2012 General Chemistry I

Transcript 2012 General Chemistry I

Chapter 3. MOLECULAR SHAPE AND STRUCTURE

THE VSEPR MODEL

3.1 The Basic VSEPR Model 3.2 Molecules with Lone Pairs on the Central Atom 3.3 Polar Molecules

VALENCE-BOND THEORY

3.4 Sigma and Pi Bonds 3.5 Electron Promotion and the Hybridization of Orbitals 3.6 Other Common Types of Hybridization 3.7 Characteristics of Multiple Bonds

THE VSEPR MODEL (Sections 3.1-3.3)

3.1 The Basic VSEPR Model

– Lewis structure: showing the linkages between atoms and the presence of lone pairs, but not the 3D arrangement of atoms

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Valence-shell electron-pair repulsion model (VSEPR model, first devized by Sidgwick and Powell, later modified by Gillespie and Nyholm) is based on Lewis structures. Molecular shapes are predicted by use of several rules that account for bond angles.

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Ideal Molecular Geometries According to VSEPR

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Generic “VSEPR formula”, AX

E

m - A = central atom; X n = n atoms bonded to central atom E m = m lone pairs on central atom E.g. BF 3 (AX 3 ), SO 2 (AX 2 E), SO 3 2 (AX 3 E), CH 4 (AX 4 ), PCl 5 (AX 5 )

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Rule 1. Regions of high electron concentration (bonds and lone pairs on the central atom) repel one another and, to minimize their repulsions, these regions move as far apart as possible while maintaining the same distance from the central atom.

This gives rise to the basic geometries, which are related to the total number (X n + E n ) of electron pairs around the central atom (next slide).

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Rule 2. There is no distinction between single and multiple bonds: a multiple bond is treated as a single region of high electron concentration.

E.g.

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Self-Test 3.1B

Predict the shape of a formaldehyde molecule, CH 2 O.

Solution Although written, CH 2 O, carbon is the central atom in a H 2 CO skeleton, hence Lewis structure is O H C H C is bonded to 3 atoms, so the molecule is trigonal planar.

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3.2 Molecules with Lone Pairs on the Central Atom

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Rule 3. All regions of high electron density, lone pairs and bonds, are included in a description of the electronic arrangement, but only the positions of atoms are considered when identifying the

shape of a molecule.

E.g.

Trigonal pyramidal 2012 General Chemistry I Angular (bent or V-shape)

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Rule 4 The strength of repulsions are in the order lone pair-lone pair > lone pair-atom > atom-atom.

- AX 3 E type trigonal pyramidal - AX 4 E type more stable seesaw shaped axial equatorial - AX 3 E 2 type 2012 General Chemistry I T-shaped - AX 4 E 2 type square planar

Self-Test 3.3B

(a) Give the VSEPR formula of the ClO

2 -

ion. Predict (a) its electron arrangement and (c) its shape.

Solution The Lewis structure is _ O ..

O , showing two 2 atoms and 2 lone pairs on the central Cl atom.

(a) Hence its VSEPR formula is AX 2 E 2 .

(b) The electron arrangement is tetrahedral.

(c) Its actual shape is angular (bent or V-shape).

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Predicting the molecular shape of SF 4 Step 1 Draw the Lewis structure.

Step 2 Assign the electron arrangement around the central atom.

Step 3 Identify the molecular shape. AX 4 E.

Step 4 Allow for distortions.

2012 General Chemistry I (Bent) seesaw shape

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3.3 Polar Molecules

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Polar molecule: a molecule with a nonzero dipole moment. A diatomic molecule is polar if its bond is polar (if the electronegativity difference between the two atoms > 0.3) - see Chapter 2.12.

E.g. HCl is a polar molecule (dipole moment of 1.1 D).

H 2 is a nonpolar molecule (zero dipole moment).

A polyatomic molecule is polar if it has polar bonds arranged in space in such a way that the dipole moments associated with the bonds do not cancel.

See Table 3.1 for dipole moments of selected molecules.

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nonpolar polar 2012 General Chemistry I polar

nonpolar

Molecular geometries giving non-polar or polar molecules: AX 2 – AX 3 E 2 (Fig. 3.7) 2012 General Chemistry I

Molecular geometries giving non-polar or polar molecules: AX 4 – AX 6 (Fig. 3.7) 2012 General Chemistry I

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VALENCE-BOND (VB)THEORY (Sections 3.4-3.7)

The VB theory is based on the Lewis model: each bonding electron pair is localized between two bonded atoms.

