Molecular Geometry and Bonding Theories

Molecular Geometry and Bonding Theories

H H CH4 H C H H molecular formula structural formula H

C 109.5o H H molecular shape H C H H H

tetrahedral shape of methane tetrahedron ball-and-stick model 109.5oo Tetrahedron Central Atom Central Atom

Substituents Substituents Methane, CH44 Tetrahedral geometry Methane, CH44 Copyright 2007 Pearson Benjamin Cummings. All rights reserved. Methane & Carbon Tetrachloride molecular formula structural formula

molecular shape H CH4 H C ball-and-stick model H H H H

C 109.5o H H Cl CCl4 Cl C Cl Cl

space-filling model Molecular Geometry 180o H C H H Trigonal planar Linear 109.5o 109.5o

H Tetrahedral 107.3 o Trigonal pyramidal 104.5o Bent H2O CH4 AsCl3 AsF5 BeH2 BF3 CO2

A Lone Pair Pear .. .. C N O 109.5o H H

H CH4, methane lone pair electrons 107o H H .. H H 104.5o

H NH3, ammonia H2O, water .. O O O O O3, ozone H O

O Molecular Shapes (Molecular Geometries) Two electron domains B A Three electron domains B B A Can only be linear Four electron domains B

Electronic geometry: tetrahedral A B B B B B Electronic geometry: trigonal planar

Molecular geometry could be: Tetrahedral Trigonal pyramidal Bent Molecular geometry could be: Trigonal planar (120o) Linear (180o) Bent Bonding and Shape of Molecules Number of Unshared Pairs

Covalent Structure Shape Examples -Be- Linear BeCl2 Trigonal planar BF3 2

0 3 0 4 0 C Tetrahedral CH4, SiCl4 3 1

: Number of Bonds Pyramidal NH3, PCl3 2 2 B : N

O: Bent H2O, H2S, SCl2 AB2 Linear AB3 Trigonal planar AB4 Tetrahedral AB3E Trigonal pyramidal

AB2E2 Angular or Bent AB2E Angular or Bent Valence Shell Electron Pair Repulsion Theory Planar triangular Tetrahedral

Trigonal bipyramidal Octahedral Valence Shell Electron Pair Repulsion Theory Planar triangular Tetrahedral Trigonal bipyramidal Octahedral

The VSEPR Model The Shapes of Some Simple ABn Molecules .. O C .. SO2 .. O N S

O O Linear O O Bent F S F O Trigonal

planar F Trigonal pyramidal SF6 F F F Cl F F T-shaped

F F F Square planar Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 305 F F P Xe F F

F S F F F F F Trigonal bipyramidal Octahedral

Molecular Shapes AB2 Linear AB3 Trigonal planar AB2E Angular or Bent AB5 Trigonal bipyramidal AB4 Tetrahedral AB4E Irregular tetrahedral

(see saw) AB6 Octahedral AB3E2 T-shaped AB5E Square pyramidal AB3E Trigonal pyramidal AB2E2 Angular or Bent

AB2E3 Linear AB4E2 Square planar Geometry of Covalent Molecules ABn, and ABnEm Type Formula AB2 AB2E AB2E2 AB2E3 AB3 AB3E AB3E2 AB4 AB4E

AB4E2 AB5 AB5E AB6 Shared Electron Pairs Unshared Electron Pairs 2 2 2 2 3 3

0 1 2 3 0 1 Linear Trigonal planar Tetrahedral Trigonal bipyramidal Trigonal planar Tetrahedral Linear Angular, or bent Angular, or bent Linear

Trigonal planar Triangular pyramidal 3 4 2 0 Triangular bipyramidal Tetrahedral T-shaped Tetrahedral 4 1

Triangular bipyramidal Ideal Geometry Observed Molecular Shape 4 5 2 0 Octahedral Triangular bipyramidal Irregular tetrahedral (or see-saw)

Square planar Triangular bipyramidal 5 6 1 0 Octahedral Octahedral Square pyramidal Octahedral Bailar, Moeller, Kleinberg, Guss, Castellion, Metz, Chemistry, 1984, page 317. Examples

