Contents. 1. Basic Concepts. 2. The Covalent Bond. 3. The Valence-Shell Electron-Pair Repulsion Models 4. Bond theories. 5. The Metallic Bond.

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1 Chemical Bonding (II) Topic 4. Chemical Bonding (II) (II) 1

2 Contents 1. Basic Concepts. a) Molecular parameters b) Lewis Dot Symbols 2. The Covalent Bond a) Polar Covalent Bond b) Formal Charge c) Exceptions to the octet rule 3. The Valence-Shell Electron-Pair Repulsion Models 4. Bond theories a) Valence Bond Theory. Hybridization of Atomic Orbitals b) Molecular Orbital Theory 5. The Metallic Bond. Topic 4. Chemical Bonding (II) 2

3 1. BASIC CONCEPTS A) Molecular Parameters Bond Distance -Average distance between the nuclei of two bonded atoms. -Factors affecting bond length: a) Size of the bonded atoms (HI > HBr > HCl > HF) b) Multiplicity of bond (C-C > C=C ) Bond Angle Image from: molecule_dimensions.svg -The internal angle between the orbitals containing electron pairs in the valence shell of the central atom in a molecule. Bond Enthalpy It is the energy required to break one mole of bonds of a particular type of substance in gaseous state. Topic 4. Chemical Bonding (II) 3

4 1. BASIC CONCEPTS B) Lewis dot symbols Valence e - play a fundamental role in chemical bonding the octet rule Chemical bond in terms of Lewis dot symbols ionic bond e - transfer covalent bonds e- sharing C N O F Ne P As Topic 4. Chemical Bonding (II) 4

5 2. THE COVALENT BOND In covalent bond atoms share electrons to complete their valence shell (i.e. to reach the gas noble configuration) It forms between no-metalic elements. F F F F Bond Pair F F Lone Pairs Coordinate covalent Bonds The electron pair is provided by one of the atoms that form the covalent bond. Topic 4. Chemical Bonding (II) 5

6 Polar Covalent Bonds Electronegativity: describes the ability of an atom (or, more rarely, a functional group) to attract electrons (or electron density) towards itself in a covalent bond. Dipole moment µ = Q x r r = distance between atoms Q = charge Units Debye, D 1 D = C m F a) Non polar molecules b) Polar molecules H H H F F Centre of positive charge Centre of negative charge Atomic nucleus There is a shift of the electron density from H to F Topic 4. Chemical Bonding (II) 6

7 Polarity of molecules Diatomic molecules Homonuclear molecules: H 2, O 2 y F 2 Heteronuclear molecules HCl, CO y NO Polyatomic Molecules Polarity depends on: -Polarity of the chemical bonds -Geometry CCl 4 A non-polar molecule CHCl 3 A polar molecule µ is the vectorial sum of the dipolar moments of the chemical bonds of the molecule. Topic 4. Chemical Bonding (II) 7

8 Formal Charge formal charge on an atom in a Lewis structure = total number of valence electrons in - the free atom total number of nonbonding electrons total number of bonding electrons ( ) The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion. Practice Exercise. Write the Lewis structure for formaldehyde (CH 2 O) using the concept of formal charge. Then compare your prediction with the real structure. Topic 4. Chemical Bonding (II) 8

9 Exceptions to the Octet Rule Odd e - species. N=O Incomplete octets. Expanded octets. Topic 4. Chemical Bonding (II) 9

10 A:: 3. VSEPR MODEL No. of electron pairs or charge clouds Distribution of the charge clouds (Molecular Shape) Electron pairs repel each other whether they are in chemical bonds (bond pairs) or unshared (lone pairs). Electron pairs assume orientations about an atom to minimize repulsions. A) Molecules with no lone pairs on the central atom BeCl 2 BF 3 CH 4 : A : A : A Linear Trigonal planar Tetrahedral 5 PCl 5 : : A : Trigonal bipyramidal Topic 4. Chemical Bonding (II) 10 6 SF 6 : : : A : : : Octahedral

