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14.1 Shapes of molecules and ions (HL)

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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. – PowerPoint PPT presentation

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Title: 14.1 Shapes of molecules and ions (HL)


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/covm
    olec/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
Practice
  • ClF3
  • PF5
  • XeO2F2
  • SOF4
  • SCl6
  • IF4
  • ICl4-
  • T-shaped
  • Trigonal bipyramidal
  • Seesaw
  • Trigonal bipyramidal
  • Octahedral
  • Seesaw
  • Square planar

12
14.2 Hybridization.
  • 14.2.1 Describe s (sigma) and p (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 carbons 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
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17
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).

18
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
19
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).

20
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.

21
  • 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.

22
  • 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 s-bond and 2 p (pi) bonds)

23
Sigma bond (s-bond)
  • When s and/or hybrid orbitals overlap 'end-on',
    sigma bonds (s) 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, s bond is shown

24
Sigma bond (s-bond)
  • results from head-on overlap of orbitals
  • electron density is symmetric about the
    internuclear axis between nuclei.

25
p (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 s plus one p
    bond,

26
p (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.

27
Predicting shape
  • The shape is dictated by the s-bonds and the
    non-bonding electron pairs (lone pairs)
  • p-bonds do not affect the shape of the molecule
    (double bonds or triple bonds)
  • Thats why we refer to bonding sites when using
    VSEPR, not paying attention to whether it was
    single, double or triple bonded.

28
14.3 Delocalization of electrons
  • 14.3.1 Describe the delocalization of (pi) p-
    electrons and explain how this can account for
    the structure of some species

29
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

30
Resonance structures
  • the nitrate ion can be viewed as if it resonates
    between the three different structures above.
  • Nitrate doesnt change from one to the next, but
    behaves as a combination of all structures

31
  • 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.

32
Practice
  • Try to show the individual Lewis structures for
    the HCO3- ion
  • Show its resonance structure too

33
Practice drawing these resonance structures
  • 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?

34
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35
Bibliography and sites to visit
  • http//www.tutorvista.com/content/chemistry/chemis
    try-iii/chemical-bonding/types-covalent-bonds.php
  • Good site on types of covalent bonds
  • http//www.mikeblaber.org/oldwine/chm1045/notes/Ge
    ometry/VSEPR/Geom02.htm
  • Used for expanded valence shell pictures
  • http//www.kentchemistry.com/links/bonding/lewisdo
    tstruct.htm
  • Puts the lewis diagrams together and explain
    them. Including expanded shell

36
  • http//www.mpcfaculty.net/mark_bishop/resonance.ht
    m
  • Resonance structures pictures and notes
  • http//en.wikipedia.org/wiki/Delocalization
  • Notes on delocalisation of electrons
  • http//www.steve.gb.com/science/atomic_structure.h
    tml
  • Amazing website for hybrid orbitals
  • http//library.thinkquest.org/C006669/data/Chem/bo
    nding/shapes.html
  • Good review of all shapes
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