CHEMISTRY XL-14A MOLECULAR SHAPE AND STRUCTURE PART II
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CHEMISTRY XL-14A
MOLECULAR SHAPE
AND STRUCTU
REPART II
July 30, 2011 Robert Iafe
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Midterm Results2
Average: 55.8StDev: 20.3High: 91.5Low: 8
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Midterm 13
1. How many seconds are in 65 years? a) 3.3 x 107 secb) 3.4 x 107 secc) 2.0 x 109 secd) 2.1 x 109 sec
2. The molecular formula of hydroquinone is C6H6O2. What is the empirical formula of hydroquinone?
a) C9H9O3
b) C6H6O2
c) C3H3Od) CHO
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Midterm 14
3. For the following set of four quantum numbers (n, l, ml, ms) identify the set that is not valid: a) {9,5,-3,½}b) {3,3,3,½}c) {2,0,0,½}d) {4,2,-2,½}
4. How many electrons are in principle quantum number n = 3?a) 3b) 6c) 9d) 18
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Midterm 15
5. How many radial nodes are in a 3d orbital?
a) 0b) 1c) 2d) 3
6. Which of the following is isoelectronic with N3-
a) Fb) Cl-1
c) Mg2+
d) Ar
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Midterm 16
7. Which of the compounds below has bonds with more covalent character?a) CaOb) Li2Oc) MgOd) MgS
8. Which of the following is the most electronegative?a) Bb) Inc) Ted) O
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Midterm 17
9. Which of the following has the largest atomic radius?a) Beb) Sc) Ted) Sr
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Midterm 18
10. 1,3-Benzodioxol-5-ol, otherwise known as sesamol, is a natural organic compound found in sesame seed oil. The combustion analysis of a 100 mg sample of 1,3-benzodioxol-5-ol shows its composition is: 60.87% C, 4.38% H, and 34.75% O. a) What is the empirical formula of 1,3-benzodioxol-5-
ol? b) If the molecular weight of 1,3-benzodioxol-5-ol is
138.12 g/mol, what is the molecular formula?
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Midterm 19
11. Sodium metal reacts readily with water to produce hydrogen gas and sodium hydroxide in solution. a) Write a balanced reaction.b) If 3.00 x 10-1 grams of sodium react with
0.500 L of water (density = 1 g/mL), what is the theoretical yield of hydrogen gas?
c) If exactly 1.20 x 10-2 grams of hydrogen gas is collected, what is the percent yield?
d) Knowing the percent yield of the reaction, what is the predicted concentration of NaOH in the solution assuming the volume of water does not change?
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Midterm 110
12. Write the noble gas electron configuration for the following atoms/ions. How many unpaired electrons does each have? a) Pb) Ag+1
13. Identify the M3+ ion for each ground state electron configuration [Ar]3d7
[Ar]3d10
14. Briefly explain why oxygen has a lower first ionization energy than nitrogen.
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Midterm 111
15. Calculate the wavelength of the radiation emitted by a hydrogen atom (one electron atom) when an electron makes a transition from n = 5 to n = 2.
16. The work function for chromium metal is 4.37 eV or 7.00 x 10-19 J. Light with a wavelength of 11.0 nm is shined on the metal. What is the velocity of the ejected electron?
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Midterm 112
17. Draw the boundary surface for the three 3p orbitals, making sure to name and label each axis correctly. How many radial nodes does a 3p orbital have?
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Midterm 113
Provide the Lewis dot structure for the following compounds. Make sure every electron is drawn. Indicate atoms with formal charge. If resonance structures can be drawn, draw all valid structures. Make sure to indicate formal charge when necessary.a) CH3CH2CH2OH b) CN- c) HNO3
d) POCl3
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Valence shell electron pair repulsion theory (VSEPR)
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Lewis Theory Connectivity, electron tracking
VSEPR Theory 3-D Structure around an atom
Valence Bond Theory Molecular Orbital Theory
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Review – VSEPR Theory15
Valence-Shell Electron-Pair Repulsion (VSEPR) Accounts for 3D shapes of molecules Based on electron-electron repulsion Determine bond angles shape
Rules are based on experimental observation:
1. Areas of electron concentration (bonds and lone pairs) around the central atom repel each other.
2. Bonds and lone pairs stay as far away from each other as possible (without changing distance)
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2 3 4 5 6
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Review - Polar Molecules17
A polar covalent bond has a non-zero dipole moment
Polar molecules – molecules with non-zero dipole moment: ≠ 0 D
Examples: HCl HF HBr
A non-polar molecule has no dipole moment: = 0 D
Examples: O2 CH4 SF6
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Polar Molecules18
Polar bonds, Non-polar molecule?
