CHem

• Chemical bonding


  Formation of bonds between atoms, molecules, or ions
• Chemical bonds


  Attractive forces which hold the atoms, molecules, or ions together in the resulting compound
• Ions


  Atoms or group of atoms that has an electric charge
• Two types of ions

  • Cations (net positive charge: more protons than electrons)
  • Anions (net negative charge: more electrons than protons)
• Groups


  A column in the periodic table. The elements in each group have the same number of valence electrons (similar chemical properties) determined by the outermost electrons.
• Period


  The horizontal rows in the periodic table. They represent elements having the same number of electron shells or energy levels. They are arranged according to the increasing atomic number of the elements.
• Valence


  A property of element which determines the number of other atoms with which an atom of the element can combine.
• Valence electron


  Negatively charged particles in the outermost region of atoms that undergo formation of chemical bonds.
• Groups in the periodic table

  • Group 1 (Alkali metals, 1 valence electron)
  • Group 2 (Alkaline earth metals, 2 valence electrons)
  • Groups 3-12 (Transition metals, d and f block metals have two valence electrons)
  • Group 13 (Boron group or earth metals, 3 valence electrons)
  • Group 14 (Carbon group or tetrels, 4 valence electrons)
  • Group 15 (Nitrogen group or Pnictogens, 5 valence electrons)
• Antoine Lavoisier tried grouping the elements as metals and nonmetals


  1789
• Johann Wolfang Döbereiner arranged the elements in groups of three in increasing order of atomic weight and called them triads


  1829
• Revision of list of elements and their atomic masses at the First International Conference of Chemistry in Karlsruhe, Germany

  1860
• John Newlands arranged the elements in the periodic table with increasing order of atomic mass (Law of Octaves)


  1865
• Newland's Law of Octaves

  • Elements are arranged in increasing order of Atomic Mass
  • States that the properties of every eighth element starting from any element are a repetition of the properties of the starting element
  • True only for elements up to Calcium
• Dmitri Mendeleev created a framework that became the modern periodic table, leaving gaps for elements that were not yet discovered


  1869
• Lother Meyer produced a version of periodic table similar to Mendeleev's, leaving gaps for some elements to be discovered


  1870

  
• The Royal Society of London awarded Mendeleev and Meyer the Davy Medal


  1882
• Henry Moseley used X-rays to measure the wavelengths of elements and correlated these measurements to their atomic numbers


  1913
• Octet rule


  States that an atom will be most stable when surrounded by 8 electrons in the outermost shell, achieved by gaining or losing electrons
• Octet rule


  • Helium and hydrogen both do not follow the octet rule, they just need two electrons in the valence shell to be stable
  • Transition metals do not normally follow the rule, because they have d-orbitals and can have more electrons in the valence shell
  • Lithium often will lose an electron to have the same configuration as helium
  • Some compounds disobey the octet rule, the further down an atom is in the periodic table, the more likely it is to not have eight electrons in the valence shell and still be stable (d-f)
• Exceptions to the octet rule

  • Sodium Chloride (NaCl) - Ionic
  • Hydrochloric acid (HCl) - Covalent
• Lewis structure


  A very simplified representation of the valence shell electrons in a molecule, with electrons represented by dots and bonding electrons represented by lines between atoms
• Pattern for Lewis structure

  • Fill up the sides ONE AT A TIME
  • Take note of Hund's Rule in distributing the valence electrons
• Lewis structure for Iodine (I)

  • Atomic number: 53

  Electron configuration: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 -4p6 5s2 4d10 5p5

  Outermost shell: 5th

  Valence electrons: 7
• Lewis structure for Chlorine (Cl)

  • Atomic number: 17

  Electron configuration: 1s2 2s2 2p6 3s2 3p5

  Outermost shell: 3rd

  Valence electrons: 7
• Lewis structure for Argon (Ar)

  • Atomic number: 18

  Electron configuration: 1s2 2s2 2p6 3s2 3p6

  Outermost shell: 3rd

  Valence electrons: 8
• Ionic bonding


  Formed due to the attraction between ions (+ and -), commonly between a metal (loss of electrons) and a non-metal (gain of electrons)
• Examples of ionic bonding


  • NaCl (Sodium Chloride)

  NaBr (Sodium Bromide)

  CaCl (Calcium Chloride)

  CaO (Calcium Oxide)
• Covalent bonding


  Involves sharing of electrons because both elements have high electronegativities, usually between two non-metals
• Types of covalent bonding

  • Non-polar - equal sharing of electrons

  Polar - unequal sharing of electrons
• Covalent bonding in diatomic molecules

  • H2, O2, F2 Br2, I2, N2, Cl2
• Examples of covalent bonding

  • H2 (Hydrogen)

  O2 (Oxygen)

  N2 (Nitrogen)

  CH4 (Methane)
• Metallic bonding


  Bond that occurs between metallic elements, formed by the valence electrons moving freely through the metal lattice
• Examples of metallic bonding

  • Au (Gold)

  Ag (Silver)

  Al (Aluminum)

  Na (Sodium)