CdBr2 SnCl2, PbI2 OH2, OF2, SCl2, TeI2 XeF2 BCl3, BF3, GaI3 NH3, NF3, PCl3, AsBr3 ClF3, BrF3 CH4, SiCl4, SnBr4, ZrI4 SF4, SeCl4, TeBr4 XeF4 PF5, PCl5(g), SbF5 ClF3, BrF3, IF5 SF6, SeF6, Te(OH)6, MoF6 Predicting the Geometry of Molecules Lewis electron-pair approach predicts number and types of bonds between the atoms in a substance and indicates which atoms have lone pairs of electrons but

gives no information about the actual arrangement of atoms in space Valence-shell electron-pair repulsion (VSEPR) model predicts the shapes of many molecules and polyatomic ions but provides no information about bond lengths or the presence of multiple bonds Introduction to Lewis Structures Lewis dot symbols 1. Used for predicting the number of bonds formed by most elements in their compounds 2. Consists of the chemical symbol for an element surrounded by dots that represent its valence electrons 3. A single electron is represented as a single dot Lewis Structures 1) Count up total number of valence electrons 2) Connect all atoms with single bonds

- multiple atoms usually on outside - single atoms usually in center; C always in center, H always on outside. 3) Complete octets on exterior atoms (not H, though) 4) Check - valence electrons math with Step 1 - all atoms (except H) have an octet; if not, try multiple bonds - any extra electrons? Put on central atom Molecules with Expanded Valence Shells Atoms that have expanded octets have AB5 (trigonal bipyramidal) or AB6 (octahedral) electron domain geometries. Trigonal bipyramidal structures have a plane containing three electron pairs. The fourth and fifth electron pairs are located above and below this plane. In this structure two trigonal pyramids share a base. F

F P F F For octahedral structures, there is a plane containing four electron pairs. F F F S F F F Similarly, the fifth and sixth electron pairs are located above and below this plane.

Two square pyramids share a base. F Trigonal Bipyramid F F P F F The three electron pairs in the plane are called equatorial. F The two electron pairs above and below this plane are called axial. The axial electron pairs are 180o apart and 90o from to the equatorial electrons.

The equatorial electron pairs are 120o apart. To minimize electron-electron repulsions, nonbonding pairs are always placed in equatorial positions, and bonding pairs in either axial or equatorial positions. F Octahedron F F S The four electron pairs in the plane are 90o to each other. F F F

The remaining two electron pairs are 180o apart and 90o from the electrons in the plane. Because of the symmetry of the system, each position is equivalent. The equatorial electron pairs are 120o apart. If we have five bonding pairs and one nonbonding pair, it doesnt matter where the nonbonding pair is placed. The molecular geometry is square pyramidal. If two nonbonding pairs are present, the repulsions are minimized by pointing them toward opposite sides of the octahedron. F F The molecular geometry is square planar. Xe F F

Electron-Domain Geometries Number of Electron Domains 2 Arrangement of Electron Domains B A B Electron-Domain Geometry Predicted Bond Angles

Linear 180o Trigonal planar 120o Tetrahedral 109.5o Trigonalbipyramidal 120o 90o

Octahedral 90o B A 3 B B B 4 A B

5 B B B 6 Ba B B A B B

B Be A Be Ba Be Acetic Acid, CH3COOH H H O C

C O 3 4 H H Number of electron domains Electron-domain geometry Predicted bond angles Hybridization of central atom Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 314

4 Tetrahedral Trigonal planar Tetrahedral 109.5o 120o 109.5o sp3 sp2 none

Intermolecular Forces + Ion-ion (ionic bonds) Ion-dipole Dipole-dipole + + +

H H Hydrogen bonding O H H London dispersion forces +

O H O H

London Dispersion Forces + + + London dispersion forces are created when on molecule with a temporarily dipole causes another to become temporarily polar. Molecular Polarity

Molecular Structure Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Electronegativity + H Cl 0

H 0 H Ionic vs. Covalent Ionic compounds form repeating units. Covalent compounds form distinct molecules. Consider adding to NaCl(s) vs. H2O(s): Na Cl Na Cl Na

Cl Na Cl Na Cl H Na Cl H O O

H H O H H NaCl: atoms of Cl and Na can add individually forming a compound with million of atoms. H2O: O and H cannot add individually, instead molecules of H2O form the basic unit. Holding it together Q: Consider a glass of water. Why do molecules of water stay together? A: There must be attractive forces.

