Inorganic Chemistry

• Electron configuration


  The manner in which electrons fill the orbitals
• Effective nuclear charge (Zeff)

  The electrostatic attraction between the protons in the nucleus and electrons in the outer shells, taking into account shielding and repulsion effects
• Zeff for valence electrons increases as you move along a period, and decreases in the order s > p > d > f
• Atomic and ionic radii


  Measured as metallic radius (half distance between metallic nuclei) or covalent radius (non-metallic nuclei)
• Atoms tend to increase in size from one period to the next in a group, and decrease in size from left to right across a period
• Anions are much larger than their parent atoms, while cations are much smaller than their parent atoms
• Electronegativity


  The ability of an atom in a molecule to attract electrons
• Electronegativity increases along a period and decreases down a group, with non-metals generally having higher electronegativity than metals
• Ionization energy


  The ease with which an electron can be removed from an atom
• The first ionization energy (removal of electron from highest occupied orbital of neutral atom) varies systematically throughout the periodic table, with the smallest (easy to remove) at lower left (near caesium)
• Electronegativity


  Measure of the ability of an atom to attract shared electrons in a chemical bond
• Bonds are predominantly covalent


  If the electronegativity values are high (non metallic elements)
• Bonds are metallic


  If the electronegativity values are low (metallic elements)
• Bonds with intermediate differences in electronegativity

  Give rise to polar bonds with negative end of the bond associated with the element with the higher of the two electronegativity values
• First ionization energy


  • Smallest (easy to remove) at lower left (near caesium) and largest upper right (hard to remove) (near fluorine)
• Ionization energy


  Generally follows similar trends to effective nuclear charge (Zeff) and atomic radii. Atoms with small atomic radii generally have high ionization energies
• Successive ionization energies are always greater than the previous one
• Electron gain enthalpy (ΔH) / electron affinity (Ea)


  The change in enthalpy when a gas phase atom gains an electron, the electron entering the highest unoccupied molecular orbital
• Electron affinities can be positive (energy released) or negative (energy absorbed) but are generally positive
• If the electron affinity of atom A is higher than that of atom B, then the energy change on gaining an electron is more exothermic for A than B
• Elements with high electron affinity


  • The additional electron can enter a shell that has a strong effective nuclear charge, this is the case for the non metallic elements in the top right hand section of the periodic table around F
• Hydrogen


  The most abundant element in the Universe and the fifteenth most abundant on earth where it has a rich chemistry
• Hydrogen forms more compounds than any other element, with water being the most important
• Hydrogen


  • Unique electronic structure (1s1) and lack of an inner shell of electrons (resulting in a very small atomic radius) is responsible for the exceptional properties of this element and its compounds
• Hydrogen


  Resembles the alkali metals as it can lose an electron and be oxidized to H+, and the halogens as it can gain an electron and be reduced to hydride ions (H-) achieving the 1s2 helium noble gas configuration
• H+ is the smallest cation, H- is amongst the largest of the anions
• Hydrogen has a high electronegativity (2.2) resulting from a combination of high ionization energy and high electron affinity
• Hydrogen


  • Forms ionic hydrides with many metals, non-polar covalent bonds with non metallic elements such as carbon, sulphur, and phosphorus which have similar electronegativity values, and polar bonds with the more electronegative non metallic elements such as oxygen, nitrogen, fluorine and chlorine
• s block elements


  Occupy groups 1 and 2 of the periodic table with ns1 and ns2 valence shell configurations respectively
• Group 1 elements (alkali metals)


  • Metallic with low electronegativities, all atoms lose their valence electron to form +1 cations
• Group 2 elements (alkaline earth metals)


  • Metallic with low electronegativities, all atoms lose both valence electrons to form +2 cations
• Atomic radii of s block elements

  Increases as you go down each group, resulting in decreasing charge density on the atom (charge to size ratio) with increasing atomic number
• Diagonally related s block elements

  Have similar charge densities, resulting in similarities in chemical and physical properties
• Biological activity or potential drug action of s block elements is most likely pharmacodynamic in nature, arising from the tendency of the bonding of the compounds in this group to be ionic, and the ease of substitution of one element for another one of similar size and charge
• Group 18 elements (noble gases)


  • Very stable and, except for Xenon, do not form any compounds
• Group 14 elements


  • Carbon is the only non-metal, silicon and germanium are semi-metals while tin and lead are metals
• Group 13 elements


  • Boron is a semi-metal and the remainder are metallic
• f block elements (Lanthanides and Actinides)


  Metallic with very similar chemical properties, arising from the lack of involvement of the f electrons in bonding. They form cations in the +3 oxidation state and are about as reactive as the alkali earth metals
• f block and d block elements

  • Characterized by well defined magnetic and spectroscopic properties, the presence of stable unpaired electron (free radical) configurations, and the existence of many radio isotopes. These properties form the basis for the extensive use of these metals and their compounds as metal drugs and as diagnostic reagents
• d block or transition metal elements

  Metallic looking, malleable, ductile, conduct heat and electricity. Higher densities, melting and boiling points than main group metals. More electronegative than main group metals (like to form covalent bonds). Form compounds in more than one oxidation state