Chapter 11 Groups II and VII

11.1 Physical Properties of Group II Elements The elements of group 2 are referred to as the ALKALINE EARTH METALS Electronic configuration ends with 2 electrons in their outermost shell (s subshell). Metallic radius increases down the group

Melting points decrease down group Atomic number increases down group Density decreases down group

11.2 Reactions of Group II Elements Group 2 Metals form Ionic compounds. When they react, they loose the outer 2 electrons in their sub-shell. Mg Mg^2+ 2e-

Ionization Energies of Group II Elements Decreases down the group

Magnesium Powder is used in flares! Reactions With Oxygen These metals burn in air, forming white solid oxides. Example: 2Mg(s) + O2(g)  2MgO(s) Some group II metals burn with flame colors; this is caused by the 2+ ions formed in the reaction. Calcium -> Brick Red Color Strontium -> Scarlet Red Color Barium -> Apple Green Color Magnesium Powder is used in flares!

Reactions With Water Magnesium reacts slowly with cold water but eventually forms a weakly alkaline solution. Mg(s) + 2H20(l) -> Mh(OH)2 + H2(g) Hot Magnesium Reacts vigorously with water Mg(s) + H2O(g) -> MgO(s) +H2(g) Calcium reacts more readily with water Ca(s) + 2H20(l) -> Ca(OH)2(aq) +H2(g) Going down the group hydrogen is released more rapidly by the reaction of the element and water. Solutions get more alkaline.

11.3 Thermal Composition of Group II Carbonates and Nitrates Carbonates and nitrates decompose when heated. Carbonates break down and form the metal oxide and give off carbon dioxide gas. MgCO3(s) -->(heat) MgO(s) +CO2(g) The temperature at which thermal decomposition takes place increases going down group 2 for both carbonates and nitrates. Nitrates also undergo thermal decomposition: 2Ca(NO3)2(s) (heat) 2CaO(s) + 4NO2(g) +O2

11.4 Some Uses of Group II Compounds Limestone is made of calcium carbonate Marble Used to make cement Magnesium oxide can be used to neutralize excess acid in the stomach.

11.5 Physical Properties of Group VII Elements Called the Halogens Atoms have 7 electrons in their outermost shell. Are all non-metals At room temperature they exist as diatomic molecules. There is a single covalent bong between the two atoms in each molecule. Melting and boiling points increase going down the group. Values are relatively low because of the simple molecular structures and weak van der Waals forces. The colors of the halogens get darker down the group.

11.6 Reactions of Group VII Elements Halogens react with metallic elements because they only need one more electron to complete their outer shell. Ca(s) + Cl2(g)  CaCl2(s) The halogens are oxidizing agents. Cl2  2Cl- +2e- Halogens also react with many non-metals forming covalent bonds H2(g) + Cl(g)  2HCl(g) The halogens get less reactive going down group VII The halogens get less powerful as oxidizing agents down group VII. FLOURINE IS THE MOST ELECTRONEGATIVE OF ALL THE ELEMENTS

(cont.) Displacement Reactions We can judge the reactivity of the halogens by looking at their displacement reactions with other halide ions in solution. A more reactive halogen can displace a less reactive halogen from a halide solution of the less reactive halogen. Chlorine water is added to sodium bromide Cl2(aq)+2NaBr(aq)  2NaCl(aq)+Br2(aq) Chlorine has “displaced” bromine from the solution Summarized ionic equation: Cl2(aq) +2Br(aq)  2Cl(aq) +Br(aq) The colors in these reactions are difficult to identify so they add cyclohexane which gives the reactions different colors.

(cont.) Reactions With Hydrogen Halogens from hydrogen halides with hydrogen gas. Hydrogen Halides are less thermally stable down the group. Bond Energies H-F: 562 H-Cl: 431 H-Br: 366 H-I: 299 The longer the bond, the weaker it is and the less energy required to break it.

11.7 Reactions with Halide Ions We tell the halide ions apart by using chemical tests. If an unknown compound is dissolved in dilute nitric acid and silver nitrate solution is added, a precipitate will be formed. They add ammonia solution to verify the results. AgNO3(aq) + X(aq)  AgX(s) + NO3(aq) Reactions of halide ions with sulfuric acid: NaCl(s) + H2SO4(l)  NaHSO4(s) + HCl(g) NaBr(s) +H2SO4(l)  NaHSO4(s) + HBr(g)

11.8 Disproportionation This term can be thought of as a “self reduction reaction.” When chlorine reacts with dilute alkali some chlorine atoms are reduced and some are oxidized. Chlorine in cold aklali: Cl2(aq) + 2NaOH(aq)  NaCl(aq) +NaClO(aq) + H20(l) Chlorine in hot alkali: 3Cl2(aq) + 6NaOH(aq)  5NaCl(aq) + NaClO3(aq) + 3H20(l)

11.9 Uses of the Halogens and Their Compounds Adding chlorine to water kills bacteria and makes water safer to drink. Bleach (NaCl) Plastic PVC Halogenated hydrocarbons used as solvents in refrigerants and aerosols.