Intramolecular forces are much stronger Intramolecular forces occur between atoms Intermolecular forces occur between molecules Intermolecular forces are not considered in ionic bonding because there are no molecules. The type of intramolecular bond determines the type of intermolecular force. Im not stealing, Im sharing unequally We described ionic bonds as stealing electrons In fact, all bonds share equally or unequally. Note how bonding electrons spend their time:

H2 H H 0 0 covalent (non-polar) HCl + H Cl

polar covalent LiCl [Li]+ [ + ionic Bonding electrons are shared in each compound, but are NOT always shared equally. The greek symbol indicates partial charge. Cl ] Dipole Moment Direction of the polar bond in a molecule. Arrow points toward the more

electronegative atom. + H Cl Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem - Dipole-induced dipole attraction

The attraction between a dipole and an induced dipole. Oxygen, O22 Nonpolar Oxygen, O22 Water, H22O + - Water, H22O +

- + - + + + - - + -

+ + - + + + - + + + -

+ + + - Dipole Dipole + induced induced dipole dipole + +

- + + + - + + + - +

+ + - + - + + - + + +

- - + - - + - + + Dipole Dipole

- + induced induced dipole dipole - + + Polar Polar Copyright 2007 Pearson Benjamin Cummings. All rights reserved.

Nonpolar Nonpolar Copyright 2007 Pearson Benjamin Cummings. All rights reserved. Determining Molecular Polarity Depends on: + H Cl dipole moments molecular shape

+ + + + Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Determining Molecular Polarity Nonpolar Molecules Dipole moments are symmetrical and cancel out.

F BF3 B F Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem F Determining Molecular Polarity Polar Molecules Dipole moments are asymmetrical and dont cancel . O H2O

H Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem H net dipole moment Determining Molecular Polarity Therefore, polar molecules have... asymmetrical shape (lone pairs) or asymmetrical atoms H CHCl3

Cl Cl Cl Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem net dipole moment Dipole Moment Bond dipoles C O

In H2O the bond dipoles are also equal in magnitude but do not exactly oppose each other. The molecule has a nonzero overall dipole moment. O .. .. Overall dipole moment = 0 O Nonpolar The overall dipole moment of a molecule is the sum of its bond dipoles. In CO 2 the bond dipoles are equal in magnitude but

exactly opposite each other. The overall dipole moment is zero. k q1 q2 F d2 m=Qr Bond dipoles H H Overall dipole moment Dipole moment, Coulombs law m

Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 315 Polar Polar Bonds .. Polar N H F F H B

F Polar F Cl F H Polar F .. O Cl

H .. H H F Nonpolar Polar Cl Cl F Cl

Xe F F Nonpolar H C C Cl Cl Nonpolar H H Polar

A molecule has a zero dipole moment because their dipoles cancel one another. HF HCl HBr HI How is the electron density distributed in these different molecules? Based on your comparison of the electron density distributions, which molecule should have the most polar bond, and which one the least polar? Arrange the molecules in increasing order of polarity. Mark Wirtz, Edward Ehrat, David L. Cedeno*

CH3Cl CH2Cl2 CHCl3 CCl4 Describe how is the electron density distributed in these different molecules? Based on your comparison of the electron density distributions, which molecule(s) should be the most polar, and which one(s) the least polar? Arrange the molecules in increasing order of polarity. Mark Wirtz, Edward Ehrat, David L. Cedeno* Benzene Mark Wirtz, Edward Ehrat, David L. Cedeno*

NO3- Nitrobenzene 2s Mark Wirtz, Edward Ehrat, David L. Cedeno* 2p (x, y, z) carbon How does H2 form? The nuclei repel But they are attracted to electrons They share the electrons +

+ Hydrogen Bond Formation Energy (KJ/mol) Potential Energy Diagram - Attraction vs. Repulsion 0 balanced attraction & repulsion no interaction increased attraction increased - 436 repulsion

0.74 A H H distance (internuclear distance) Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 318 Covalent bonds Nonmetals hold onto their valence electrons. They cant give away electrons to bond. Still want noble gas configuration. 1s22s22p63s23p6eight valence electrons (stable octet) Get it by sharing valence electrons with each other. By sharing both atoms get to count the electrons toward noble gas configuration. Covalent bonding Fluorine has seven valence electrons A second atom also has seven

By sharing electrons both end with full orbitals 8 Valence electrons F F 8 Valence electrons Single Covalent Bond A sharing of two valence electrons. Only nonmetals and Hydrogen. Different from an ionic bond because they actually form molecules. Two specific atoms are joined.

