Group 2

• Group 2 Metals


  Also known as alkaline earth metals. They're all in s-block and have high melting and boiling points due to strong metallic bonding and a giant metallic structure. The strength of the metallic bonding decreases down the group. They lose 2 electrons from the highest energy s orbital to form cations with a 2+ charge. They are oxidised. They are reducing agents because they are oxidised and reduces another species.
• Alkaline earth metals


  Name comes from the alkaline properties of the metal hydroxides. The elements are reactive and do not occur in their elemental form naturally. On earth, they're found in stable compounds such as CaCO3
• Electronic configuration of group 2 elements

  All in s-block. Each group 2 element has two outer shell electrons, two more than the electron configuration of a noble gas. Two electrons are in the outer s sub-shell and are lost first because they are from the highest energy s orbital.
• Oxidation


  Increase in oxidation number, loss of electrons.
• Reduction


  Decrease in oxidation number, gain of electrons.
• Cation Formation


  Group 2 metals form 2+ ions by losing electrons.
• Reactions of group 2 elements and redox reaction trends

  Each new electron added is further from the nucleus. There's a decrease in attraction between outer electrons and nucleus as atomic radius increases and shielding increases. Outer electrons are lost more easily as you move down the group. Reactivity increases down the group as it becomes easier to lose electrons down the group.
• General redox reaction for a group 2 element

  M -> M2+ + 2e-

  Oxidation number of M- (0)

  Oxidation number of M2+- (+2)

  M is oxidised as it loses electrons.
• Atomic Radius


  Increases down Group 2, affecting reactivity.
• Ionisation Energy


  The energy needed to remove and electron from the highest energy sub-shell to form positive ions.
• First Ionisation Energy


  The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.
• General equation for first ionisation energy

  M(g) -> M+(g) + e-
• Second Ionisation Energy


  The energy required to remove one electron from each ion in one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions.
• General equation for second ionisation energy

  M(g) -> M2+(g) + e-
• Reasons for greater ionisation energies than other group 2 elements

  Less shielding

  Greater attraction between electrons and the nucleus

  More energy is required to remove its outermost electron

  Atomic radius is smaller
• Trend in ionisation energies

  Ionisation energies decrease down group 2 due to increasing atomic radius meaning more shielding between the nucleus and the outermost electron resulting in a weaker attraction between the two.
• Reactivity Trend


  Increases down Group 2 due to lower ionisation energy. As the nuclear attraction on the outer electrons decreases. This is due to increased atomic radius and increased shielding. It becomes easier to remove electrons as reactivity increases as it's easier to give away electrons.
• Shielding Effect


  Increased shielding meaning the atomic radius increases, so there's a weaker attraction between outermost electrons and the nucleus.
• Redox Reactions


  Involve electron transfer; oxidation and reduction occur.
• Group 2 elements with oxygen

  Vigorous and exothermic reactions, they release energy in the form of heat. The product is a basic oxide. The reactions become more violent going down the group as the reactivity of the metals increases. The group 2 metal oxides form, MO, and have a giant ionic lattice structure and high melting points.
• General equation when group 2 elements react with oxygen

  MO = M2+ + O2-
• Basic oxide


  A base that reacts with an acid. They're proton acceptors and don't dissolve in water.
• Group 2 elements in water

  Group 2 elements react with water to form alkaline metal hydroxides (M(OH)2 and release H2 gas. The M2+ ions and hydroxide OH- ions dissolve in water forming aqueous ions. These release energy in the form of heat, exothermic reaction.
• Magnesium anomaly


  The reaction between magnesium and liquid water is slow but the reactions of the other group 2 metals become more vigorous down the group as reactivity increases. Magnesium is the lease reacts. and so only reacts rapidly with steam.
• Equation for magnesium hydroxide

  Mg(s) + H2O(g) -> MgO(s) + H2(g)
• Beryllium Anomalies


  Beryllium's reactions differ from other Group 2 elements.
• Stable Compounds


  Group 2 elements found in stable compounds like CaCO3.
• Electron Configuration


  Arrangement of electrons in an atom's orbitals.
• Trend in Electronegativity


  Electronegativity decreases down Group 2.
• Uses of calcium hydroxide

  Ca(OH)2, lime water which is used to prove presence of CO2 solution which turns milky. Ca(OH)2(aq) + CO2(g) -> CaCO3(s) + H2O(l)
• Metallic Bond Strength


  Decreases down the group due to larger atomic radius.
• Ionisation Energy Values


  Measured in kJ/mol for Group 2 elements.
• Group 2 elements with dilute acids

  Group 2 elements react with dilute acids to form a salt and hydrogen. This is a redox reaction and not neutralisation because water is not formed. The reactions become more violent down the group as the reactivity of the metal increases.
• General equation for group 2 elements reacting with dilute acids

  M(s) + H2SO4(aq) -> MSO4(s) + H2(g)
• Ionic Radius


  Size of an ion compared to its atom.
• Giant Ionic Lattice


  Structure of metal oxides with high melting points.
• Redox Reaction


  Simultaneous oxidation and reduction processes.
• Calcium Oxide


  CaO, a basic oxide formed from calcium.
• Strontium Oxide


  SrO, a basic oxide formed from strontium.
• Barium Oxide


  BaO, a basic oxide formed from barium.