Presentation 4 Inorganic-aj

Chapter 2 Molecular Structure and Bonding Lecture 4 09/10/2009 Dr. Andrei Jitianu

Atkins • Overton • Rourke • Weller • Armstrong

Shriver & Atkins

Inorganic Chemistry Fourth Edition

Chapter 2 Molecular Structure and Bonding

Copyright © 2006 by D. F. Shriver, P. W. Atkins, T. L. Overton, J. P. Rourke, M. T. Weller, and F. A. Armstrong

Outline • Lewis structures – The octet rule; – The structure and bond properties; – The VSEPR model;

• Valence-bond theory – The hydrogen molecule; – Homonuclear diatomic molecules; – Polyatomic molecules;

• Molecular orbital theory – – – –

An introduction in theory Homonuclear diatomic molecules; Heteronuclear diatomic molecules; Bond properties;

Outline • Lewis structures – The octet rule; – The structure and bond properties; – The VSEPR model;

• Valence-bond theory – The hydrogen molecule; – Homonuclear diatomic molecules; – Polyatomic molecules;

• Molecular orbital theory – – – –

An introduction in theory Homonuclear diatomic molecules; Heteronuclear diatomic molecules; Bond properties;

Lewis structures • A covalent bond is formed when two neighbouring atoms share an electron pair. – Single bond – one shared electron pair A:B or A-B – Double bond – two shared electron pairs – A::B or A=B – Triple bond – three shared electron pairs – A:::B or A≡B

An unshared of valence electrons on an atom (:A) lone pair

The octet rule

Key point: • Atoms share electron pairs until they have acquired an octet of valence electrons. Each atom shares electrons with neighbouring atoms to achieve a total of eight valence electrons (an ‘octet’)

The steps of constructing a Lewis structure: 1. Decide on number of electrons that are to be included in the structure by adding together the numbers of all valence electrons provided by the atoms. 2. Write the chemical symbols of the atoms in the arrangement that shows which atoms are bonded together. 3. Distribute the electrons in pairs so that there is one pair of electrons forming a single bond between each pair of atoms bonded together , and then supply electron pairs (to form lone pairs or multiple bonds) until each atom has an octet

Exercises Write the structure of the BF4-

Tetrahedral

Write a Lewis structure for PF3 molecule

Trigonal pyramidal

Resonance

Key points: • Resonance between Lewis structures lowers the calculated energy of the molecule and distributes the bonding character of electrons over the molecules; • Lewis structures with similar energies provide the greatest resonance stabilization.

This Lewis structure indicates incorrectly that one O-O bond is different from the other. In fact O-O bond from O3 is 128 pm between O-O which is 148 pm and O=O which is 121pm. The concept resonance shows that the actual structure is a superposition, or average of all feasible Lewis structures corresponding to a given atomic arrangement. Resonance

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• In quantum mechanical terms: – The electron distribution of each structure is presented by a wavefunction – The actual wavefunction is the superposition of the individual wavefunctions for each contributing structure.

ψ = ψ(O-O=O)+ψ(O=O-O) - The overall wavefunction is written as a superposition with equal contributions from both structures because the two structures have identical energies

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The blended structure of two or more Lewis structures is called a resonance hybrid. – –

The resonance occurs between structures that differ only in allocation of electrons; The resonance does not occur between structures in which the atoms themselves lie in different positions.

• Resonance effects: 1.Resonance averages the bond characteristics over the molecules 2.The energy of resonance hybrid structure is lower than that of any single contributing structure.

The energy of O3 resonance hybrid is lower than that of either individual structure alone.

Resonance is important when several structures with the same energy are present. For this case all the structures with the same energy contribute equally to the overall structure. Structure with different energies may also contribute to an overall resonance hybrid.

Higher energy

Formal Charge

Key points: • The formal charge is the charge an atom would have if electron pairs were shared equally; • Lewis structures with low formal charges typically have the lowest energy.

The decision about the best arrangement of electrons to give a Lewis structure with the lowest energy can be estimated quantitatively by formal charge, f , on each atom.