In an ionic solid you cant tell which atom the electrons moved from or to. Sigma bonding orbitals From s orbitals on separate atoms + + s orbital s orbital + + + + Sigma bonding molecular orbital Sigma bonding orbitals

From p orbitals on separate atoms p orbital p orbital Sigma bonding molecular orbital

Pi bonding orbitals P orbitals on separate atoms Pi bonding molecular orbital Sigma and pi bonds All single bonds are sigma bonds

A double bond is one sigma and one pi bond A triple bond is one sigma and two pi bonds. Atomic Orbitals and Bonding Bonds between atoms are formed by electron pairs in overlapping atomic orbitals Example: H2 (H-H) E 1s : 1s 1s Use 1s orbitals for bonding Example: H2O

From VSEPR: bent, 104.5 angle between H atoms Use two 2p orbitals for bonding? 1s 2p 90 1s 2p 2p E 2s How do we explain the structure predicted by VSEPR using atomic orbitals?

Overlapping Orbitals Draw orbital diagrams for F + F, H + O, Li + F 1s 2s 2p 2p F2 2s 1s 1s 1s

2s 2p 1s H2O electron transfer Li 1+ 1s 2s 2p

2s LiF is ionic (metal + non-metal) 1s F 1- lithium atom Li lithium ion Li+ ee- 3p

+ e- loss of one valence electron e- 3p+ e- e- fluorine atom F e-

e- e- e- e- e- e- fluoride ion F1- gain of one valence electron

e- 9p+ e e - e- 10p+ e- e- - e

- e- e - e e- - e- Formation of Cation lithium atom

Li lithium ion Li+ ee- 3p + e- loss of one valence electron e-

3p+ e- Formation of Anion fluorine atom F egain of one valence electron ee- fluoride ion F1- e- e-

e- e- e- e- 9p+ e e - 10p+ e- e-

- ee- e- e- e e- - e- Formation of Ionic Bond fluoride ion

F1- lithium ion Li+ ee- e- e- e- 3p+ e- e- 9p+

e- e- e e- - e- First, the formation of BeH2 using pure s and p orbitals. Be = 1s22s2 H BeH2 Be

s p atomic orbitals H No overlap = no bond! atomic orbitals The formation of BeH2 using hybridized orbitals. atomic orbitals H Be s

Be H p H hybrid orbitals H Be s p Be BeH2

sp p All hybridized bonds have equal strength and have orbitals with identical energies. Hybrid Orbitals Ground-state Be atom 1s 2s 2p Be atom with one electron promoted 1s 2s

2p Energy hybrid orbitals px py pz n=2 sp s 1s

sp 2p Be atom of BeH2 orbital diagram n=1 hybridize s orbital H p orbital two sp hybrid orbitals sp hybrid orbitals shown together (large lobes only)

Be H Hybrid Orbitals Ground-state B atom 2s 2p B atom with one electron promoted 2s 2p Energy hybrid orbitals

px py pz sp2 sp2 s 2p B atom of BH3 orbital diagram H hybridize B

s orbital H p orbitals three sps hybrid orbitals sp hybrid orbitals shown together (large lobes only) 2 H Hybridization the blending of orbitals Valence bond theory is based on two assumptions: 1. The strength of a covalent bond is proportional to the amount of overlap between atomic orbitals; the greater

the overlap, the more stable the bond. 2. An atom can use different combinations of atomic orbitals to maximize the overlap of orbitals used by bonded atoms. We have studied electron configuration notation and the sharing of electrons in the formation of covalent bonds. Lets look at a molecule of methane, CH4. Methane is a simple natural gas. Its molecule has a carbon atom at the center with four hydrogen atoms covalently bonded around it. Carbon ground state configuration What is the expected orbital notation of carbon