11 3. VSEPR MODEL B) Molecules with lone pairs on the central atom bond pairs vs. Bond pairs lone pairs vs. bond pairs Lone pairs vs. Lone pairs Topic 4. Chemical Bonding (II) 11

12 3. VSEPR MODEL B) Molecules with lone pairs on the central atom Number of bond pairs ( x ) Number of Lone pairs of electrons( y ) Number of group of electrons (n = x+y) VSEPR Notation (AB x E y ) Electron group geometry Example Ideal bond angles Molecular geometry AB 2 Linear BeCl 2 180º Linear AB 2 E 1 Trigonal planar SO 2 120º Angular 4 0 AB 4 CH 4 Tetrahedal AB 3 E 1 Tetrahedral NH º Piramidal 2 2 AB 2 E 2 H 2 O Angular 5 0 AB 5 PCl 5 Trigonal pyramidal 4 1 AB Trigonal 5 4 E 1 SF 4 Seesaw 90º; 120º 3 2 AB Bipyramidal 3 E 2 ClF 3 T- shaped 2 3 AB 2 E 3 XeF 2 Linear 6 0 ICl AB 5 E 5, BrF 5 1 Octahedral XeOF 4 AB 6 SF 6 90º Octahedral Square pyramidal 4 2 AB 4 E 2 XeF 4; ; ICl 4 - Square planar Topic 4. Chemical Bonding (II) 12

13 4. BOND THEORIES Valence Bond Theory Atomic orbital overlap describes covalent bonding. Area of overlap of orbitals is in phase. A localized model of bonding. H 2 S : Principles and Moderns Applications (R.H. Petrucci) Topic 4. Chemical Bonding (II) 13

14 4. BOND THEORIES Valence Bond Theory Hybridization of Atomic Orbitals CH 4 Hybridization: term applied to the mixing of 4 σ bonds: C(sp 3 )-H(1s) atomic orbitals in a atom (i.e. linear combination of the wavefunctions). Bond angle: º -Promotion -Hybridization CH 4 molecule Topic 4. Chemical Bonding (II) 14

15 4. BOND THEORIES Valence Bond Theory BF 3 Orbital diagram B (ground state) Promotion of one electron (2s 2p) (Excitation) Mixing of orbitals (2s & 2p) (Hybridization) BF 3 molecule 3sp 2 orbitals + Unhybridised 2p 3 σ bonds: B(sp 2 )-F(2p) Bond angle: 120º Topic 4. Chemical Bonding (II) 15

16 4. BOND THEORIES Valence Bond Theory BeCl 2 Orbital diagram Be (ground state) Promotion of one electron (2s 2p) (Excitation) Mixing of orbitals (2s & 2p) (Hybridization) BeCl 2 molecule 2sp orbitals + 2 Unhybridised 2p 2 σ bonds: Be(sp)-Cl(3p) Bond angle: 180º Topic 4. Chemical Bonding (II) 16

Mixing of orbitals (3s, 3p & 3d) (Hybridization) 3 3 sp

17 4. BOND THEORIES SF 6 Orbital diagram S (ground state) 3 3 Promotion of 2 electron (3s 3p, 3d) (Excitation) Mixing of orbitals (3s, 3p & 3d) (Hybridization) 3 3 sp 3 d 2 orbitals SF 6 molecule Topic 4. Chemical Bonding (II) 17

18 4. BOND THEORIES Valence Bond Theory Hybridization of Atomic Orbitals. Geometry of Organic Molecules. B) Geometry in acethylene molecule C 2 H 2 H C C H σbond C(sp)-C(sp) σbond C(sp)-H(1s) π Bonds C(2p)-C(2p) Topic 4. Chemical Bonding (II) 18