It is a localized electron model, devized originally by Heitler and London, and later modified by Pauling, Slater and others.

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3.4 Sigma (

) and Pi (

) Bonds

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The

-bond is formed by ‘head-on’ overlap: there is no nodal surface containing the interatomic (bond) axis.

It is cylindrical or “sausage” shaped.

E.g. H 2 ,

(1s, 1s) HF,

(1s, 2p z ) See next slide N 2 ,

(2p z , 2p z ) 2012 General Chemistry I

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The

-bond is formed by ‘sideways’ overlap: there is a nodal plane containing the interatomic (bond) axis.

It is composed of two cylindrical shapes (lobes), one above and the other below the nodal plane.

N 2 one

-bond with two perpendicular

-bonds - In General, single bond (one

-bond), double bond (one

- and one

-bond) triple bond (one

- and two

-bonds) 2012 General Chemistry I

Self-Test 3.6A

How many

-bonds and how many in (a) CO 2 and (b) CO?

-bonds are there Solution

p s

H C

,with two

-bonds (a) CO 2 and two is

-bonds.

(b) CO is O C

with one

-bond and two

-bonds.

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3.5 Electron Promotion and the Hybridization of Orbitals

all C-H bonds equivalent

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Electron promotion: electron relocated to a higher-energy orbital promotion 2012 General Chemistry I hybridization

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Hybrid orbitals: produced by hybridizing orbitals of a central atom C [He]2s 1 2p x 1 2p y 1 2p z 1 4 sp 3 hybrids h 1 h 3 = s + p x = s - p x + p + p y y + p - p z z h h 2 4 = s - p x = s + p x - p y - p y + p z - p z each C-H bond sp 3 hybrid orbital 2012 General Chemistry I

Examples of sp 3 hybrid orbitals in bonding - Ethane, C 2 H 6 Each C atom uses sp 3 hybrid orbitals.

The C-C bond is formed as

(C2sp 3 , C2sp 3 ).

Each C-H bond is formed as

(C2sp 3 , H1s).

Ammonia, NH 3 Four sp 3 hybrid orbitals; one occupied by a lone pair, and the other three forming N-H

-bonds 2012 General Chemistry I

3.6 Other Common Types of Hybridization

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sp 2 hybrid orbitals in BF 3

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sp hybrid orbitals in CO 2 2012 General Chemistry I

sp 3 d hybrid orbitals in PCl 5 - sp 3 d 2 hybrid orbitals in SF XeF 4 6 and 2012 General Chemistry I

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Self-Test 3.8A

Describe (a) the electron arrangement, (b) the molecular shape, and (c) the hybridization of the central atom in chlorine trifluoride.

Solution ..

Cl ..

(a) The Lewis structure of ClF 3 is F F The VSEPR formula of ClF 3 is AX 3 E 2 arrangement is trigonal bipyramidal.

F , hence the electron (b) The molecular geometry is T-shaped.

(c) The hybridization on Cl is sp 3 d.

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3.7 Characteristics of Multiple Bonds

Ethene, CH 2 =CH 2 C-C

bond,

(C2sp 2 , C2sp 2 ) C-C

bond,

(C2p, C2p) each C-H bond formed as

(C2sp 2 , H1s) 2012 General Chemistry I

- Benzene, C 6 H 6 C-C

C-C

bond, bond,

s p

(C2sp 2 , C2sp (C2p, C2p) 2 ) Each C-H bond formed as

(C2sp 2 , H1s) - Ethyne (acetylene), C 2 H 2 C-C

bond,

(C2sp, C2sp) Two C-C

bond,

(C2p x , C2p x ),

(C2p y , C2p y ) Each C-H bond formed as

(C2sp, H1s) 2012 General Chemistry I

Self-Test 3.9A

Describe the structure of the carbon suboxide molecule, C 3 O 2 , in terms of hybrid orbitals, bond angles, and

-bonds. The atoms lie in the order OCCCO.

- and Solution ..

C C C ..

O : The VSEPR model predicts all three central C atoms to be of the type AX 2 , and hence bonding to all three is linear.

All three C atoms are sp hybridized and oxygen is sp 2 hybridized.

All bond angles are 180 o .