If the dipole moments cancel out (Vector sum = 0):
Each C-O bond is polar
Two equally polar bonds, with opposing dipole moments
CO2
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Polar Molecules19
H2OEach O-H bond is polar
Overall molecule is ??????
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Polar Molecules20
C2H2Cl2
cis-CH2Cl2 trans-CH2Cl2
non-polar
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Polar Molecules21
non-polar polar
CCl4 CHCl3
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Valence Bond (VB) Theory22
Lewis Theory Connectivity, electron tracking
VSEPR Theory 3-D Structure around an atom
Valence Bond Theory Extended 3-D Structure Information Delocalization in Molecules Illustrates Multiple Bonding Prediction of Reactivity
Molecular Orbital Theory
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Answers to HΨ=EΨ23
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Answers to HΨ=EΨ24
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Bonding With Orbitals25
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Bonding with Orbitals26
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Valence-Bond Theory27
So far, we’ve been thinking of molecules using Lewis’ Theory:
Bonding electrons are located in between bonded atoms – electrons are localized
But….
We learned in Ch. 1 that we learned to think of electrons as wavefunctions, which are described by atomic orbitals
Valence Bond Theory – quantum mechanical view of bonding
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‘Types’ of Bonds – σ bond28
Lets start with H2, the simplest molecule:
Ground state H has one 1s electron
When the 2 H atoms bond, the atomic orbitals merge, forming a σ-bond
σbond (sigma bond) – along bond axis.
We say the atomic orbitals overlap
More overlap = Stronger bond
All single covalent bonds are σ-bonds
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‘Types’ of Bonds – σ bond29
All single covalent bonds are σ-bonds
Can have σ-bonds between any types
of orbitals:
Two s orbitals
Two p orbitals
An s and a p orbital
Etc….
Example: σ-bond between 1s orbital of H and 2pz orbital of F
H –– F
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‘Types’ of Bonds – Multiple bonds
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What happens in N2?
Remember the bond in N2 is a triple bond…
Each N atom has 3 unpaired 2p electrons
Lets look at the atomic orbitals of N:
pzpypx
N –– N
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‘Types’ of Bonds – Multiple bonds
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What happens in N2?
Remember the bond in N2 is a triple bond…
Each N atom has 3 unpaired 2p electrons
How do the atomic orbitals bond?
pz pz
Between the 2pz and 2pz orbitals, we have bonding on the bond axis
This is a σ-bond!
N –– N
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‘Types’ of Bonds – Multiple bonds
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What happens in N2?
Remember the bond in N2 is a triple bond…
Each N atom has 3 unpaired 2p electrons
How do the atomic orbitals bond?
Between 2px and 2px orbitals, the bonding is not on the bond axis
A different type of bond!
π-bond –nodal plane along bond axis
Bonding occurs above and below the bond axis
px px
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Bonding of N233
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Bonding of N234
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Hybridization of Orbitals35
What about methane (CH4)?
According to Valence Bond Theory:
C should only make 2 bonds!
But we know that C can make 4 bonds
And CH4 has a tetrahedral shape according to VSEPR:
How do get a tetrahedral shape from the 2s and 2p orbitals?
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Hybridization of Orbitals36
We solve the 4 bond problem by promoting an electron:
It takes energy to promote an electron to a higher E orbital
But, overall the energy is lower if C can make 4 bonds instead of 2.
Ok, what about the geometry problem?
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Hybridization of Orbitals37
By promoting an electron, we can now make 4 bonds
But the geometry of the px, py, and pz orbitals don’t match the tetrahedral shape of CH4
s and p orbitals are described by a wave-like model of the e
If we think of the orbitals as interfering with each other, we can define new hybrid orbitals:
h1 = s + px + py + pz
h2 = s – px – py + pz
h3 = s – px + py – pz
h4 = s + px – py – pz
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sp3 Hybridization38
C in CH4 uses 4 hybrid orbitals:
We took 1 s orbital and 3 p orbitals to make 4 sp3 orbitals.