19 Hybridization Ideal bond angles Molecular geometry Stick and balls Model sp 180º Linear sp 2 120º Trigonal sp º Tetrahedral sp 3 d 90º; 120º Trigonal pyramidal sp 3 d 2 90º Octahedral Topic 4. Chemical Bonding (II) 19

20 4. BOND THEORIES Molecular Orbital Theory Explains the formation of the H-H bond in H 2 molecule. Description of the model: When the 1s wave functions of the two hydrogen atoms (H) are combined, one sigma (σ) bonding orbital (named σ 1s ) is obtained. It is known as linear combination of atomic orbitals (LCAO). -The sum of the two 1s orbitals (one for each atom) yields the bonding orbital. In terms of wave mechanics, this corresponds to a constructive interaction. -The difference of the two orbitals forms the antibonding orbital, s 1s *, i.e. in terms of wave mechanics, this corresponds to a destructive interference. Topic 4. Chemical Bonding (II) 20

21 4. BOND THEORIES Molecular Orbital Theory Bonding interaction + + Addition + 1s A + 1s B σ Distance Antibonding interaction + + Substraction + - Distance 1s A - 1s B σ* Topic 4. Chemical Bonding (II) 21

22 4. BOND THEORIES Molecular Orbital Theory Images from molecularorbital/section1.html Topic 4. Chemical Bonding (II) 22

4. BOND THEORIES Molecular Orbital Theory Molecular orbital for Homonuclear Diatomic Molecules of the second period (Li 2 - N 2 ) Molecular orbital for Homonuclear Diatomic Molecules of the second

23 4. BOND THEORIES Molecular Orbital Theory Molecular orbital for Homonuclear Diatomic Molecules of the second period (Li 2 - N 2 ) Molecular orbital for Homonuclear Diatomic Molecules of the second period (O 2, F 2, Ne 2 ) Image from Atomic Orbitals (O) Molecular Orbitals (O 2 ) Atomic Orbitals (O) Topic 4. Chemical Bonding (II) 23

24 4. BOND THEORIES Molecular Orbital Theory Image from Topic 4. Chemical Bonding (II) 24

5. THE METALLIC BOND First approach to explain the metallic bond: the electron sea model Nuclei are in a sea of electrons (the valence electrons) It explains metallic luster and malleability

25 5. THE METALLIC BOND First approach to explain the metallic bond: the electron sea model Nuclei are in a sea of electrons (the valence electrons) It explains metallic luster and malleability Force applied Topic 4. Chemical Bonding (II) 25

26 5. THE METALLIC BOND Bond theories: the Band theory It is an extension of the MO theory N atoms give N orbitals with very close energies a band is formed. Energy Energy band Li Li 2 Li 3 Li 4 Li n (n N av ) Band theory in alkaline metals Outer shell electron configuration: ns 1 N orbitals, N levels N/2 filled levels Conduction band Band theory in alkaline earths metals Outer shell electron configuration: ns 2 N orbitals, N levels N filled levels p and s bands overlap in energy. Conduction band Valence band Valence band Topic 4. Chemical Bonding (II) 26

27 5. THE METALLIC BOND Band gap E g Band gap E g a) Metal b) Metal c) Semiconductor d) Insulator Elements of Group 14 Tin Germanium Silicon C (Diamond) Gap (ev): 0,08 0,72 1,11 6,0 Topic 4. Chemical Bonding (II) 27

28 Bibliography R. H. Petrucci, W. S. Harwood, F. G. Herring, J. Madura.. Principles and Modern Applications. Prentice Hall, Ninth edition. Chapters R. Chang. Chemistry. McGrawHill, Ninth edition. Chapters J. McMurry, R.C. Fay. Chemistry. Pearson, Fourth edition. Chapters 6, 7. Topic 4. Chemical Bonding (II) 28

Molecular Geometry and Chemical Bonding Theory

Molecular Geometry and Chemical Bonding Theory The Valence -Shell Electron -Pair Repulsion (VSEPR) Model predicts the shapes of the molecules and ions by assuming that the valence shell electron pairs