Each carbon forms one neighbor.

p s

- and one

p p

-bond, to each C or O O

p s

s p

O 2012 General Chemistry I

Chapter 3. MOLECULAR SHAPE AND STRUCTURE

MOLECULAR ORBITAL THEORY

3.8 The Limitations of Lewis’s Theory 3.9 Molecular Orbitals 3.10 Electron Configurations of Diatomic Molecules 3.11 Bonding in Heteronuclear Diatomic Molecules 3.12 Orbitals in Polyatomic Molecules

MOLECULAR ORBITAL THEORY (Sections 3.8-3.12)

3.8 The Limitations of Lewis’s Theory

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Valence Bond (VB) theory deficiencies - Cannot explain paramagnetism of O

(existence of unpaired electrons) - Difficulty treating electron-deficient compounds such as B 2 H 6 - No simple explanation for spectroscopic properties of compounds

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Molecular Orbital (MO) theory advantages Addresses all of the above shortcomings of VB theory - Provides a deeper understanding of electron-pair bonds - Accounts for the structure and properties of metals and semiconductors - Universally used in calculations of molecular structures 2012 General Chemistry I

3.9 Molecular Orbitals

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Molecular orbitals (MOs) contain the valence electrons in molecules: they are delocalized over the whole molecule.

- In VB theory, bonding electrons are localized on atoms or between pairs of atoms.

- n molecular orbitals must be constructed from n atomic orbitals:

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MOs are formed by linear combination of atomic orbitals (LCAO-MO): - Bonding orbital (constructive interference)

A1s +

B1s → overall lowering of energy - Antibonding orbital (destructive interference)

A1s -

B1s → overall raising of energy 2012 General Chemistry I

Molecular orbital energy-level diagram MO energy-level diagrams are an integral part of MO theory: they give the following information.

1. Relative energies of original AOs and resulting MOs 2. Orbital location of electrons and electron spin (using arrows)

H

2 - In H 2 , two 1s-orbitals merge to form the bonding orbital

1s and the antibonding orbital

1s * 2012 General Chemistry I

3.10 Electron Configurations of Diatomic Molecules

Constructing MOs (by combining AOs) must follow a number of rules,the first three of which are similar to those of the ‘Building-up Principle’ for constructing atomic electron configuration: 1. Electrons are accommodated in the lowest-energy MO, then in orbitals of increasingly higher energy.

2. According to the Pauli exclusion principle, each MO can accommodate up to two electrons. If two electrons are present

in one orbital, they must be paired.

3. If more than one MO of the same energy is available, the electrons enter them singly and adopt parallel spins ( Hund’s rule).

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4. Only atomic orbitals of the same symmetry along the bond axis can be combined.

s + p (head-on) p + p (head-on) + (sideways) p p Not allowed: + or + 5. Atomic orbitals of the same or similar energy interact more strongly than those with very different energies. Low energy AOs contribute more to bonding MOs; high energy AOs contribute more to antibonding MOs.

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Homonuclear diatomic molecules of Period 2 elements (Li

– F

)

General Features

Two 2s atomic orbitals (one on each atom) overlap to form two

orbitals: one bonding (

2s -orbital) and the other antibonding (

2s *-orbital).

Six 2p atomic orbitals (three on each atom) overlap to form six MOs: two 2p z bonding and antibonding (

2p ,

2p *) MOs, and four 2p x , 2p y orbitals to form two

2p orbitals to form

and two

2p *MOs.

See next slide.

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Four 2p x , 2p y AOs form two

2p two

2p * MOs and Two 2p z AOs form bonding

2p and antibonding

2p * MOs Two 2s AOs form one bonding (

2s ) MO and one antibonding (

2s * ) MO 2012 General Chemistry I

For O

and F

2 41

- For O 2 and F 2 , the energy levels of 2s and 2p are well-separated.

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from Shriver

- From Li 2 to N 2 , the energy levels of 2s and 2p are close, and thus the 2s orbital also participates in forming

2p orbitals.

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from Shriver

MO theory account of bonding in N

2 - In N 2 , each atom supplies five valence electrons.

A total of ten electrons fill the MOs.

The ground configuration is, b = ½(8-2) = 3

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Bond order (b): net number of bonds Bond order, b = 1 2 (N e _ N e *) N e = number of electrons in bonding MOs N e * = number of electrons in antibonding MOs 2012 General Chemistry I

LUMO HOMO

MO theory account of bonding in O

2 In O 2 , each atom supplies six valence electrons.

A total of twelve electrons fill the MOs.

