Chemistry Alevel Aqa PMT

• Partial charges


  Caused by polar bonds
• Dipole moment


  Permanent dipole moment shown using an arrow with a line through it
• Electron pairs


  • Exist in pairs in their orbitals
  • Each pair of electrons around an atom will repel all other electrons
  • Pairs of electrons will take up positions as far apart as possible to minimise repulsion

  
• Two pairs of electrons


  1. The two bonding pairs arrange themselves at 180° to each other

  2. The molecule is described as being linear

  
• Three pairs of electrons

  1. The 3 pairs arrange themselves as far apart as possible

  2. They all lie in one plane at 120° to each other

  3. The arrangement is called trigonal planar
• Four pairs of electrons

  1. Four electron pairs arrange themselves in a tetrahedral arrangement

  2. A tetrahedron is a regular triangularly-based pyramid

  3. The carbon atom would be at the centre and the hydrogens at the four corners

  4. All the bond angles are 109.5°

  
• Five pairs of electrons

  1. The 5 electron pairs take up a shape described as a trigonal bipyramid

  2. Three of the chlorines are in a plane at 120° to each other

  3. The other two are at right angles to this plane

  4. The trigonal bipyramid therefore has two different bond angles - 120° and 90°

  
• Six pairs of electrons

  1. 6 electrons in the outer level of the sulphur, plus 1 each from the six fluorines, makes a total of 12 - in 6 pairs

  2. Because the sulphur is forming 6 bonds, these are all bond pairs

  3. They arrange themselves entirely at 90°, in a shape described as octahedral

  
• Ammonia, NH3


  • Nitrogen is in group 5 and so has 5 outer electrons
  • Each of the 3 hydrogens is adding another electron to the nitrogen's outer level, making a total of 8 electrons in 4 pairs
  • Because the nitrogen is only forming 3 bonds, one of the pairs must be a lone pair
  • The shape is that of a triangular pyramid

  
• Bond pair – lone pair repulsion


  • Bond pair – lone pair repulsion is greater than bond-pair – bond pair repulsion
  • This extra repulsion reduces the bond angles from 109.5° to 107°

  
• Water, H2O


  • Oxygen has four pairs of electrons, two of which are lone pairs
  • These will again take up a tetrahedral arrangement
  • The bond angle closes slightly more to 104°, because of the repulsion of the two lone pairs
  • The shape isn't described as tetrahedral, because we only "see" the oxygen and the hydrogens - not the lone pairs
  • Water is described as bent or V-shaped

  
• Chlorine Tetrafluoride, ClF4


  • Chlorine is in group 7 and so has 7 outer electrons
  • The four fluorines contribute one electron each, making a total of 12 - in 6 pairs
  • The chlorine is forming three bonds - leaving you with 4 bonding pairs and 2 lone pairs, which will arrange themselves into a square bipyramid
  • The ion has a flat square shape described as square planar with 90° bond angles

  
• Molecular shapes


  • Tetrahedral
  • Trigonal bipyramid
  • Octahedral
• For each molecule or ion

  1. Show its Lewis structure (dot-cross)

  2. Draw its 3D structure

  3. Name its shape

  4. Give its bond angles

  
• Covalent Bonding


  When two non-metal atoms bond, one can not supply electrons to fill the outer shell of another. Instead they form covalent bonds.
• Covalent Bond


  A pair of electrons shared between two atoms. The atoms are held together because the electron pair is attracted by both of the nuclei.

  
• Unpaired Electron


  An electron that occupies an orbital of an atom singly.
• Lone Pair


  A pair of electrons sharing an orbital in the outer shell of an atom.
• Formation of Covalent Bond

  Unpaired electrons from two different atoms are shared.

  
• Draw dot and cross diagrams for the following molecules

  • H2O
  • NH3
  • CO2

  
• For each molecule identify

  1. Bonding pairs (pairs of electrons forming covalent bonds)

  2. Lone pairs

  
• Write the formula for each of the compounds below and state the type of bonding found

  • Aluminium oxide
  • Hydrogen fluoride
  • Potassium sulfide
  • Ethane

  
• Draw a dot and cross diagram for

  1. Aluminium oxide

  2. Hydrogen fluoride

  
• Boron trifluoride, BF3


  • Boron atom only has 3 electrons in its outer level, and there is no possibility of it reaching a noble gas structure by simple sharing of electrons. The boron forms the maximum number of bonds that it can in the circumstances.
• Boron trifluoride molecule


  • Has a trigonal planar shape. The electron pairs in the covalent bonds repel each other, and therefore the bond angles are all 120°.

  
• Phosphorus(V) chloride, PCl5


  • Phosphorus forms two chlorides - PCl3 and PCl5. When phosphorus burns in chlorine both are formed, with the majority product depending on how much chlorine is available.
• Draw the structural formula of the following molecules showing their lone pairs

  1. Ammonia

  2. Water

  3. Hydrogen chloride

  
• Coordinate (Dative) Bond


  A covalent bond in which both electrons come from the same atom.

  
• Reaction between ammonia and boron trifluoride, BF3

  1. Name the type of bond formed

  2. Describe how the bond is formed

  3. Show the structural formula of the molecule formed

  
• Electronegativity


  A measure of the tendency of an atom to attract a bonding pair of electrons

  
• Pauling scale


  The most commonly used scale for electronegativity, where fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to caesium and francium which are the least electronegative at 0.7
• Electronegativity


  • Depends on the nuclear charge, the distance between the nucleus and the outer electrons, and the shielding of the nuclear charge by electrons in inner shells
  • Larger nuclear charge = larger electronegativity
  • Larger distance = smaller electronegativity
  • More shells = smaller electronegativity
• Polarity


  The unequal sharing of electrons between atoms that are bonded together covalently

  
• If two atoms of equal electronegativity bond together, the bonding pair of electrons will be found on average half way between the two atoms

  
• If atom B is slightly more electronegative than atom A, B will attract the electron pair rather more than A does, resulting in a polar bond
• If atom B is a lot more electronegative than atom A, the electron pair will be dragged right over to B's end of the bond, forming ions
• Types of intermolecular forces

  • van der Waals forces
  • Dipole-dipole forces
  • Hydrogen bonding

  
• Dipole-dipole forces


  Intermolecular forces that act between molecules that have permanent dipoles

  
• Carbon dioxide (CO2) is a linear molecule with two polar bonds, but the dipoles cancel out resulting in no overall permanent dipole moment, making it a non-polar molecule

  
• Water (H2O) is a bent molecule with two polar bonds, and the dipoles do not cancel out, resulting in a permanent dipole moment, making it a polar molecule