups about central

atom) Ex. Hexaamminechromium(III) ion [CrNH3)6]3+

Six hybrid orbitals are needed to make the ion The six lowest energy orbitals of the Cr3+ ion

Two 3d, one 4s, three 4p

mix and become six equivalent d2sp3 hybrid orbitals that point to the corners of an octahedron

The six d2sp3 hybrid orbitals are filled with the six electron pairs from the six NH3 ligands

Note the lowest 6 energy levels for Cr3+ involve both n=3 & n=4 sublevelsThe 3d orbitals are of lower energy than the 4s and 4p orbitalsThe hybrid designation, d2sp3, follows this orderIf all the orbitals had the same “n” value, the order would have been sp3d2

ParamagneticUnpaired e-

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Application of Valence Bond Theory to Complex Ions Square Planar Complexes (four electron groups about

central atom) Metal ions with a d8 configuration usually form square

planar complexes In the [Ni(CN)4]2- ion, the model proposes

one 3d, one 4s, two 4p for Ni2+

to from four dsp2 hybrid orbitals pointing the corners of a square accepting one electron pair from each of the four CN- orbitals

Note the filling of the first 4 unhybridized 3d orbitals after one 3d, one 4s and two 4p orbitals combine to form the four dsp2 hybrid orbitals

ParamagneticUnpaired e-

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Application of Valence Bond Theory to Complex Ions Tetrahedral Complexes (four electron groups about central

atom) Metal ions that have a filled d sublevel, such as Zn+2

[Ar] 3d10

often form Tetrahedral complexes In the [Zn(OH)4]2- ion, the model proposes the lowest

available Zn2+ orbitals

one 4s, three 4p

mix to become four sp3 hybrid orbitals that point to the corners of a tetrahedron, occupied by four lone pairs, one from each of the four OH- ligands

Diamagnetic

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Crystal Field Theory

Valence Bond Theory pictures and rationalizes bonding and shape of molecules

VB theory gives little insight into the colors of coordination compounds and can be ambiguous with regard to magnetic properites

Crystal Field Theory explains color and magnetism

Highlights the “effects” on the d-orbital energies of the metal ion as the ligands approach

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What is Color?

White light is electromagnetic radiation consisting of “all” wavelengths () in the “visible” range

Objects appear “colored” in white light because they absorb certain wavelengths and reflect or transmit others

Opaque objects reflect light

Clear objects transmit light

If the object absorbs all visible wavelengths, it appears “black”

If the object reflects all visible wavelengths, it appears “white”

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Crystal Field Theory What is Color?

Each color has a “complimentary” color An object has a particular color for two

reasons It reflects (or transmits) light of that color or It absorbs light of the “complimentary”

color

Ex. If an object absorbs only red (compliment of green), it is interpreted as “green” Colors with approximate wavelength ranges

Complimentary colors, such as red and green,lie opposite each other

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In CF Theory, the properties of complexes result from the splitting of d-orbital energies

Split d-orbital energies arise from “electrostatic” interactions between the positively charged metal ion cation and the negative charge of the ligands

The negative charge of the ligand is either partial as in a polar neutral ligand like NH3, or full, as in an anionic ligand like Cl-

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Crystal Field Theory The ligands approach the metal ion along the mutually

perpendicular x, y, and z axes (octahedral orientation), minimizing the overall energy of the system

B & C Lobes of the dx2-y2 and dz2 orbitals lie directly in line with the approaching ligands and have stronger repulsions

D, E, F lobes of the dxy, dxz, and dyz orbitals lie “between” the approaching ligands, so the repulsion are weaker

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Crystal Field Theory An energy diagram of the orbitals shows all five d orbitals are

higher in energy in the forming complex than in the free metal ion, because of the repulsions from the approaching ligands

Crystal Field Splitting Energy - The d orbital energies are“split” with the two dx2-y2 and dz2 orbitals (eg orbital set) higher in energy than the dxy, dxz, and dyz orbitals (t2g orbital set)

