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14.1 Shapes of molecules and ions (HL)

14.1 Shapes of molecules and ions (HL). 14.1.1 State and predict the shape and bond angles using the VSEPR theory for 5 and 6 negative charge centers. . Molecules with more than 4 electron pairs. Molecules with more than 8 valence electrons [expanded valence shell]

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14.1 Shapes of molecules and ions (HL)

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  1. 14.1 Shapes of molecules and ions (HL) 14.1.1 State and predict the shape and bond angles using the VSEPR theory for 5 and 6 negative charge centers.

  2. Molecules with more than 4 electron pairs • Molecules with more than 8 valence electrons [expanded valence shell] • Form when an atom can ‘promote’ one of more electron from a doubly filled s- or p-orbital into an unfilled low energy d-orbital • Only in period 3 or higher because that is where unused d-orbitals begin

  3. Why does this ‘promotion’ occur? • When atoms absorb energy (heat, electricity, etc…)their electrons become excited and move from a lower energy level orbital to a slightly higher one. • How many new bonding sites formed depends on how many valence electrons are excited.

  4. Exceptions to the octet rule. Shows sulphur achieving 8, 10 and 12 valence electrons due to energy input and excited electrons. • http://www.saskschools.ca/curr_content/chem20/covmolec/exceptns.html

  5. Trigonal Bipyramidal (5 pairs of V.E.)

  6. Trigonal Bipyramidal • Normally would have 3 bp, but the lone pair has moved from the p-orbital to include the d-orbital, allowing for 2 additional bonding sites. • Ex: PCl5

  7. Octahedral (6 pairs of V.E.)

  8. BrF5 is square pyramidal SF6 is octahedral XeF4 is square planar

  9. Bond angles • In general, the greater the bond angle, the weaker the repulsions. • Equatorial- equatorial (120 o) repulsions are weaker than axial- equatorial (90o) repulsions. • Equatorial: lie on the trigonal plane (straight across) • Axial: lies above and below the trigonal plane (up and down)

  10. Remember that lone pairs cause more repulsion than bonding sites, so expect the bond angle to be changed should there be lone pairs, or double or triple bonds involved (multiple bonds also cause more repulsion than expected)

  11. ClF3 PF5 XeO2F2 SOF4 SCl6 IF4+ ICl4- T-shaped Trigonal bipyramidal Seesaw Trigonal bipyramidal Octahedral Seesaw Square planar Practice:

  12. 14.2 Hybridization. 14.2.1 Describe σ (sigma) and π (pi) bonds 14.2.2 State and explain the meaning of the term hybridization 14.2.3 Discuss the relationships between Lewis structures, molecular shapes and types of hybridization (sp, sp2, sp3).

  13. hybridization • the concept of mixing atomic orbitals to form new hybrid orbitals • Used to help explain some atomic bonding properties and the shape of molecular orbitals for molecules. • The valence orbitals (outermost s and p orbitals) are hybridised (mathematically mixed) before bonding, converting some of the dissimilar s and p orbitals into identical hybrid spn orbitals • We must know sp, sp2, and sp3 hydrid orbitals

  14. Hybrid orbitals • Carbon has 4 valence electrons. • 2 electrons paired up in the s-orbital, and 2 electrons unpaired in the p-orbital. • So why does it commonly make 4 bonding sites?

  15. One of carbon’s paired s-orbital electrons is ‘promoted’ to the empty p-orbital • This produces a carbon in an excited state which has 4 unpaired electrons (4 equivalent bonding sites)

  16. sp3 hybrid orbital • formed by mixing the outermost s- and all three outermost p- orbitals to form four sp3 hybrids. • The furthest these four [negatively charged, and therefore repulsive] orbitals can get from each other is the corners of a tetrahedron (109°).

  17. Overlap four s-orbitals from four hydrogens (blue) with four sp3 hybrids on carbon leads to formation of bonds, each containing one electron from the carbon and one from the hydrogen

  18. Examples of sp3 hybrids • Methane, ammonia, water and hydrogen fluoride. • Note that the orbitals not involved in bonding to hydrogen are still hybridised, but end up as lone pairs of electrons (symbolised by the two dots in the diagram above).

