module three: chapter seven: periodicity

• reading from left to right of the periodic table, the elements are arranged in order of increasing atomic number.
• the elements on a periodic table are arranged in horizontal rows called periods, the number of the period gives the number of the highest energy electron shell in an elements atoms.
• across period two, the 2s sub shell fills with two electrons, followed by the 2p sub shell with six electrons
• what is the name of group one on the periodic table?

  alkali metals
• what is the name of group two on the periodic table?

  alkaline earth metals
• what is the name of group three to twelve on the periodic table?

  transition elements
• what is the name of group five on the periodic table?

  pnictogens
• what is the name of group six on the periodic table?

  chalcogens
• what is the name of group seven on the periodic table?

  halogens
• what is the name of group zero on the periodic table?

  noble gases
• ionisation energy measures how easily an atom loses electrons to form positive ions.
• first ionisation energy is the energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.
• first ionisation energy example:

  Na (g) = Na+ (g) + e-
• electrons are held in their shells by attraction from the nucleus. the first electron lost will be in the highest energy level and will experience the least attraction from the nucleus.
• atomic radius is the greater the distance between the nucleus and the outer electrons, the less the nuclear attraction. the force of attraction falls off sharply with increasing distance, so atomic radius has a large effect.
• nuclear charge is the more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons.
• electron shielding is when electrons are negatively charged and so inner shell electrons repel outer shell electrons, this repulsion, called the shielding effect, reduces the attraction between the nucleus and the outer electrons.
• an element has as many ionisation energies as there are electrons, for example, helium has two electrons and two ionisation energies.
• second ionisation energy example:

  He+ (g) = He2+ (g) + e-
• the second ionisation energy is the energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions.
• successive ionisation energies allow predictions to be made about: the number of electrons in the outer shell, the group of the element, the identity of an element.
• as you go down the noble gases, atomic radius increases, more inner shells so shielding increases, nuclear attraction on outer electrons decreases, and first ionisation energy decreases.
• as you go across period two, nuclear charge increases, same shell so similar shielding, nuclear attraction increases, atomic radius decreases, and first ionisation energy increases.
• at room temperature all metals apart from mercury are solid. all metals can conduct electricity.
• metallic bonding and structure: in a solid metal structure, each atom has donated it's negative outer shell electrons to a shared pool of electrons which are delocalised throughout the whole structure. the positive ions left behind consist of the nucleus and the inner electron shells of the metal atoms.
• metallic bonding is the strongest electrostatic attraction between cations and delocalised electrons. the cations are fixed in position maintaining the structure and shape of the metal. the delocalised electrons are mobile and are able to move throughout the structure, only the electrons move.
• metallic bonding example: if there are twelve cations, each with a 1+ charge, this balances the charge. for metals containing 2+ cations, twice as many negatively charged electrons are present to balance the charge.
• most metals have strong metallic bonds, high electrical conductivity and high melting and boiling points
• metals conduct electricity in solid and liquid. the delocalised electrons can move through the structure carrying charge. however in ionic compounds they have no mobile charge carriers as a solid.
• the melting point depends on the strength of the metallic bonds holding together the atoms in the giant metallic lattice
• for most metals, high temperatures are necessary to provide large amounts of energy needed to overcome the strong electrostatic attraction between cations and electrons. the strong attraction results in most metals having high melting and boiling points.
• carbon and silicon are in group four, carbon and silicon use their four electrons to form covalent bonds to other carbon or silicon atoms. the result is a tetrahedral structure.
• giant covalent lattices have high melting and boiling points. this is because covalent bonds are strong, high temperatures are necessary to provide the large quantity of energy needed to break the strong covalent bonds.
• giant covalent lattices are insoluble in almost all solvents, the covalent bonds holding together the atoms in the lattice are far too strong to be broken by interaction with solvents.
• giant covalent lattices are non conductors of electricity, only exceptions are graphene and graphite
• in carbon and silicon, all four outer shell electrons are involved in covalent bonding, so none are available for conducting electricity
• carbon is special in forming several structures in which one of the electrons is available for conductivity. graphene and graphite are able to conduct electricity
• the oxides of group two elements react with water, releasing hydroxide ions, OH- and forming alkaline solutions of the metal hydroxide. CaO (s) + H2O (l) = Ca2+ (aq) + 2OH- (aq)
• the group 2 hydroxides are only slightly soluble in water . when the solution becomes saturated, any further metal and hydroxide ions will form a solid precipitate: Ca2+ (aq) + 2OH- (aq) = Ca(OH)2 (s)
• the solubility of hydroxides in water increases down group 2, so the resulting solutions contain more OH- ions and more alkaline.