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Molecular Structure

Molecular Structure

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Molecular Structure

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  1. Molecular Structure Chapter 9

  2. VSEPR MODEL • The shape of a molecule is described by reporting the locations of its atoms.

  3. In the VSEPR model, we focus on one atom at a time, examining the bonds it forms and any lone pairs it may possess. In order to understand and study the properties of molecules, we must be able to recognize and understand the orientation of atoms within a molecule. 

  4. The atoms in a molecule that are bonded to each other share a pair of valence electrons.  Some molecules also have atoms with nonbonding electron pairs.  Electron pairs repel each other, and they want to take up a position in space that will minimize their interactions with other electron pairs, bonded or non bonded.  • The goal of the VSEPR model is to arrange the electron pairs around the central atom so that there is the least amount of repulsions among them.  This occurs when the electron pairs are as far away from each other as possible. 

  5. To report shapes more precisely, we give the bond angles, which are the angles between bonds visualized as straight lines joining the atom centers.

  6. Molecules with multiple bonds • The VSEPR model does not distinguish between a single and multiple bonds. A multiple bond is treated just like other region of high electron concentration.

  7. Class Practice • Suggest a shape for the ethyne molecule,

  8. To understand any molecule, one must first complete a Lewis dot structure. It is then possible to predict the molecular shape using Two basic Principles: • 1. The shapes of molecules are determined by the repulsion between electron pairs in the outer shell of the central atom. Both bond pairs (electron pairs shared by two atoms) and lone pairs (those located on a central atom but not shared) must be considered. • 2.Lone pairs repel more than bond pairs.

  9. The molecule BF3 has a dot symbol as shown below. Here the B atom has 3 bonded pairs in its outer shell. Minimizing the repulsion causes this molecule to have a trigonal planar shape, with the F atoms forming an equilateral triangle about the B atom. The F-B-F bond angles are all 120°, and all the atoms are in the same plane. Central atoms with octet configurations:

  10. The molecule NH3 has a dot symbol much like that for BF3. However, there is a lone pair in the outer shell of the central N atom.  • In NH3 the N has 3 bond pairs and 1 lone pair, (4 total pairs). The shape is called trigonalpyramidal(approximately tetrahedral minus one atom).

  11. 3 groups of electrons (3 bonds around a central atom, no lone pairs) • BF3 Trigonal Planar Molecule • 4 groups of electrons (4 bonds around a central atom, no lone pairs)CF4 Tetrahedral Molecule

  12. 4 groups of electrons (3 bonds around a central atom, 1 lone pair)NH3 Trigonal pyramid • 4 groups of electrons (2 bonds around a central atom, 2 lone pairs)H2O Bent Molecule

  13. 4 groups of electrons (1 bond around a central atom, 3 lone pairs)HF Linear Molecule  • 5 groups of electrons (5 bonds around a central atom, no lone pairs)PF5 Trigonal Bipyramid 

  14. 5 groups of electrons (4 bonds around a central atom, 1 lone pair)SF4 See-saw Molecule • 5 groups of electrons (3 bonds around a central atom, 2 lone pairs)ClF3 T-shaped Molecule

  15. 5 groups of electrons (2 bonds around a central atom, 3 lone pairs)XeF2 Linear Molecule  • 6 groups of electrons (6 bonds around a central atom, no lone pairs)SF6 Octahedral Molecule 

  16. 6 groups of electrons (5 bonds around a central atom, 1 lone pair)BrF5 Square Pyramid • 6 groups of electrons (4 bonds around a central atom, 2 lone pairs)XeF4 Square Planar

  17. Polar bonds • A covalent bond in which the electron pair is shared unequally has partial ionic character and is called a polar covalent bond. • A polar covalent bond is a bond between two atoms that have partial electric charges arising from their difference in electronegativity. Partial charges give rise to an electric dipole moment.

  18. Polar molecules • A diatomic molecule is polar if its bond is polar. A polyatomic molecule is polar if it has polar bonds arranged in space in such a way that their dipoles do not cancel. • Water is a polar molecule because of the way the atoms bind in the molecule such that there are excess electrons on the oxygen side and a lack or excess of positive charges on the Hydrogen side of the molecule.

  19. Water is a polar molecule with positive chargeson one side and negative on the other.

  20. Examples of polar molecules of materials that are gases under standard conditions are: Ammonia (NH3), Sulfur Dioxide (SO2) and Hydrogen Sulfide (H2S).

  21. Home work • Page 413 • 9.52,9.58,9.60,

  22. Bond Strengths in diatomic molecules • The strength of a bond between two atoms is measured by the bond enthalpy . The bond enthalpy typically increases as the order of the bond increases, decreases as the number of lone pairs on neighboring atoms increases, and decreases as the atomic radius increases. • The greater the bond strength the harder it will be to break the bond.

  23. Bond strengths in polyatomic molecules • The first step in predicting the strength of a bond in polyatomic molecule is to identify the atoms and the bond order. An electronegative atom can pull electrons toward itself from more distant parts of the molecule. So all the atoms in the molecule experience a slight pull. Because these variations in strength are not very great, the average bond enthalpy ∆HB is a guide to the strength of a bond in any molecule.

