Important definitions Mod 2

• Isotopes


  Atoms of the same element with the same number of proton/electrons but different number of neutrons
• Relative atomic mass


  Weighted mean mass of an atom of an element relative to 1/12th of the mass of an atom of carbon-12
• Relative isotopic mass


  Mass of an isotope relative to 1/12th of the mass of an atom of cabon-12 and it is always a whole number.
• Empirical Formula


  Simplest whole number ratio of atoms of each element in a compound
• Water of crystallisation 


  Water molecules that are bonded into a crystalline structure of a compound
• Anhydrous


  Contains no water
• Standard solution

  Solution of known concentration
• Yield


  Amount of product that a chemical reaction produces
• Strong acid 


  Completely dissociates in an aqueous solution, releases all its H+ ions

  eg. HCl

  
  

  Bronsted Lowry acid -> Proton donor
• Weak acid 


  Partially dissociates in aqueous solution

  Eg. Carboxylic acids -> CH3COOH
• Base 


  Neutralises an acid to form a salt

  Bronsted Lowry -> Proton acceptor

  Eg. Metal oxides, metal hydroxides, metal carbonates and ammonia
• Alkali


  A base that dissolves in water releasing OH- ion into the solution
• Salt


  Product of a reaction where the H+ ions are replaced by a metal ion or ammonium ion
• Homogenous Solution


  Ensuring that the solution is the same concentration everywhere ?
• Oxidation


  Losing electrons -> positive ion

  Gaining oxygen

  Increasing oxidation number

  Lose of hydrogen
• Reduction


  Gaining electrons -> negative ion

  Losing oxygen

  Decreasing oxidation number

  Gain of hydrogen
• Electronegativity


  measure of the tendency of an atom to attract a shared pair of electrons within a covalent bond
• Orbitals


  Region around the nucleus that can hold up to 2 electrons with opposite spins
• Ionic bonding


  Electrostatic attraction between oppositely charged ions
• Covalent Bonding


  Strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atom
• Dative covalent bond


  When an atom uses a lone pair of electrons to form a covalent bond.

  One of the bonded atom gives both the shared pair of electrons in the covalent bond
• Metallic bonding 


  Electrostatic attraction between the positive metal ions and negative electrons
• * Induced dipole dipole interaction


  London dispersion forces

  Weakest intermolecular forces, exists in all types of molecules

  Caused by the uneven distribution of electrons which establishes temporary dipoles in molecule. This induces a temporary dipole on the neighbouring molecules.
• *Permanent Dipole - dipole interactions


  Strongest intermolecular

  Exists between polar molecules
• *Hydrogen bonding


  Type of permanent dipole - dipole interaction found when Hydrogen is bonded to either Oxygen, Nitrogen or Fluorine.

  Caused by the large difference in electronegativity between H and O, N and F.
• *Ranking of intermolecular forces


  Strongest

  1. Hydrogen Bonding
  2. Permanent dipole-dipole interaction
  3. Induced dipole-dipole interaction/London Dispersion Forces

  Weakest
• First Ionisation energy


  Energy required to remove one mole of electrons from one mole of atoms in a gaseous state to produce one mole of gaseous 1+ ions

  Equation: $A\ \left(g\right)\ \rightarrow\ A^+$$\left(g\right)\ +$$\ e^-$

  (State symbols are important)