Chemistry

• Chemical bond


  A mutual attraction between two atoms resulting from the simultaneous attraction between their nuclei and the outer electrons

  
• The energy of the combined atoms is lower than that of the individual atoms resulting in higher stability
• Main types of chemical bonds


  • Ionic Bonds
  • Metallic Bonds
  • Covalent Bonds

  
• Bond length


  The average distance between the nuclei of two bonded atoms
• The stronger the bond


  The shorter the bond length
• The greater the number of bonds (single, double, triple) between bonded atoms


  The greater the bond strength (and therefore the shorter the bond length)
• The larger the bonded atoms


  The longer the bond length (and therefore the weaker the bond strength)
• Bond energy


  The energy needed to break one mole of its molecules into separate atoms
• The shorter and stronger the bond


  The more energy required to break the bond
• Ionic bond


  A bond formed when there is a transfer of electrons to form ions which attract each other with an electrostatic force
• An ionic bond forms between a metal and a non-metal atom
• Metallic bond


  A bond formed due to the force of attraction between positive nuclei and a sea of delocalized valence electrons

  
• Metallic bonds form between two metal atoms
• Covalent bond


  A bond formed by the overlapping of half-filled orbitals in non-metals resulting in the sharing of electrons

  
• A covalent bond forms between two non-metal atoms
• Molecule


  A group of two or more atoms covalently bonded and that function as a unit
• Valence electrons


  The electrons in the highest energy level of an atom in which there are electrons
• Lewis (dot) diagram


  A structural formula in which valence electrons are represented by dots (or crosses) (1 dot/cross = 1 valence electron)

  
• Rule 1: Different atoms, each with an unpaired valence electron, can share these electrons to form a chemical bond

  
• Bond order


  The number of unpaired valence electrons in an atom, equal to the number of bonds the atom can form

  
• Multiple bonds between atoms is possible
• Bonding pair


  A pair of electrons that is shared between two atoms in a covalent bond
• Lone pair


  A pair of electrons in the valence shell of an atom that is not shared with another atom

  
• Couper notation


  A Lewis diagram which shows only the bonding electron pairs using a line (1 line = 1 bonding pair)

  
• Electronegativity


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

  
• The difference in electronegativity (ΔEN) between two atoms can give an indication of the type of bond that will form

  
• Non-polar covalent bonds


  Bonds that form between (non-metal) atoms of the same element OR between atoms with the same electronegativity

  
• In non-polar bonds, bonding electrons are shared equally between the atoms

  
• Interatomic forces


  Forces that exist between atoms within a molecule
• Polar covalent bonds


  Bonds that form between (non-metal) atoms of different elements AND different electronegativity

  
• In polar bonds, bonding electrons are shared unequally between atoms with electrons being pulled closer to the atom with a higher electronegativity

  
• Dipole


  The formation of a partial negative charge (δ-) on the atom with the stronger pull and a partial positive charge (δ+) on the atom with the weaker pull
• Intermolecular forces (IMFs)


  Forces that exist between (atoms in different) molecules
• Rule 2: Different atoms with paired valence electrons, called lone pairs, cannot share these four electrons and cannot form a chemical bond
• Interatomic forces


  • Ionic bonds
  • Metallic bonds
  • Covalent bonds
• Rule 3: Atoms with an empty valence shell can share a lone pair of electrons from another atom to form a dative covalent bond (aka a coordinate covalent bond)
• Intermolecular forces


  • Ion-dipole forces
  • Dipole-dipole forces
  • Hydrogen bonds
  • Dipole-induced dipole forces
  • London forces
• VSEPR theory


  A theory used to predict the 3D shape of a molecule

  
• Relative Strengths of Forces


  • Ionic Bonds (600 − 1 100 kJ.mol-1)
  • Covalent Bonds (60 − 700 kJ.mol-1)
  • Metallic Bonds (100 − 350 kJ.mol-1)
  • Hydrogen Bonds (10 − 40 kJ.mol-1)
  • Dipole-Dipole Forces (5 − 25 kJ.mol-1)
  • London Forces (0.05 − 40 kJ.mol-1)
• According to VSPER theory, the electron pairs (bonding and lone) will orientate themselves to minimise the forces of repulsion between them