The Alkali Metals – Li, Na, K, Rb, Cs (Fr is unstable and scarce) -Low density largest atoms in each period with lowest mass in each period -Soft weak metallic -Malleable bonding b/c only -Low melting points one valence e - -Melting points decrease as you move down the group since atom size increases as you move down and the attraction between the valence e - and the nucleus decreases

Physical Properties -Properties that can be observed without changing the identity of the substance -Melting points, boiling point, density, softness, malleability, electronegativity, ionization energy, and atomic radii Chemical Properties Properties that describe how a substance changes into a completely different substance are called chemical properties. Flammability, corrosion/oxidation resistance, acidity, reactivity are examples of chemical properties.

All alkali metals are very reactive -They only have to lose one electron that is not held strongly by the nucleus -Reactivity increases as you move down a group since the valence electron is easier to lose -Tarnish when exposed to air -Form oxides with the general formula M 2 O Example: 4Na (s) + O 2(g)  2Na 2 O (s) -form halides with the general formula MX Example: 2 Li (s) + Cl 2(g)  2 LiCl (s)

All alkali metals will react vigorously with water to produce hydrogen gas and a alkaline solution (basic solution) Example: 2 Li (s) + 2H 2 O (l)  2LiOH (aq) + H 2(g) Li + + OH -

Halogens – F, Cl, Br and I (At is scarce & unstable) -very reactive  only needs one electron to complete valence shell -Reactivity decreases as you go down the group since the electron gained is at a higher energy level that is farther away from the nucleus and therefore not as strongly attracted. -All exist as diatomic molecules -All highly electronegative -State at room temperature: -Cl 2 and F 2 are gases -Br 2 is a liquid at room temp. -I 2 is a solid (changes to a gas when heated) - the different states are due to the increase atom size which results in an increase in london dispersion forces as you move down the group

Halogens are not very soluble in water Solutions of chlorine have a green tinge Solutions of bromine vary from yellow to orange to brown as concentration increases Iodine is brown water or ethanol (polar solvents) Iodine is purple in paint thinner or hexane (non-polar solvents)

Halogens form acidic solution when dissolved in water: X 2(g) + H 2 O  HOX (aq) + X - (aq) + H + (aq) Example: Br 2(l) + H 2 O  HOBr (aq) + Br - (aq) + H + (aq) weak acid Cl 2(g) + H 2 O  HOCl (aq) + Cl - (aq) + H + (aq) weak acid- turns litmus red and then colourless NaOCl - used in bleach and disinfectants (kills microbes)

Reaction with Metals -Forms salts that are white and usually soluble in water giving clear solutions PbI 2 is not soluble and bright yellow AgCl  white ppt. that darkens when exposed to light AgBr  cream ppt. AgI  pale yellow ppt. AgF  soluble

Since reactivity decreases as you move down a group, a higher halogen will displace a lower halogen in a reaction Example: Cl 2(aq) + 2 I - (aq)  I 2(aq ) + 2 CI - (aq) I 2(aq ) + 2 CI - (aq)  no reaction

Metal oxides tend to form basic solution when dissolved in water Non-metal oxides tend to form acidic solutions when dissolved in water Oxides of Period Three The oxides we'll be looking at are: Na 2 O MgO Al 2 O 3 SiO 2 P 4 O 10 SO 3 Cl 2 O 7 P 4 O 6 SO 2 Cl 2 O ClO 2 Notice in the first row of oxides that each successive element bonds to an extra half oxygen NOTE: to the right of the period more than one possible oxide exists A simple summary of the trend in acid-base behaviour The trend is from strongly basic oxides on the left-hand side to strongly acidic ones on the right, via an amphoteric oxide (aluminium oxide) in the middle. An amphoteric oxide is one which shows both acidic and basic properties.

Sodium oxide Na 2 O is an ionic compound and is therefore a solid with high melting and boiling point. Molten Na 2 O can conduct electricity. Sodium oxide is a strongly basic oxide. It is basic because it contains the oxide ion, O 2-, which is a very strong base with a high tendency to combine with hydrogen ions. Reaction with water Sodium oxide reacts exothermically with cold water to produce sodium hydroxide solution. Depending on its concentration, this will have a pH around 14. Na 2 O (s) + H 2 O (l)  2Na + (aq) + 2 OH - (aq) Reaction with acids As a strong base, sodium oxide also reacts with acids. Na 2 O (s) + 2HCl (aq)  2Na + (aq) + 2 Cl - (aq) + H 2 O (l)

Magnesium oxide Magnesium oxide is again a basic oxide, because it also contains oxide ions. However, it isn't as strongly basic as sodium oxide because the oxide ions aren't so free. In the sodium oxide case, the solid is held together by attractions between 1+ and 2- ions. In the magnesium oxide case, the attractions are between 2+ and 2-. It takes more energy to break these. Therefore reactions involving magnesium oxide will always be less exothermic than those of sodium oxide. MgO is an ionic compound as is therefore a solid with high m.p and b.p Molten MgO can conduct electricity. Reaction with water When MgO is dissolved in water a slightly alkaline solution is produced (around pH = 9) There is a slight reaction with the water to produce magnesium hydroxide. However this is almost insoluble - and so not many hydroxide ions actually get into solution. MgO(s) + H 2 O(l)  2Mg(OH) 2(s) Reaction with acids MgO(s) + 2HCl(aq)  2Mg +2 (aq) + 2 Cl - (aq) + H 2 O(l)