The ground configuration is, b = ½(8-4) = 2 LUMO HOMO accounts for paramagnetism of O 2 2012 General Chemistry I

Self-Test 3.10A

Deduce the electron configuration and bond order of the ion C 2 2 .

Solution The ion C 2 2 has the configuration of C 2 , but with two extra electrons, which enter the

2p bonding MO.

Configuration is

2s 2

2s * 2

2p 4

2p 2 Bond order is 1/2(8 _ 2) = 3 E

2p *

2s *

2s 2012 General Chemistry I

3.11 Bonding in Heteronuclear Diatomic Molecules

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A diatomic molecule built from atoms of two different elements is polar, with the electrons shared unequally by the two atoms.

In a nonpolar covalent bond, c A 2 = c B 2 - In an ionic bond, the coefficient belonging to one ion is zero.

- In a polar covalent bond, the AO belonging to the more electronegative atom has the lower energy, and so it makes the larger contribution to the lowest energy (bonding) MO.

- Conversely, the contribution to the highest energy (most antibonding) orbital is greater for the higher-energy AO, which belongs to the less electronegative atom.

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- For HF, electronegativity of F (4.0), H (2.2) mainly H1s-orbital → partial (+) charge on H - For EO (E = C or N), 2012 General Chemistry I mainly F2p z -orbital → partial (–) charge on F

Self-Test 3.11B

Write the configuration of the ground state of the cyanide ion, CN , assuming that its molecular orbital energy diagram is the same as that for CO.

Solution There are 4 + 5 + 1 = 10 electrons to accommodate into the MOs of Fig. 3.35.

Hence the configuration is 1

2 2

* 2 1

4 3

2 (

2s 2

2s * 2

2p 4

2p 2 ) E

2p *

2s *

2s 2012 General Chemistry I

3.12 Orbitals in Polyatomic Mol ecules

The MOs spread over all atoms in the molecule.

Experimentally studied by using ultraviolet and visible spectroscopy.

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Water, H

O

- A water molecule with six atomic orbitals (one O2s, three O2p, and two H1s) O2p x antibonding orbitals nonbonding orbitals H1s-O2p y -H1s H1s-(O2s,2p z )-H1s 2012 General Chemistry I bonding orbitals

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Benzene, C

H

6 - All thirty C2s-, C2p-, and H1s-orbitals contribute to MOs.

- The orbitals in the ring plane: C2s-, C2p x , C2p y , and six H1s -orbitals → delocalized

-orbitals for C-C and C-H - six C2p z -orbitals perpendicular to the ring → delocalized

-orbitals spreading the ring From VB, each C atom with sp 2 orbitals hybrid forming

-bonds and 120 ° angles.

From MO, the six C2p z -orbitals delocalized

-orbitals.

form six 2012 General Chemistry I

Great stability: the

-electrons occupy only orbitals with a net bonding effect.

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Hypervalent compounds (a central atom forms more bonds than allowed by the octet rule) From VB, SF 6 has S with sp 3 d 2 hybridization From MO, four orbitals of S and six of F, a total of 10 AOs → 10 MOs 12 electrons occupy bonding and nonbonding Orbitals.

Average bond order of each S-F is 2/3.

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Colors of vegetation highest occupied molecular orbital (HOMO) lowest unoccupied molecular orbital (LUMO) → excited an electron from a HOMO to a LUMO, by the photons with the energy of visible light Beta-carotene and lycopene contain many conjugated C=C bonds.

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ULTRAVIOLET AND VISIBLE SPECTROSCOPY

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The Technique - The electrons in the molecule can be excited to a higher energy state, by electromagnetic radiation .

Bohr frequency condition,

E = h

- UV-vis absorption gives us information about the electronic energy levels of molecules .

i.e. Chlorophyll absorbs red and blue light, leaving the green light present in white light to be reflected.

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Chromophores are characteristic groups of atoms in molecules that absorb certain wavelengths of uv or visible light –

* transition in nonconjugated double bonds ~ 160 nm, but for molecules with many conjugated double bonds it is the visible region.

– n-

* transition in the carbonyl group ~ 280 nm LUMO HOMO - d-to-d transition in d-metal complexes in visible ranges - charge transfer transition in d-metal complexes electrons migrate from the ligands to the metal atom or vice versa.

E.g. deep purple color of MnO 4 2012 General Chemistry I