14.1 Shapes of molecules and ions (HL)

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Exceptions: a) Fewer electrons (incomplete octet) if the central atom is a small

1 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 3 rd 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).

2 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.

3 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 ClF 3 Three: Linear I 3 -

4 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 BrF 5 Two non-bonding pairs: square planar XeF 4

5 Rivi 1 Rivi 2 Rivi 3 Rivi 4

6 Ex. Predict the shape and bond angles of : PF 5, PF 6 - Homework: p.129 Ex.14, 15,16 p.155 Ex. 11, 16 a) and b), 17 d)

7 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.

8 Molecular orbital theory When a bond is formed, atomic orbitals overlap to form new molecular orbitals that are lower in energy.

9 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 ).

10 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.

11 sp 3 -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 sp 3 -orbitals arrange themselves tetrahedrally, bond angle 109,5º, and four equal σ bonds are formed with the hydrogen.

12 Other compounds with sp 3 hybridization Hybridization is not just restricted to carbon compounds. Tetrahedral compounds (i.e. 4 negative charge centers) have sp 3 -hybridization: NH 4+, NH 3, PCl 3, H 2 O

13 sp 2 -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 120 o.

14 The remaining p orbitals of carbon form a pi bond. Therefore, the double bond consists of one sigma and one pi bond.

15 Other compounds with sp 2 hybridization Planar triangular compounds (i.e. 3 negative charge centers),e.g. BF 3, SO 3, SO 2, propanone

16 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 180 o.

17 The remaining 2 p orbitals of carbon form two pi bonds. Therefore, the triple bond consists of one sigma and two pi bonds.

18 Other compounds with sp hybridization Linear compounds (i.e. 2 negative charge centers), e.g. N 2, HCN

19 14.3 Delocalization of electrons

Delocalization of electrons in benzene Delocalization of electrons can occur whenever alternate double and single bonds occur between carbon atoms.

20 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 sp 2 -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.

21 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.

22 Other common compounds with resonance structures Resonance occurs when more than one valid Lewis structure can be drawn for a molecule. NO 3 - NO 2 -

23 CO 3 - O 3 RCOO -

24 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).

25 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.

26 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.

27 3. Electrical conductivity Both metals and graphite have delocalized electrons spread out through the entire structure and thus conduct electricity.

General and Inorganic Chemistry I.

General and Inorganic Chemistry I. Lecture 1 István Szalai Eötvös University István Szalai (Eötvös University) Lecture 1 1 / 29 Outline István Szalai (Eötvös University) Lecture 1 2 / 29 Lewis Formulas