Chemistry module 2.5

    Cards (39)

    • Molecules
      Can adapt the following shapes and bond angles
    • Molecules of different shapes
      • Corresponding bond angles
    • Valence shell electron pair repulsion theory (VSEPR)
      • Predicts the shape and bond angles of molecules
      • Electrons are negatively charged and will repel other electrons when close to each other
      • Bonding pairs of electrons will repel other electrons around the central atom forcing the molecule to adopt a shape in which these repulsive forces are minimised
    • Valence shell electrons are those electrons that are found in the outer shell
    • Electron pairs
      • Repel each other as they have the same charge
      • Lone pair electrons repel each other more than bonded pairs
      • Repulsion between multiple and single bonds is treated the same as for repulsion between single bonds
      • Repulsion between pairs of double bonds are greater
    • The most stable shape is adopted to minimize the repulsion forces
    • Lone pairs of electrons
      • Have a more concentrated electron charge cloud than bonding pairs of electrons
      • The cloud charges are wider and closer to the central atom's nucleus
    • Order of repulsion
      Lone pairlone pair > lone pair – bond pair > bond pair – bond pair
    • Electronegativity
      The power of an atom to attract the pair of electrons in a covalent bond towards itself
    • The electron distribution in a covalent bond between elements with different electronegativities will be unsymmetrical
    • Pauling scale

      Used to assign a value of electronegativity for each atom
    • Fluorine is the most electronegative atom on the Periodic Table, with a value of 4.0 on the Pauling Scale
    • Factors affecting electronegativity
      • Nuclear charge
      • Atomic radius
      • Shielding
    • Down a group
      There is a decrease in electronegativity going down the group
    • Across a period
      Electronegativity increases across a period
    • Polarity
      • When two atoms in a covalent bond have the same electronegativity the covalent bond is nonpolar
      • When two atoms in a covalent bond have different electronegativities the covalent bond is polar and the electrons will be drawn towards the more electronegative atom
    • Dipole moment
      • A measure of how polar a bond is
      • The direction of the dipole moment is shown by the sign in which the arrow points to the partially negatively charged end of the dipole
    • To determine whether a molecule with more than two different atoms is polar, the polarity of each bond and how the bonds are arranged in the molecule must be considered
    • CH3Cl
      • Has four polar covalent bonds which do not cancel each other out, causing CH3Cl to be a polar molecule
    • Though CCl4 has four polar covalent bonds, the individual dipole moments cancel each other out
    • Dipole moment
      The direction of the dipole moment and the arrow points to the delta negative end of the dipole
    • Assigning polarity to molecules
      1. Consider the polarity of each bond
      2. Consider how the bonds are arranged in the molecule
    • Some molecules have polar bonds but are overall not polar because the polar bonds in the molecule are arranged in such way that the individual dipole moments cancel each other out
    • CH3Cl has four polar covalent bonds which do not cancel each other out causing CH3Cl to be a polar molecule; the overall dipole is towards the electronegative chlorine atom
    • CCl4 has four polar covalent bonds, but the individual dipole moments cancel each other out causing CCl4 to be a nonpolar molecule
    • Intramolecular forces

      Forces within a molecule, usually covalent bonds
    • Intermolecular forces
      Weaker forces between molecules
    • Types of intermolecular forces
      • Induced dipole - dipole forces (London dispersion forces or van der Waals' forces)
      • Permanent dipole - dipole forces (van der Waals' forces)
      • Hydrogen Bonding
    • Intramolecular forces are stronger than intermolecular forces
    • Induced dipole-dipole forces
      • Exist between all atoms or molecules, also known as van der Waals' forces or London dispersion forces
      • Caused by the constant movement of the electron charge cloud, creating temporary dipoles that can induce dipoles in neighbouring molecules
      • Stronger for molecules with more electrons or higher relative molecular mass
    • Permanent dipole-dipole forces
      Forces between two molecules that have permanent dipoles, where the δ+ end of one dipole and the δ- end of another are attracted
    • Permanent dipole-dipole forces are stronger than induced dipole-dipole forces in smaller molecules with an equal number of electrons
    • Hydrogen Bonding
      • The strongest form of intermolecular bonding, a type of permanent dipole-dipole bonding
      • Requires a species with an O, N or F atom bonded to H, where the H becomes highly δ+ and can bond with the lone pair of an O, N or F in another molecule
    • Properties of water
      • High melting and boiling points
      • High surface tension
      • Anomalous density of ice compared to water
    • The high melting and boiling points of water are caused by the strong intermolecular forces of hydrogen bonding between the molecules
    • Ice has a lower density than liquid water due to the 'more open' structure of the hydrogen-bonded network in the solid state
    • Iodine
      • Weak intermolecular forces (instantaneous dipole-induced dipole) hold the molecular lattice together
      • Tends to sublime at temperatures approaching 114°C due to the weak intermolecular forces
      • Forms a purple vapour when it sublimates
    • Solubility
      • 'Like dissolves like' - non-polar substances dissolve in non-polar solvents, polar covalent substances dissolve in polar solvents
      • Larger polar covalent molecules may have decreased solubility as the polar part is a smaller proportion of the overall structure
      • Giant covalent substances generally don't dissolve in any solvents
    • Conductivity
      • Covalent substances cannot conduct electricity in solid or liquid state as they lack freely moving charged particles
      • Some polar covalent molecules can ionise and conduct electricity
      • Some giant covalent structures can conduct due to delocalised electrons