Strong-field ligands, such as CN- lead to larger splitting energy Weak-field ligands such as H2O lead to smaller splitting energy

Crystal Field Splitting Energy

Forming Complex

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Explaining the Colors of Transition Metals

Diversity in colors is determined by the energy difference () between the t2g and eg orbital sets in complex ions

When the ions absorbs light in the visible range, electrons move from the lower energy t2g level to the higher eg level, i.e., they are “excited” and jump to a higher energy level

E electron = Ephoton = hv = hc/

The substance has a “color” because only certain wavelengths of the incoming white light are absorbed

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Crystal Field Theory Example – Consider the [Ti(H2O)6]3+ ion – Purple in

aqueous solution Hydrated Ti3+ is a d1 ion, with the d electron in one of the

three lower energy t2g orbitals The energy difference (A) between the t2g and eg

orbitals corresponds to the energy of photons spanning the green and yellow range

These colors are absorbed and the electron jumps to one of the eg orbitals

Red, blue, and violet light are transmitted as purple

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For a given “ligand”, the color depends on the oxidation state of the metal ion – the number of “d” orbital electrons available

A solution of [V(H2O)6]2+ ion is violet

A solution of [V(H2O)6]3+ ion is yellow

For a given “metal”, the color depends on the ligand

[Cr(NH3)6]3+ (yellow-orange)

[Cr(NH3)5]2+ (Purple)

Even a single ligand is enough to change the color

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Crystal Field Theory Spectrochemical Series

The Spectrochemical Series is a ranking of ligands with regard to their ability to split d-orbital energies

For a given ligand, the color depends on the oxidation state of the metal ion

For a given metal ion, the color depends on the ligand As the crystal field strength of the ligand increases, the

splitting energy () increases (shorter wavelengths of light must be absorbed to excite the electrons

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Practice ProblemRank the following ions in terms of the relative value of and of the energy of visible light absorbed

[Ti(H2O)6]3+ Ti(NH3)6]3+ Ti(CN)6]3+

Ans:

Oxidation State of Ti is +3 in all formulas

From the spectrochemical series table, the ligand strength is in the order:

CN- > NH3 > H2O

Relative size of , thus, the energy of light absorbed is

Ti(CN)6]3+ > Ti(NH3)6]3+ > [Ti(H2O)6]3+

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Explaining the Magnetic Properties of Transition Metal Complexes The splitting of energy levels influence magnetic

properties Affects the number of unpaired electrons in the

metal ion “d” orbitals According to Hund’s rules, electrons occupy orbitals

one at a time as long as orbitals of “equal energy” are available

When “all” lower energy orbitals are “half-filled (all +½ spin state)”, the next electron can Enter a half-filled orbital and pair up (with a –½ spin

state electron) by overcoming a repulsive pairing energy (Epairing) or

Enter an empty, higher energy orbital by overcoming the crystal field splitting energy ()

The relative sizes of Epairing and () determine the occupancy of the d orbitals

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Crystal Field Theory Explanation of Magnetic Properties

The occupancy of “d” orbitals, in turn, determines the number of unpaired electrons, thus, the paramagnetic behavior of the ion

Ex. Mn2+ ion ([Ar] 3d5) has 5 unpaired electrons in 3d orbitals of equal energy

In an octahedral field of ligands, the orbital energies split

The orbital occupancy is affected in two ways: Weak-Field ligands (low ) and High-Spin

complexes Strong-Field ligands (high ) and Low-Spin

complexes(from spectrochemical series)

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Weak-Field ligands and High-Spin complexes Ex. [Mn(H2O)6]2+ Mn2+ ([Ar] 3d5) A weak-field ligand, such as H2O, has a “small” crystal

field splitting energy () It takes less energy for “d” electrons to move to

the “eg” set (remaining unpaired) rather thanpairing up in the “t2g” set with its higherrepulsive pairing energy (Epairing)