in itsYou ground should state? conclude that carbon only has TWO electrons available for bonding. That is not enough! 2p 2s 1s Can you see a problem with this? (Hint: How many unpaired electrons does this carbon atom have available for bonding?) How does carbon overcome this problem so that it may form four bonds? Carbons Empty Orbital The first thought that chemists had was that carbon promotes

one of its 2s electrons 2p 2s 1s to the empty 2p orbital. 2p 2p 2s 1s 2s 1s Non-hybridized orbital

hybridized orbital However, they quickly recognized a problem with such an arrangement 1s 1s 1s 1s 2p 2s 1s Three of the carbon-hydrogen bonds would involve

an electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom. But what about the fourth bond? A Problem Arises Unequal bond energy The fourth bond is between a 2s electron from the carbon and the lone 1s hydrogen electron. 1s 1s 1s 1s 2p

2s 1s Such a bond would have slightly less energy than the other bonds in a methane molecule. Unequal bond energy #2 This bond would be slightly different in character than the other three bonds in methane. This difference would be measurable to a chemist by determining the bond length and bond energy. But is this what they observe? The simple answer is, No. Measurements show that all four bonds in methane

are equal. Thus, we need a new explanation for the bonding in methane. Chemists have proposed an explanation they call it hybridization. Hybridization is the combining of two or more orbitals of nearly equal energy within the same atom into orbitals of equal energy. Enter Hybridization In the case of methane, they call the hybridization sp3, meaning that an s orbital is combined with three p orbitals to create four equal hybrid orbitals. These new orbitals have slightly MORE energy than the 2s orbital and slightly LESS energy than the 2p orbitals.

sp3 Hybrid Orbitals Carbon 1s22s22p2 Carbon could only make two bonds if no hybridization occurs. However, carbon can make four equivalent bonds. B A B B Energy hybrid orbitals

px py B pz s Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 321 sp3 sp3 C atom of CH4 orbital diagram Hybridization of s and p Orbitals The combination of an ns and an np orbital gives rise to two equivalent sp hybrids oriented

at 180. Combination of an ns and two or three np orbitals produces three equivalent sp2 hybrids or four equivalent sp3 hybrids. Copyright 2007 Pearson Benjamin Cummings. All rights reserved. Hybridization of s and p Orbitals Both promotion and hybridization require an input of energy; the overall process of forming a compound with hybrid orbitals will be energetically favorable only if the amount of energy released by the formation of covalent bonds is greater than the amount of energy used to form the hybrid orbitals. Copyright 2007 Pearson Benjamin Cummings. All rights reserved. Hybridization Involving d Orbitals

promote 3s 3p 3d unhybridized P atom P = [Ne]3s23p3 3s hybridize Ba Be F F

five sp3d orbitals A 3d Be F F 3d vacant d orbitals F P 3p

Be Ba Trigonal bipyramidal degenerate orbitals (all EQUAL) Pure atomic orbitals of central atom Hybridization of the central atom Number

of hybrid orbitals Shape of hybrid orbitals s,p sp 2 Linear s,p,p sp2 3

Trigonal Planar s,p,p,p sp3 4 Tetrahedral s,p,p,p,d sp3d 5 Trigonal Bipyramidal

s,p,p,p,d,d sp3d2 6 Octahedral Hybridization Animation, by Raymond Chang Hybridization Animation, by Raymond Chang Bonding Single bonds Overlap of bonding orbitals on bond axis Termed sigma or bonds Double bonds

Sharing of electrons between 2 p orbitals perpendicular to the bonding atoms Termed pi or bonds Bond Axis of bond 2p 2p One bond Multiple Bonds promote 2s hybridize 2p

2s 2p sp2 2p C2H4, ethene H H C C H

H one s bond and one p bond H H s H s C H s s

C H C s C H H Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 325-326 Two lobes of one p bond H

C C Multiple Bonds promote 2s hybridize 2p 2s 2p sp2

C2H4, ethene HH H 2p p sp HH p 2 C

sp2 sp 2 C sp2 H sp sp2 p p

one s bond and one p bond H H s H s C H s s C

H C s C H H Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 325-326 Two lobes of one p bond 2 H

HH p bond Internuclear axis p p s bonds H H C C

C H H C6H6 = benzene Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329 H C C H C