  24. The average bond enthalpies can be used to estimate reaction enthalpies and to predict the stability of a molecule.

  25. Class Practice • Decide whether the reaction : CH3CH2I(g) +H2O- CH3CH2OH(g) +HI(g) is endothermic or exothermic.

  26. Bond lengths • A bond length is the distance between the centers of two atoms joined by a chemical bond. Bond lengths affect the overall size of a molecule. • Bond length is directly related to bond order, when more electrons participate in bond formation the bond will get shorter. Bond length is also inversely related to bond strength and the bond dissociation energy, as a stronger bond is also a shorter bond. In a bond between two identical atoms half the bond distance is equal to the covalent radius.

  27. The covalent radius of an atom is the contribution it makes to the length of the covalent bond. Covalent radii are added together to estimate the lengths of bonds in molecules.

  28. Sigma and Pi bonds • Pi bonds involve the electrons in the leftover p orbital for each carbon atom. Those p orbitals are the electron clouds or orbitals that are shown going up above and below each carbon atom

  29. The distinction between a sigma bond and a pi bond is shown below. The sigma bond has orbital overlap directly between the two nuclei. The pi bond has orbital overlap off to the sides of the line joining the two nuclei.

  30. A single bond is a σ- bond. • A double bond is a σ bond plus one π bond. • A triple bond is a σ bond and two π bonds. • In valence bond theory a bond forms when unpaired electrons in valence shell atomic orbitals on two atoms pair. The atomic orbitals they occupy overlap end to end to form σ bonds or side by side to form π bonds.

  31. Hybridization • Valence Bond theory cannot account for bonding in polyatomic molecules like CH₄. • The electronic configuration of carbon is [He] 2s²2px¹ 2py¹.The 2s orbital is filled in the ground state of the carbon atom. Consider forming hybrids by mixing the 2s orbital with the 2pz orbital. • The configuration of carbon would be [He] 2s¹2px¹ 2py¹2pz¹

  32. Without promotion , a carbon atom has only two unpaired electrons and so can form only two bonds ; after promotion it has four unpaired electrons and can form four bonds. Each bond releases energy as it forms. Despite the energy cost of promoting the electron, the overall energy of methane molecule is lower than it would be if carbon formed only two C-H bonds.

  33. The carbon in methane is now sp³ hybridised. Four hydrogen should form 4 bonds, one with the 2s orbital of carbon and three with the 2p orbitals in carbon.

  34. The four half-filled orbitals have equal energy and are called sp3 hybrid orbitals. The s has a superscript of 1 and the p has a 3, thus meaning that there are four sp3 orbitals, all with the same amount of energy. The hydrogen atoms can now bond to these orbitals. This is how methane forms.

  35. Hybridization also occurs in boron atoms, where it forms sp2 orbitals. One electron from the 2s orbital gets promoted to an empty 2p orbital, thus forming three sp2 hybrid orbitals. • .

  36. The sp hybridization, occurs in beryllium. In this example, one of the two electrons in the 2s orbital gets 'promoted' to an empty 2p orbital, thus forming two sp hybrid orbitals.

  37. Hybrids including d-orbitals • This usually occurs when the central atom is an element in period 3 or later, it uses its d orbital to expand the valence shell to accommodate more than 4 electron pairs. • The atom now uses one d orbital in addition to four s and p orbitals of the valence shell. The resulting orbital is called as sp³d hybrid orbital. This alignment is TrigonalBipyramidal because the five outer atoms form two pyramids (one on top, the other upside down on the bottom) around the central atom. In this model, all atoms are 90 degrees or 120 degrees apart and have the orbital hybridization of "sp3d"

  38. sp3d2 • The final alignment is octahedral, not because it has eight outer atoms, it actually has six, but because it forms eight 90 degree angles. The orbital hybridization is "sp3d2

  39. Paramagnetic and dimagnetic substances • Paramagnetic substance means that it is attracted to a magnetic field. It is a property of unpaired electrons which act as tiny magnets. Paramagnetic substances tend to move into a magnetic field • Most substances contain fully paired electrons and are repelled by a magnetic field ;such substances are diamagnetic.

  40. Limitations of Lewis Theory • Lewis theory cannot account for electron deficient compounds or the paramagnetism of oxygen. In molecular orbital theory, electrons occupy orbitals that spread throughout the molecule. The Pauli exclusion principle allows no more than two electrons to occupy each molecular orbital.

  41. Bonding and antibondingorbitals • Constructive interference produces MO's that places the bonding electrons between the nuclei of the bonding atoms. Electrons are simultaneously attracted by both nuclei which lowers their energies. The MO's produced are called bonding orbitals. • Destructive interference produces MO's that places the electrons away from the space between the nuclei. The nuclei repel which makes the bond less stable (i.e. higher energy). These MO's are called antibonding orbitals.

  42. Molecular orbitals in period 2 diatomic molecules • Consider two helium atoms approaching. Two electrons go into the sigma 1s bonding MO, and the next two into the sigma star antibonding MO. Bond order=(#of electrons in bonding orbital-#of electrons in antibonding orbitals)/2

  43. As antibonding MOs are more antibonding than bonding MOs are bonding, He2 (dihelium), is not expected to exist • And dihelium has a bond order of zero.

  44. Home work • Page 412 • 9.50,9.52,9.58