Thus, the number of unpaired electrons in aweak-field ligand complex is the same as inthe free ion

Weak-Field Ligands create high-spin complexes,those with a maximum of unpaired electrons

Generally Paramagnetic

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Strong-Field Ligands and Low-Spin Complexes Ex. [Mn(CN)6]4-

Strong-Field Ligands, such CN-, cause large crystal field splitting of the d-orbital energies, i.e., higher ()

() is larger than (Epairing) Thus, it takes less energy to pair up in the “t2g“ set than

would be required to move up to the “eg” set The number of unpaired electrons in a

Strong-Field Ligand complex is less thanin the free ion

Strong-Field ligands create low-spin complexes,i.e., those with fewer unpaired electrons

Generally Diamagnetic

Fewerunpaired electrons

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Crystal Field Theory Explaining Magnetic Properties

Orbital diagrams for the d1 through d9 ions in octahedral complexes show that both high-spin and low-spin options are possible only for:

d4 d5 d6 d7 ions With three “lower” energy t2g orbitals available,

the d1, d2, d3 ions always form high-spin (unpaired) complexes because there is no need to pair up

Similarly, d8 & d9 ions always form high-spin complexes because the 3 orbital t2g set is filled with 6 electrons (3 pairs)The two t2g orbitals must have either two d8 or one d9 unpaired electron

04/21/23 87

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Explaining Magnetic Properties

high spin: weak-field

ligand

low spin: strong-

field ligand

high spin: weak-field

ligand

low spin: strong-

field ligand

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Practice ProblemIron(II) forms an essential complex in hemoglobin

For each of the two octahedral complex ions

[Fe(H2O)6]2+ [Fe(CN)6]4-

Draw an orbital splitting diagram, predict the number of unpaired electrons, and identify the ion as low-spin or high spin

Ans:

Fe2+ has the [Ar] 3d6 configuration

H2O produces smaller crystal field splitting () than CN-

The [Fe(H2O)6]2+ has 4 unpaired electrons (high spin)

The [Fe(CN)6]4- has no unpaired electrons (low spin)

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Crystal Field Theory Four electron groups about the central atom

Four ligands around a metal ion also cause d-orbital splitting, but the magnitude and pattern of the splitting depend on the whether the ligands are in a “tetrahedral” or “square planar” arrangement

Tetrahedral – AX4 Octahedral – AX4E2 (2 ligands along “z” axis

removed)

Splitting of d-orbital energies by

a square planar field of ligands.

Splitting of d-orbital energies by a tetrahedral field of ligands

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Crystal Field Theory (Splitting) Tetrahedral Complexes

Ligands approach corners of a tetrahedron None of the five metal ion “d” orbitals is directly in

the path of the approaching ligands Minimal repulsions arise if ligands approach the

dxy, dyz, and dyz orbitals closer than if they approach thedx2-y2 and dz2 orbitals (opposite of octahedral case)

Thus, the dxy, dyz, and dyz orbitals experience more electron repulsion and become higher energy

Splitting energy of d-orbital energies is “less” in a tetrahedral complex than in an octahedral complex

tetrahedral < octahedral

Only high-spin tetrahedral complexes are known because the magnitude of () is small (weak)

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Crystal Field Theory (Splitting) Square Planar Complexes

Consider an Ocatahedral geometry with the two z axis ligands removed, no z-axis interactions take place

Thus, the dz2, dxz an dyz orbital energies decrease The two ‘d” orbitals in the xy plane (dxy, dx2-y2)

interact most strongly with the approaching ligands The (dxy, dx2-y2) orbital has its lobes directly on the

x,y axis and thus has a higher energy than the dxy orbital

Square Planar complexes are generally strong field – low spin and generally diamagnetic

D8 metals ions such as [PdCl4]2- have 4 pairs of the electrons filling the lowest energy levels and are thus, “diamagentic”

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1/19/20151 George Mason University General Chemistry 212 Chapter 23 Transition Elements...

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