2p atomic orbitals Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329 s bonds H H C C C C H

Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329 H H C C H and p bonds s bonds H

H C C H H C C H C C

C H Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329 C H H s bonds H H C C H

H C C H C C C H Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329 C H H

N2O4 hn 2 NO2 nitrogen dioxide dinitrogen tetraoxide (free radical) O OO N NN O OO red-brown

colorless Energy-level diagram for (a) the H2 molecule and (b) the hypothetical He2 molecule (a) Energy s*1s 1s 1s H atom H atom s1s H2 molecule

(b) Energy s*1s 1s 1s He atom He atom s1s He2 molecule Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 332 Bond Order Bond order = (# or bonding electrons - # of antibonding electrons)

A bond order of 1 represents a single bond, A bond order of 2 represents a double bond, A bond order of 3 represents a triple bond. Because MO theory also treats molecules with an odd number of electrons, Bond orders of 1/2 , 3/2 , or 5/2 are possible. A bond order of 0 means no bond exists.

Energy-level diagram for the Li2 molecule s*2s Li = 1s 2s 2 1 2s1 Energy 2s1 s2s s*1s 1s2

1s2 Li Li Li 2 Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 334 s1s Energy-level diagram for molecular orbitals of second-row homonuclear diatomic molecules. s*2p p*2p 2p

2p p2p s2p s*2s 2s 2s s2s Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 337 Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 338 Increasing 2s 2p interaction Energy of p2p

molecular orbitals s2p s*2s s2s O2, F2, Ne2 Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 338 B2, C2, N2 Small 2s 2p interaction Large 2s 2p interaction B2 C2

N2 s*2p s*2p p*2p p*2p s2p p2p p2p s2p s*2s

s*2s s2s s2s O2 F2 Ne2 1 2 3

2 1 0 Bond enthalpy (kJ/mol) 290 620 941 495 155

----- Bond length (angstrom) 1.59 1.31 1.10 1.21 1.43 ----- Paramagnetic

Diamagnetic Bond order Magnetic behavior Paramagnetic Diamagnetic Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 339 Diamagnetic _____ s2s

p2px p2py s2p s*2s p*2px p*2py s*2p Arrange the atomic and molecular orbitals in order of increasing energy. How many orbitals are per molecule? Can you distinguish the bonding from the antibonding MOs? Mark Wirtz, Edward Ehrat, David L. Cedeno*

C2 Magnetic Properties of a Sample PARAMAGNETISM molecules with one or more unpaired electrons are attracted into a magnetic field. (appears to weigh MORE in a magnetic field) DIAMAGNETISM substances with no unpaired electrons are weakly repelled from a magnetic field. (appears to weigh LESS in a magnetic field) Experiment for determining the magnetic properties of a sample sample The sample is first weighed in

the absence of a magnetic field. N S When a field is applied, a diamagnetic sample tends to move out of the field and appears to have a lower mass. N S A paramagnetic sample is drawn into the field and thus appears to gain mass. Paramagnetism is a much stronger effect than is diamagnetism.

Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 339 Experiment for determining the magnetic properties of a sample sample The sample is first weighed in the absence of a magnetic field. N S When a field is applied, a diamagnetic sample tends to move out of the field and appears to have a lower mass. N

S A paramagnetic sample is drawn into the field and thus appears to gain mass. Paramagnetism is a much stronger effect than is diamagnetism. Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 339 Electron Domains lone Pair single bond double bond triple bond

: : Cl : : Cl : B Cl : : There are 3 electron domains about the central atom: no lone pairs and three single bonds. Three electron domains arrange themselves in a trigonal plane, with 120o angles. We predict a trigonal planar geometry. :

Cl : : B : : : Cl : : Cl : : :

Determine the shape of the BCl3 molecule: Electron-domain geometry: trigonal planar Molecular geometry (shape): trigonal planar sp2 hybrid orbitals shown together (large lobes only) One s orbital Hybridize Two p orbitals Three sp2 hybrid orbitals

Copyright 2007 Pearson Benjamin Cummings. All rights reserved. Ammonia, NH33 Ammonia, NH33 Triangular pyramidal Copyright 2007 Pearson Benjamin Cummings. All rights reserved.

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