Unit 1 Chemistry

• Across a period from left to right


  The covalent radius decreases
• As you move from left to right across the periodic table, atoms have more electrons in their outer energy level and more protons in their nucleus
• As you move down a group in the periodic table


  The covalent radius increases
• Ionisation energy


  The energy involved in removing one mole of electrons from one mole of atoms in the gaseous state
• The second ionisation energy is the energy required to remove a second mole of electrons
• Across a period from left to right

  The ionisation energy increases
• Going down a group

  The ionisation energy decreases
• Electronegativity


  A measure of an atom's attraction for the electrons in a bond
• Across a period from left to right


  The electronegativity of atoms increases
• Going down a group

  Electronegativity decreases
• Both of these trends show that fluorine is highly electronegative (it pulls a shared pair of bonding electrons towards itself)
• Metallic bonding


  Occurs between the atoms of metal elements. The outer electrons are delocalised (free to move). This produces an electrostatic force of attraction between the positive metal ions and the negative delocalised electrons. This delocalised 'sea of electrons' is responsible for metal elements being able to conduct electricity
• Covalent molecules


  Discrete covalent molecules are small groups of atoms held together by strong covalent bonds inside the molecule and weak intermolecular forces between the molecules. The covalent bond itself is a shared pair of electrons electrostatically attracted to the positive nuclei of two non-metal atoms. The atoms achieve a stable outer electron arrangement (a noble gas arrangement) by sharing electrons
• Most of the discrete covalent molecules are diatomic elements


  • Hydrogen (H2)
  • Nitrogen (N2)
  • Oxygen (O2)
  • Fluorine (F2)
  • Chlorine (Cl2)

  
• There are also some larger covalent molecular elements


  • Phosphorous (P4)
  • Sulfur (S8)
  • Fullerenes (C60)
• Covalent network


  Covalent networks are large, rigid three-dimensional arrangements of atoms held together by strong covalent bonds. They have high melting points because they only contain strong bonds. Examples include carbon in the forms of diamond and graphite
• Monatomic elements


  Group 0 elements (the noble gases) exist as single, unattached particles. They are stable atoms. They have fuller outer energy levels so they do not usually form molecules with other atoms. They have low melting and boiling points as they are easily separated by overcoming the weak forces of attraction between the atoms
• Summary of bonding


  • Monatomic (noble gas)
  • Covalent molecular (N2, P4, S8)
  • Covalent network
  • Metallic lattice (all metals)

  
• Monatomic (noble gas): Low density, low melting point, non-conductor
• Covalent molecular (N2, P4, S8): Low density, low melting point, non-conductor
• Covalent network: Very high density, very high melting point, non-conductor (except graphite)
• Metallic lattice (all metals): High density, high melting point, conductor
• Intramolecular bonding


  Bonding inside molecules
• Intermolecular bonding


  Bonding between molecules
• Pure covalent bond


  A shared pair of electrons, electrostatically attracted to the positive nuclei of two atoms. Atoms can share electrons to achieve a stable outer electron arrangement (a noble gas arrangement). Pure covalent bonds exist between two atoms with the same electronegativities and have no ionic character
• Polar covalent bond


  A bond formed when a shared pair of electrons is not shared equally, due to one of the elements having a higher electronegativity than the other. This makes one atom slightly negative and the other slightly positive, creating a dipole
• Ionic bond


  Formed between a metal and non-metal with a large difference in electronegativity. The atom with the lower electronegativity loses an electron to the atom with the higher electronegativity, forming positive and negative ions. The ionic bond is the electrostatic force of attraction between these ions, arranged in a 3D ionic lattice
• Pure covalent bonds, polar covalent bonds and ionic bonds all exist as part of the same bonding continuum, determined by the differences in electronegativity between the elements involved
• London dispersion forces


  The weakest type of intermolecular bond, caused by an uneven distribution of electrons within an atom creating a temporary dipole. The strength depends on the size of the molecule or atom
• Permanent dipole interactions


  Stronger than London dispersion forces, occur between polar molecules with a permanent dipole
• Hydrogen bonding


  The strongest type of intermolecular bond, a specific type of permanent dipole to permanent dipole attraction that occurs when a hydrogen atom is covalently bonded to a highly electronegative element such as nitrogen, oxygen or fluorine
• Permanent dipole interactions


  Molecules with a permanent dipole are polar. Polar molecules display attractions between the oppositely charged ends of the molecules. This type of intermolecular bond is stronger than London dispersion forces (which can be called temporary dipoles).
• Hydrogen bonding


  Hydrogen bonding is the strongest type of intermolecular bond. It is a specific type of permanent dipole to permanent dipole attraction that occurs when a hydrogen atom is covalently bonded to a highly electronegative element such as nitrogen, oxygen or fluorine. As with permanent dipole to permanent dipole attractions, the oppositely charged ends of molecules attract.
• Molecules containing hydrogen bonding

  • Water
  • Ammonia
  • Alcohols
  • Alkanoic acids
• Bonding strength


  • Covalent bonds > Hydrogen bonds > Permanent dipole interactions > London dispersion forces
• Bromine molecules


  Contain pure covalent bonds so are held together by London dispersion forces (temporary dipoles caused by momentary uneven sharing of electrons). Bromine is a liquid at room temperature (melting point -7°C).

  
• Iodine monochloride (I-Cl)


  Contains a polar covalent bond. Since this molecule is a dipole it displays permanent dipole to permanent dipole interactions. It exists as a solid at room temperature (melting point 27°C).
• Polar molecules


  Molecules that contain polar covalent bonds and have a non-symmetrical shape, resulting in an overall polarity of the molecule.
• Non-polar molecules


  Molecules that contain polar covalent bonds but have a symmetrical shape, resulting in no overall polarity of the molecule.
• Polar molecules


  • Water