14.1 Shapes of molecules and ions (HL) · 14.1 Shapes of molecules and ions (HL) The octet is the...
Transcript of 14.1 Shapes of molecules and ions (HL) · 14.1 Shapes of molecules and ions (HL) The octet is the...
14.1 Shapes of molecules and ions (HL)
● The octet is the most common electron arrangement because of its stability.
● Exceptions:a) Fewer electrons (incomplete octet) if the central atom is a small atoms, e.g. Be and B
b) More than eight electrons (expanded octet) if the central atom is a 3rd row element or below, e.g. P and S
Species with five negative charge centres
● The shape of a molecule or ion can be predicted by the valence shell electron pair repulsion theory (VSEPR).
● Pairs of electrons (=negative charge centres) arrange themselves around the central atom so that they are as far apart from each other as possible.
● If a molecule has five charge centres and they all are bonding electrons, the shape is triangular bipyramidal.
● If one or more of these five negative charge centres is a non-bonding pair, this will influence the final shape of the molecule.
● One: Tetrahedron
● Two: T-shaped ClF3
● Three: Linear I3
-
Species with six negative charge centres
● Molecules with six charged centres that are all bonding have an octahedral shape, e.g. SF
6.
● One non-bonding pair: square pyramidal BrF5
● Two non-bonding pairs: square planar XeF4
Rivi 1 Rivi 2 Rivi 3 Rivi 40
2
4
6
8
10
12
● Ex. Predict the shape and bond angles of :
PF5, PF
6
-
● Homework: p.129 Ex.14, 15,16● p.155 Ex. 11, 16 a) and b), 17 d)
14.2 Hybridization
● The Lewis structure is a useful model, but it makes one false assumption:
– It assumes that all eight electrons are equal.
● The energy of the electrons are not equal, since some of them exist in s sub-levels and other in p sub-levels.
● A more advanced model of bonding is called the molecular orbital theory.
Molecular orbital theory
● When a bond is formed, atomic orbitals overlap to form new molecular orbitals that are lower in energy.
Sigma (σ) bonds
● A sigma bond is formed when two atomic orbitals on different atoms overlap along a line drawn through the two nuclei (”head on”).
pi (π) bonds
● A pi bond is formed when two p orbitals overlap ”sideways on”.
● The electron density is concentrated in two regions, above and below the plane of the bond axis.
sp3-hybridization
● Methane contains four equal C-H bonds.
● When the carbon bonds to hydrogens, the one 2s and the three 2p orbitals hybridize to form four new energetically equal hybrid orbitals.
● These four sp3-orbitals arrange themselves tetrahedrally, bond angle 109,5º, and four equal σ bonds are formed with the hydrogen.
http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/bom5s2_6.swf
● Hybridization is not just restricted to carbon compounds.
● Tetrahedral compounds (i.e. 4 negative charge
centers) have sp3-hybridization:
NH4
+, NH3, PCl
3 , H
2O
Other compounds with sp3 hybridization
sp2-hybridization
● In ethene, the one 2s orbital hybridizes with two 2p orbitals.The remaining 2p orbital stays as it is.
● These orbitals form three sigma bonds with carbon and 2 hydrogens, bond angle 120o.
● The remaining p orbitals of carbon form a pi bond.
● Therefore, the double bond consists of one sigma and one pi bond.
Other compounds with sp2 hybridization
● Planar triangular compounds (i.e. 3 negative charge
centers),e.g.
BF3 , SO
3, SO
2, propanone
sp hybridization
● In ethyne, the 2s orbital hybridizes with one 2p orbital. The remaining two 2p orbitals stay as they are.
● These orbitals form two sigma bonds with carbon and hydrogen, bond angle 180o.
● The remaining 2 p orbitals of carbon form two pi bonds.
● Therefore, the triple bond consists of one sigma and two pi bonds.
Other compounds with sp hybridization
● Linear compounds (i.e. 2 negative charge
centers), e.g. N2, HCN
14.3 Delocalization of electrons
● http://www.youtube.com/watch?v=eDiDl-ByUd4
Delocalization of electrons in benzene● Delocalization of electrons can occur whenever alternate
double and single bonds occur between carbon atoms.
● The carbon atoms in benzene are sp2-hybridized and each carbon has a p orbital containing one electron.
● Instead of forming double and single bonds, the electrons are delocalized over all six carbon atoms.
Resonance structures
● When writing the Lewis structure for some molecules, it is possible to write more than one structure.
● For example benzene can be written:
● These two structures are known as resonance hybrides. The true structure lies somewhere in between the two.
Other common compounds with resonance structures
● Resonance occurs when more than one valid Lewis structure can be drawn for a molecule.
● NO3
-
● NO2
-
● CO3
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● O3
● RCOO-
Properties of species with delocalized electrons
1. intermediate bond lengths and strengths
● All affected bonds have equal bond strengths and the bond lengths are intermediate between those of single and double bonds.
● Bond length is dependend on the number of bonds between the atoms: triple < double < single bond
● Bond order: number of shared electron pairs
number of bonding positions
● The higher the bond order, the greater the electron density (=the shorter the bond).
2. Stability
● Delocalization spreads the electrons as far apart as possible and therefore minimizes the repulsion between them.
● Ex. 1 The delocalization makes the benzene molecule more stable by ca. 150 kJ/mol. This is called the delocalization enthalpy or resonance energy.
● This makes the molecule less chemically reactive, since this extra energy has to be put in to break the bonds.
● Ex 2. The relative stability of R-O- ions depends on to what extent the negative charge is delocalized between the two bonds:
● The more stable the ion, the more likely is it formed in a reaction.
3. Electrical conductivity
● Both metals and graphite have delocalized electrons spread out through the entire structure and thus conduct electricity.