  19. sp2 hybrid orbital • formed when only one s- and two p-orbitals are involved. • This leaves one remaining p orbital, which may be involved in forming a double bond.

  20. The furthest these orbitals can get from one another is a trigonal bipyramid, with the sp2 hybrids arranged at 120° to each other in a plane. • This is characteristic of molecules with double bonds.

  21. Finally, sp hybrids are formed using just one s and one p orbital. • Two sp hybrids are formed from them, and the two p-orbitals remaining may contribute to a triple bond. • These arrange themselves at the corners of an octahedron, with the two sp hybrids diametrically opposite one another. • sp hybridisation is characteristic of the triple bond. (1 σ-bond and 2 π (pi) bonds)

  22. Sigma bond (σ-bond) • When s and/or hybrid orbitals overlap 'end-on', sigma bonds (σ) are formed • They have a single area of electron density between the nuclei of the two atoms whose orbitals are overlapping. • In the diagrams below, σ bond is shown

  23. Sigma bond (σ-bond) • results from head-on overlap of orbitals • electron density is symmetric about the internuclear axis: between nuclei.

  24. π (pi) bonds • p orbitals can overlap sideways too: when this happens two lobes of electron density are formed between the atoms. • From the diagram, you can see that the double bond in ethene is composed of one σ plus one π bond,

  25. π (pi) bonds • results from sideways overlap of orbitals • bonds resulting from the combination of parallel p orbitals • electron density is above and below the internuclear axis.

  26. Predicting shape • The shape is dictated by the σ-bonds and the non-bonding electron pairs (lone pairs) • π-bonds do not affect the shape of the molecule (double bonds or triple bonds) • That’s why we refer to bonding sites when using VSEPR, not paying attention to whether it was single, double or triple bonded.

  27. 14.3 Delocalization of electrons 14.3.1 Describe the delocalization of (pi) π- electrons and explain how this can account for the structure of some species

  28. Delocalised electrons • The term 'delocalised' refers to an electron which is not 'attached' to a particular atom or to a specific bond. • Delocalized electrons are contained within an orbital that extends over several adjacent atoms. • Classically, delocalized electrons can be found in double bonds and in aromatic systems • Double bonds = 1 sigma and 1 pi bond • Delocalisation is often represented with resonance structures or resonance hybrid

  29. Resonance structures • the nitrate ion can be viewed as if it resonates between the three different structures above. • Nitrate doesn’t change from one to the next, but behaves as a combination of all structures

  30. Resonance is possible whenever a Lewis structure has a multiple bond and an adjacent atom with at least one lone pair. • The following is the general form for resonance in a structure of this type.

  31. Practice • Try to show the individual Lewis structures for the HCO3- ion • Show its resonance structure too

  32. NO3- NO2- CO32- O3 RCOO- Benzene (C6H6) TOK Kekule claimed that the inspiration for the cyclic structure of benzene came from a dream. What role do the less rational ways of knowing play in the acquistion of scientific knowledge? Practice drawing these resonance structures:

  33. Bibliography and sites to visit • http://www.tutorvista.com/content/chemistry/chemistry-iii/chemical-bonding/types-covalent-bonds.php • Good site on types of covalent bonds • http://www.mikeblaber.org/oldwine/chm1045/notes/Geometry/VSEPR/Geom02.htm • Used for expanded valence shell pictures • http://www.kentchemistry.com/links/bonding/lewisdotstruct.htm • Puts the lewis diagrams together and explain them. Including expanded shell

  34. http://www.mpcfaculty.net/mark_bishop/resonance.htm • Resonance structures pictures and notes • http://en.wikipedia.org/wiki/Delocalization • Notes on delocalisation of electrons • http://www.steve.gb.com/science/atomic_structure.html • Amazing website for hybrid orbitals • http://library.thinkquest.org/C006669/data/Chem/bonding/shapes.html • Good review of all shapes

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