Bonding and Intermolecular Forces

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Shapes of Molecules and Ions

  • The main reason why covalent bonds form is because when electrons are shared, the situation is more stable than when there are two seperate atoms. 
  • We draw dot and cross diagrams to see how the electrons are shared in a compound.

                                   (http://www.bbc.co.uk/schools/gcsebitesize/science/images/diag_ammonia.gif)

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Shapes of Molecules and Ions (cont)

  • The electron pair repulsion theory is a model that predicts the shape of a molecule around a central atom. 
  • Using a dot and cross diagram, we count the number of bonding electron pairs and lone electron pairs. From this we can determine the bond angle and the molecular shape.
  • The electron pairs repel eachother to form a shape of minimum repulsion and maximum separation. 
  • Lone pairs repel more than bonding pairs because they are attracted to a single nucleus (not 2 atoms).
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Shapes of Molecules and Ions (cont)

(http://www.examstutor.com/chemistry/resources/studyroom/bonding/shapes_of_molecules/pictures/table_of_shapes.gif)

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Shapes of Molecules and Ions (cont)

  • For determining the molecular shape and bond angle, electrons in a double/triple bond count only as 'one pair', just like a single bond.
  • Single bonds are sigma bonds.
  • Double bonds are 1 sigma + 1 pi bond. 
  • There is no free rotation about a double bond because otherwise this would break the pi bond.
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Shapes of Molecules and Ions (cont)

              (http://www.scienceaid.co.uk/chemistry/organic/images/pisigmabonds.jpg)

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Shapes of Molecules and Ions (cont)

  • The shape of an ion is predicted in the same way as the shape of a molecule.
  • The thing to look out for is that each negative charge means there is an extra electron; and each positive charge means the ion is missing an electron.
  • For example, in NH4+ (ammonium ion), the outer shell has 8 not 9 electrons, all of which bond. 
  • Note that this means NH4+ has dative covalent bonds, in which both electrons in the bond come from the same atom. 
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Shapes of Molecules and Ions (cont)

               (http://hsc.csu.edu.au/chemistry/core/monitoring/chem944/ammonium.gif)

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Shapes of Molecules and Ions (cont)

  • Carbon has 4 allotropes, each of which differ in molecular structure due to differences in bonding.
  • In diamond, each carbon atom forms 4 identical bonds to neighbouring carbon atoms (tetrahedral).
  • Graphite consists of carbon atoms in layers. Within a layer, each carbon atom is strongly bonded to 3 others (trigonal planar). The layers are only weakly bonded to eachother by london forces. 
  • The 4th outer electron from each carbon is delocalised and free to move, making graphite a good electrical conductor. 
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Shapes of Molecules and Ions (cont)

  • Fullerenes consist of 32 or more carbon atoms. Buckminsterfullerene has 60. Each carbon atom is bonded to 3 others to make a ball-shaped molecule. The 4th electron is delocalised, however it cannot move between molecules. So, fullerens cannot conduct electricity.
  • Nanotubes are fullerenes in the form of tubes. They are extremely small and stiff. In polymer form they obtain good electrical conductivity and enormous strength. 
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Bond Polarity and Intermediate Bonding

  • The electrons shared between atoms in a covalent bond are not always shared equally.
  • The electronegativity of an element measures the ability of an atom to attract an electron pair to itself in a covalent bond.
  • The most electronegative elements are flourine, oxygen and nitrogen.
  • If there is a big enough difference in electronegativity between 2 atoms, the electron density of the bond is distorted, causing a dipole. The resulting bond is polarised (1 atom +, other -). 
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Bond Polarity and Intermediate Bonding (cont)

  • A molecule that contains polar bonds will be a polar molecule, unless the polar bonds are symetrically arranged.
  • In this case the net effect of the polar bonds cancels out, and the resulting molecule is non-polar. eg CO2 or CCl4. 

                                             (http://www.chemguide.co.uk/atoms/bonding/ccl4.GIF)

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Intermolecular Forces

  • Intermolecular forces are forces between molecules (not between atoms).
  • They are much weaker than covalent bonds. 
  • There are 3 types of intermolecular force: London forces, perminant dipole-dipole interactions, and hydrogen bonds. 
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Intermolecular Forces (cont)

  • London or Van der Waals forces exist between all molecules.
  • London forces are caused by an instantaneous unequal electron distribution in a molecule, which creates an instantaneous dipole.
  • This causes an induced dipole in the opposite direction on a neibouring molecule. The two dipoles attract. 
  • The more electrons in a molecule, the larger the induced dipole, and the stronger the london forces. 
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Intermolecular Forces (cont)

  • Perminant dipole-dipole forces occur between perminant dipoles in neighbouring polar molecules.
  • The slightly +ve part of one molecules dipole attracts the slightly -ve part of another molecules dipole.
  • These forces are stronger than london forces. 
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Intermolecular Forces (cont)

  • Hydrogen bonds are a special type of perminant dipole-dipole force. 
  • When a hydrogen atom is bonded to a highly electronegative atom (flourine, oxygen or nitrogen), the dipole formed is very large.
  • The slightly +ve H atom on one molecule forms a very strong bond with the electronegative atom on another molecule: a hydrogen bond.
  • The hydrogen bond is at 180 degrees to the normal covalent bond between oxygen and hydrogen in water. 
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Intermolecular Forces (cont)

  • The stronger the intermolecular forces between molecules, the higher the boiling/melting point of that substance.
  • Straight chain molecules (aliphatic) have a higher boiling point than branched molecules with the same number of electrons. This is because straight chains can line up with eachother easier to get a larger surface area for intermolecular forces to appear. 
  • High boiling point = low volatility 
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Solubility

  • In terms of intermolecular forces, for solubility the general rule is 'like dissolves like'.
  • The solubility of one liquid in another depends on:        - The intermolecular forces in each seperate liquid before mixing.                                                                - The potential new intermolecular forces that could form between the two liquids when mixed.
  • The energy released from making the new intermolecular bonds has to be equal to if not greater than the energy required to break the initial intermolecular forces. 
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