Formal charge

f = V – L – (1/2)P

V- number of valence electrons L- number of lone pair of electrons P- number of shared electrons

The formal charge is the difference between the number of valence electrons in the free atom and the number that the atom has in the molecule, assuming it owns one electron of each shared pair and both electrons of any lone pair.

A representation of the calculation of formal charge. The lines show the bonding electrons and lone pair electrons are apportioned to each atom (a.) In a diatomic molecule A-B and (b.) in a triatomic molecule A=B-C. The formal charge on each atom is the difference between the number of electrons obtained in this way and the number in the free, neutral atom.

Oxidation State

Key points: • Oxidation numbers are assigned by applying a set of rules;

• The oxidation number ω (omega) is a parameter obtained by exaggerating the ionic character of a bond. • Oxidation state is the physical state of the element corresponding to its oxidation number

Hypervalence

Key point: • Hypervalence and octet expansion occur for elements following Period 2. • The elements of Period 2 from Li to Ne obey the octet rule, but the elements of later periods show deviations from it.

10 e- in the valence shell

12 e- in the valence shell

A P atom can accommodate more than eight electrons if it uses its vacant 3d orbitals; In PCl5 with its five pairs of bonding electrons, at least one of 3d orbital must be used in addition to the four 3s and 3p orbitals of the valence shell.

Outline • Lewis structures – The octet rule; – The structure and bond properties; – The VSEPR model;

• Valence-bond theory – The hydrogen molecule; – Homonuclear diatomic molecules; – Polyatomic molecules;

• Molecular orbital theory – An introduction in theory – Homonuclear diatomic molecules; – Heteronuclear diatomic molecules;

Structure and bond properties a. Bond length Key points: • The equilibrium bond length in a molecule is the separation of the centers of the two bonded atoms; • Covalent radii vary through the periodic table in much the same way as metallic and ionic radii.

The equilibrium bond length in a molecule is the separation of the centers of the two bonded atoms. »X-Ray data on solids; »IR and Microwave spectroscopy;

• The contribution of an atom to a covalent bond is called the covalent radius of the element.

Bond length

Covalent radii

A covalent radius expresses the closeness of approach of bonded atoms.

The closeness of approach of nonbonded atoms in neighbouring molecules that are in contact is expressed in terms of the van der Waals radius of an element; - the internuclear separation when the valence shells of the two atoms are in nonbonding contact

b. Bond strength

Key points: • The strength of a bond is measured by its dissociation enthalpy; • Mean bond enthalpies are used to make estimates of reaction enthalpies

Thermodynamic measure of the strength of an AB bond - bond dissociation enthalpy ∆Hϴ(A-B)

AB(g) →A(g)+B(g) standard reaction enthalpy The mean bond enthalpy, B, is the average bond dissociation enthalpy taken over a series of A-B bonds in different molecules.

Mean bond enthalpy, can be used to estimate reaction enthalpies.

d. Electronegativity and bond enthalpy Key point: • The Pauling scale of electronegativity is useful for estimating bond enthalpies and for assessing the polarities of bonds.

Pauling definition: power of an atom of the element to attract electrons to itself when it is a part of a compound. The original concept related the electronegativity with the energetics of bond formation

A2(g)+B2(g) →2AB(g) ∆E-excess of energy of the A-B bond over the average energy of A-A and B-B bonds can be attributed to the presence of ionic contributions to the covalent bonding.

|χP(A)-χP(B)|=0.102(∆E/kJ mol-1)1/2 Where

∆E=B(A-B)-1/2{B(A-A)+B(B-B)}

B(A-B) is the mean A-B bond enthalpy. •

Binary compounds with ∆χ>1.7 - ionic

CsF- ∆χ=3.19; χmean=2.38 –ionic F2 - ∆χ=0; χmean=3.98 Cs - ∆χ=0; χmean=0.79 Ex: χMg=1.31; χO=3.44

A Ketelaar triangle, showing how a plot of average electronegativity against electronegativity difference can be used to classify the bond type for binary compounds