Aluminium oxide Has both ionic and covalent characteristics. It has a very high m.p and can conduct electricity in a liquid state. Aluminium oxide is amphoteric. It has reactions as both a base and an acid. Reaction with water The oxide ions are held too strongly in the solid lattice to react with the water. Reaction with acids Al 2 O 3(s) + 6H + (aq)  2Al 3+ (aq) + 3H 2 O Reaction with bases Al 2 O 3(s) + 2 OH - (aq) + 3H 2 O (l)  2Al(OH) - 4(aq)

Silicon dioxide Giant lattice structure held together by strong covalent bonds. SiO 2 is a solid with high melting and boiling points. The difference in electronegativity between silicon and oxygen is too small to form ionic bonds. Silicon dioxide has no basic properties - it doesn't contain oxide ions and it doesn't react with acids. Instead, it is very weakly acidic, reacting with hot concentrated or molten bases. SiO 2(s) + 2 NaOH (l)  Na 2 SiO 3(l) + H 2 O(g) Reaction with water Silicon dioxide doesn't react with water, because of the difficulty of breaking up the giant covalent structure.

Phosphorus Oxides White solids with low melting points Reaction with water: P 4 O 6 + 6H 2 O  4H 3 PO 3  4H + + 4H 2 PO 3 – P 4 O H 2 O  4H 3 PO 4  4H + + 4H 2 PO 4 – H 3 PO 3 (phophorous acid) is a weak acid and H 3 PO 4 (phosphoric acid) is a stronger acid Reactions with bases: P 4 O OH -  4PO 4 –3 + 6H 2 O P 4 O OH -  4HPO 3 –2 + 2H 2 O

Sulphur Oxides Gases at room temperature Reactions with water: SO 2 + H 2 O  H 2 SO 3  H + + HSO 3 – SO 3 + H 2 O  H 2 SO 4  H + + HSO 4 – H 2 SO 3 (sulphurous acid) is a weak acid H 2 SO 4 (sulphuric acid) is a strong acid Reactions with bases: SO OH -  SO H 2 O SO 3 + 2OH -  SO H 2 O

Chlorine Oxides Cl 2 O is a gas at room temperature and Cl 2 O 7 is a liquid. Reactions with water Cl 2 O 7 + H 2 O  2HClO 4 HClO 4 (perchloric acid) is a strong acid HClO 4  H + + ClO 4 – Cl 2 O + H 2 O  2HOCl (weak acid) Reactions with bases Cl 2 O OH -  2ClO 4 – + H 2 O Cl 2 O + 2 OH -  2OCl – + H 2 O

Chlorides of Period 3 NaCl- ionic compounds with high melting and boiling points -NaCl dissociates when dissolved in water with no reaction -NaCl  Na + + Cl - MgCl 2 - ionic compound with high melting and boiling points -MgCl 2 will also dissociate (MgCl 2  Mg Cl - ), -a very small percentage of MgCl 2 will react with water to form a slightly acidic solution : MgCl 2 + 2H 2 O  Mg(OH) 2(s) + 2H + + 2Cl - -Solutions of NaCl and MgCl 2 and molten NaCl and MgCl 2 will conduct electricity

AlCl 3 – solid that sublimes at 178 o C to form Al 2 Cl 6 -reacts vigorously when added to water to produce an acidic solution that will conduct electricity AlCl 3(s) + 3H 2 O  Al(OH) 3(s) + 3H + + 3Cl - If there is a large excess of water the following reaction occurs: AlCl 3(s) + 6H 2 O  [Al(H 2 O) 6 ] 3+ (aq) + 3Cl - [Al(H 2 O) 6 ] + (aq)  [Al(H 2 O) 5 OH] 2+ (aq) + H + (weak acid) -Molten AlCl 3 will not conduct electricity

Can not form giant covalent lattice structures for chlorides of silicon and phosphorus because chlorine can only form one bond. Chlorides of silicon and phosphorus have molecular covalent structures. This results in weak intermolecular forces, and therefore low melting and boiling points. Molten silicon and phosphorus chlorides do not conduct electricity

Reactions with water -produces acidic solution -produces solutions that conduct electricity SiCl 4(l) + 2 H 2 O  SiO 2(s) + 4H + + 4Cl - PCl 3(l) + 3H 2 O  H 3 PO 3(aq) + 3H + + 3Cl - PCl 5(l) + 4H 2 O  H 3 PO 4(aq) + 5H + + 5Cl - Hydrolysis: reaction with water where water splits up a molecule Note: as you move across period 3 each successive element bonds to an additional chlorine atom