The sum of the atomic orbitals = the sum of the hybrid orbitals
The 4 sp3 orbitals point in the 4 directions of the tetrahedral bonds
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sp3 Hybridization39
Back to CH4:
The 4 2sp3 hybrid orbitals on carbon make σ-bonds with the 1s orbitals of the 4 H atoms
NH3 also uses sp3 hybrid orbitals:
Whenever an atom has a tetrahedral structure, it is sp3 hybridized
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sp3 Hybridization40
Whenever an atom has a tetrahedral structure, we say it is sp3 hybridized
This includes molecules with multiple central atoms:
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Atomic Orbitals Hybrid Atomic Orbitals
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psh1 psh2
sp3 bonding works for tetrahedral-shaped molecules
What about the other VSEPR shapes?
Linear bonding can be described by sp hybridization:
2 of the p orbitals remain as they were
Unused p orbitals are available for π bonding
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Other Types of Hybridization
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sp3 bonding works for tetrahedrally shaped molecules
What about the other VSEPR shapes?
Trigonal planar can be described by sp2 hybridization:
The pz orbital is not used and remains as it was.
y1 p2sh
yx2 p2
1p2
3sh
yx3 p2
1p2
3sh
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Hybrid Orbitals43
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Hybrid Orbital Shapes44
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Look Familiar?45
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2 3 4 5 6 sp sp2 sp3 sp3d sp3d2
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Other Types of Hybridization
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N atomic orbitals always produce N hybrid orbitals
Spectroscopic data suggests terminal atoms use hybrid orbitals as well
A terminal Cl uses sp3 hybridization in the arrangement of its lone pairs?
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Examples48
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π Bond49
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Ethylene50
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Ethylene51
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Characteristics of Multiple Bonds52
double bond = 1 σ-bond + 1 π-bond
triple bond = 1 σ-bond + 2 π-bonds
σ-bonds result from head-on overlap of orbitals
π-bonds results from side-by-side overlap
Atoms in a single bond can rotate freely, whereas atoms in a double bond are much less likely to:
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Characteristics of Multiple Bonds
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The shape of ethylene (C2H4)
Experimental evidence:
All six atoms lie in the same plane with 120º bond angles
Suggests trigonal planar structure sp2 hybridization
Each C: sp2 + p
σ-bonds with the sp2 orbitals
C –– C
4 C –– H
π-bond with the leftover p orbital
C –– C
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Characteristics of Multiple Bonds54
The shape of acetylene (C2H2) is linear sp hybridization
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Hybridization and Benzene55
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Molecular Orbital Theory56
Lewis Theory Connectivity, electron tracking
VSEPR Theory 3-D Structure around an atom
Valence Bond Theory Extended 3-D Structure Information Delocalization in Molecules (LIMITED) Illustrates Multiple Bonding Prediction of Reactivity
Molecular Orbital Theory Orbitals as a function of the whole molecule Delocalization much more thorough Anti-bonding and Non-bonding electrons Behavior in a magnetic field Spectral Data
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Paramagnetic / Diamagnetic
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O O
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Constructive and Destructive Interference
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Molecular Orbitals59
MO Theory – electrons occupy molecular orbitals spread over entire molecule
All valence electrons are delocalized Molecular orbitals are built by adding together atomic orbitals Linear combination of atomic orbitals (LCAO)
H – H bond in H2 with two 1s electrons:
A1sB1s
• LCAO-MO shows constructive interference
• ELCAO-MO < EAO
• LCAO-MO are bigger than AO
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Molecular Orbitals60
“Conservation of orbitals”
H2 bonds with two 1s orbitals – it must have 2 LCAO-MOs
A1sB1s
1st LCAO-MO is lower in energy than Atomic orbitals
This is called a bonding orbital
2nd LCAO-MO - destructive interference, higher PE than atomic orbitals
Nodal plane! – much less e- density “internuclearly”
This is an antibonding orbital
A1sB1s
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Sigma Bonds61
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Molecular Orbital Energy Diagram62
E
Shows:
Relative Energies of
Atomic Orbitals
Bonding MOs
Antibonding MOs
Shapes of LCAO-MOs
Type of MO – σ or π
Molecular Electron Configuration
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Electron Configurations of Diatomic Molecules
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Shows ALL valence electrons using the Building-Up Principle:
Electron configuration of H2 σ1s2
Based on MO theory – a single electron can hold a bond together
Use “Hund’s Rule just like in atomic orbitals!
