bonding n structure

• Electronegativity

  The tendency of an atom or radical to attract electrons in a covalent bond.
• Most electronegative element

  Fluorine
• Least electronegative element

  Francium
• Ionic bond


  The electrostatic attraction between cations and anions.
• Metallic bond


  The electrostatic attraction between cations and the sea of delocalised electrons.
• Covalent bond

  The electrostatic attraction between the nuclei and the shared pair of electrons.
• Dative covalent bond

  • Also known as a 'co-ordinate' bond

  
  

  • A covalent bond where both electrons come from one atom
• Polar covalent bond

  A covalent bond in which electrons are not shared equally.
• Hydrogen bond

  • A type of intermolecular force between a delta positive hydrogen atom of one molecule and a lone pair of electrons on an atom of a neighbouring molecule.

  
  

  • A type of dipole-dipole interaction

  
  

  • The strongest intermolecular force

  
  

  • The hydrogen atom must be covalently bonded to atoms of one of the three electronegative elements fluorine, oxygen or nitrogen.

  
  

  • The atom of the electronegative element makes the hydrogen atom sufficiently delta postive.
• Simple molecule

  Two to one thousand atoms that covalently bonded. There are intermolecular forces between the molecules.
• Macromolecule

  One thousand or more atoms that are covalently bonded.
• Examples of non-polar simple molecules with non-polar bonds

  • Methane

  • Hydrogen gas

  • Chlorine gas

  • Fullerene, C60
• Examples of polar simple molecules

  • Ammonia

  • Water

  • Chloromethane
• Examples of non-polar simple molecules with polar bonds

  • Carbon dioxide

  • Boron trifluoride

  • Tetrafluoromethane

  
  

  These molecules have symmetry
• Explanation for why molecules with polar bonds can be non-polar

  • The molecules have symmetry

  • The dipoles cancel out
• Dipole

  A molecule that has two regions with opposite charges
• Describe and explain the trend in melting points for the period 2 fluorides

  • Melting points increase

  • The strength of ionic bonding increases

  • This is because the radius of the cation decreases and the charge increases
• Hydronium ion

  • H3O+

  
  

  • Formed when water reacts with H+ ions
• Describe and explain the trend in melting points for the period 2 chlorides

  • Melting points decrease

  • The group 1 halide is an ionic compound that has a high melting point

  • The group 3 halide is a simple molecule that has weak intermolecular forces and a low melting point
• Polyatomic ion

  An ion made of two or more atoms, such as sulfate
• Halide ion

  Ion of a halogen element
• London dispersion forces

  • The weakest intermolecular force between ALL molecules, both polar and non-polar.

  
  

  • Non-polar molecules and noble gases ONLY have this type of intermolecular force

  
  

  • They arise by random movements of electrons that cause molecules to temporarily have a slight charge.

  
  

  • The more electrons there are, the greater the partial charges and the forces of attraction and so more thermal energy is required to break these.
• Explanation for how London dispersion forces arise

  1) An atom or molecule has electrons that move randomly

  2) By chance, there are more electrons on one side than the other

  3) Forming a temporary dipole (uneven distribution of charge)

  4) This causes a dipole to be induced on a neighbouring atom or molecule

  5) Atoms or molecules are attracted δ- to δ+
• Molecules that have hydrogen bonds

  • Hydrogen fluoride

  • Ammonia

  • Water

  • Ethanol
• Dipole-dipole interactions

  • Attractions between oppositely charged regions of polar molecules

  
  

  • Molecules are attracted δ- to δ+

  
  

  • All molecules that have this force ALSO have London dispersion forces
• Explain why water (H2O) has a higher melting point than hydrogen disulfide (H2S)

  • Both molecules have London dispersion forces and dipole-dipole interactions

  
  

  • Water also has hydrogen bonds
• Description of the trend in melting point for the hydrogen halides

  • Melting points decrease then increase
• Explanation for the trend in melting point for the hydrogen halides

  • Hydrogen fluoride has the highest melting point as it has hydrogen bonds , weak dipole-dipole interactions and London dispersion forces. It is the only hydrogen halide to have hydrogen bonds.

  
  

  • Hydrogen chloride has the lowest melting point as it has the weakest dipole-dipole interactions and weakest London forces.

  
  

  • Melting points increase from HCl to HI because both dipole-dipole interactions and London dispersion forces increase

  
  

  • This is because the number of shells of the halogen atom increases from 2 for chlorine to 5 for iodine and so there are more electrons
• Number of lone pairs of electrons on the oxygen atom of water

  2
• Number of lone pairs of electrons on the nitrogen atom of ammonia

  1
• Lewis acid

  A substance that can accept a pair of electrons, e.g. H+
• Lewis base

  A substance that can donate an electron pair, e.g. nitrogen of ammonia
• Shape of a methane molecule

  Tetrahedral
• Shape of an ammonia molecule

  Trigonal pyramidal
• Shape of a water molecule

  Bent
• Shape of a carbon dioxide molecule

  Linear
• Shape of a sulfur hexafluoride molecule

  Octahedral
• Shape of a boron trifluoride molecule

  Trigonal planar
• Shape of a phosphorus pentafluoride molecule

  Trigonal bipyrimidal
• Bond angle in hydrogen fluoride

  There is no bond angle as the molecule does not have three or more atoms.