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Bond Order64
Valence Bond Theory: Bond order was #
bonding pairs MO Theory, bond order
is: b = ½ (N – N*)
N = # e –s in bonding orbitals
N* = # e–s in antibonding orbitals
H2 b = 1 He2 b = 0
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2nd row Diatomic Molecules65
Must include both 2s and 2p orbitals in making our MOs
Only orbitals that are close in energy will form MOs
We have 2 (2s) orbitals and 6 (2p) orbitals 8 atomic orbitals
2s orbitals overlap in the same way as the 1s orbitals of H2
2p orbitals overlap in same orientations as in Valence bond theory:
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2nd row Diatomic Molecules66
We have 2 (2s) orbitals and 6 (2p) orbitals 8 atomic orbitals
2s orbitals overlap in the same way as the 1s orbitals of H2
2p orbitals overlap in same orientations as in Valence bond theory:
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Pi Bonds67
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2nd row Homonuclear Diatomic Molecules
68
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2nd row Homonuclear Diatomic Molecules
69
Homonuclear Diatomic MO Energy diagram for all elements right of (and including) oxygen
O2 F2 ‘Ne2’
And their combinations
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2nd row Homonuclear Diatomic Molecules
70
Homonuclear Diatomic MO Energy diagram for all elements left of oxygen
Li2 Be2 B2 C2 N2
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Relative MO Energy Levels71
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MO Diagram for Li272
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MO Diagram for Be273
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MO Diagram for N274
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MO Diagram for O275
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MO Diagram for F276
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77
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Relative AO Energy Levels78
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Heteronuclear Diatomic Molecules79
Account for ΔEN between AO’s
More electronegative atoms will have lower energy atomic orbitals
Contributions are dependent on energy level
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MO diagram for HF80
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Orbitals in Polyatomic Molecules81
An electron pair in a bonding orbital helps to bind together the whole molecule – can become very complex
Occasionally some MOs are neither bonding nor antibonding –these are called nonbonding orbitals
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MO Diagram for Carbon Monoxide
82
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Organic Molecules83
Chemists commonly mix VB and MO theories when discussing organic molecules:
VB theory is good for talking about σ bonds
MO theory is better for considering π bonds
Consider Benzene (C6H6):
VB perspective – each C is sp2 hybridized
σ framework is overlap of sp2 orbitals
MO perspective – conjugation of π bonds extremely important
π MOs account for delocalization of electrons
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π Orbitals of Benzene84
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Coordination Compounds85
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Colors of Coordination Compounds
86
Aqueous solutions of [Fe(SCN)(OH2)5]2+,
[Co(SCN)4(OH2)2]2-, [Cu(NH3)4(OH2)2]2+, and [CuBr4]2-
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Biological Importance87
Metalloproteins`
Hemoglobin
cobalamin (vitamin B12)
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Coordination Compounds88
Metals in the d-block form coordination compounds
Coordination compound – electrically neutral compound in which at least on of the ions present is a complex
Ligands – ions or molecules that are attached to the central metal atom or ion
Coordination number – number of ligands attached to the central metal
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Shapes of Complexes89
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Shapes of Complexes90
Sandwich compound
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Nomenclature91
[FeCl(OH2)5]+ pentaaquachloridoiron(II) Overall complex charge = +1 Ligand charge = Cl-1 and OH2 (water) Metal charge + ligand charge = overall complex charge
Fe + (-1) = +1 Fe = +2
[Co(NH3)3(OH2)3]2(SO4)3 triamminetriaquacobalt(III) sulfate Overall complex charge = +3 Ligand charge = 0 Co = +3
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Common Ligands92
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Common Ligands93
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Prefixes94
1 mono 2 di (bis) 3 tri (tris) 4 tetra (tetrakis) 5 penta
(pentakis) 6 hexa (hexakis) 7 hepta 8 octa
Prefixes in parentheses are used when the ligand already contains a Greek prefix or is polydentate
Example: ethylenediamine oxalato
(bidentate)
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More Nomenclature95
Na2[PtCl2(ox)2] sodium dichloridobis(oxalato)platinate(IV) Overall complex charge = -2 Ligand charge = Cl-1 and oxalato C2O4
2-
Metal charge + ligand charge = overall complex charge Pt + (-6) = -2 Pt = +4
[CoBr(NH3)5]SO4 pentaamminebromidocobalt(III) sulfate Overall complex charge = -2 Ligand charge = Br-1 and NH3 Metal charge + ligand charge = overall complex charge
Co + (-1) = +2 Co = +3
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Polydentate Ligands96
Some ligands attach to the metal more than once
Bidentate ligand
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Polydentate Ligands97
Some ligands attach to the metal more than once
Hexadentate ligand
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Coordination Isomers98
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Coordination Isomers99