1 Group 17 -The Halogens -The only group containing solids, liquids and gases at room temperature -All exist as diatomic molecules -All are volatile Nonmetals - with corrosive fumes -None occur in nature as a pure element. -All are electronegative; HCl, HBr and HI are strong acids; HF is one of the stronger weak acids Cl 2 Br 2 I2I2 Melting Point Boiling Point State (at 20 °C) Density (at 20 °C) Fluorine-220 °C-188 °CGas g/cm 3 Chlorine-101 °C-34 °CGas g/cm 3 Bromine-7.25 °C58.8 °CLiquid3.123 g/cm 3 Iodine114 °C185 °CSolid4.93 g/cm 3

Electronic Properties - Form monoatomic anions with -1 charge. - 7 valence electrons: [N.G.] ns 2 np 5 - Large negative enthalpy of electronic attraction - Are good oxidizing agents ie. They are good at gaining electrons so that other elements are oxidized in their presence First Ionization Energy (kJ/mol) Enthalpy of Electronic Attraction (kJ/mol) Standard Reduction Potential (V = J/C) Fluorine Chlorine Bromine Iodine

3 Oxidation -Fluorine, chlorine and bromine are oxidizers that are strong enough to oxidize the oxygen in water! When fluorine is bubbled through water, hydrogen fluoride and oxygen gas are produced. -Hypofluorous acid (HOF) is an intermediate: Chlorine and bromine are capable of the same however they need the presence of sunlight: The oxygen produced in these reactions is particularly reactive, so moist chlorine gas is a much stronger oxidizing agent than anhydrous chlorine – as evidenced by the fact that it can bleach dyes like indigo, litmus and malachite green while anhydrous chlorine cannot.

4 Reactions with Metals

5 Reactions with P and S Halogens react with phosphorus (P 4 ) to give PX 3 or PX 5. Halogens react with sulfur (S 8 ) to give S 2 X 2 or SX 6 (depending on the ratio of reactants). Many other sulfur-halogen compounds can be made (including, but not limited to, SX 2, SX 4 and S 2 X 4 ). *See for a videohttp://boyles.sdsmt.edu/pwithcl/reaction_of_white_phosphorus_and.htm of the reaction between phosphorus and chlorine at room temperature.

Disproportionation Reactions Chlorine gas also reacts with hydroxide anions. The products of this reaction depend on the temperature of reaction: Cl OH - → ClO - + Cl - + H 2 O These are disproportionation reactions because one oxidation state of chlorine (Cl 2 ) reacts to give two different oxidation states of chlorine (Cl and an oxoanion). Other such disproportionation reactions continue: 2 ClO − → Cl − + ClO 2 − ClO − + ClO 2 − → Cl − + ClO 3 − Sodium chlorate crystals can be heated to undergo yet another reaction: 4 ClO 3 − → Cl − + 3 ClO 4 −

Oxoanions of the halogens Fluorine, chlorine and bromine only occur naturally as monoatomic ions (fluoride, chloride and bromide). Iodine occurs naturally as iodide, but also in some oxygen-containing anions (“oxoanions”). Chlorine, bromine and iodine are all capable of forming oxoanions in which an atom of halogen is surrounded by oxygen atoms. Because oxygen is more electronegative than all three of these halogens, the electron density is pulled out onto the oxygen atoms, leaving the halogen with a positive formal charge: (Note that oxidation states are assigned similarly to formal charge, but treating every bond as 100% ionic, so that both electrons in the bond go to the more electronegative atom.)

Fluorine can only form one oxoanion. Why? Draw Lewis structures for the four oxoanions of bromine, indicating the molecular geometry of each. Determine the formal charge of bromine Atom in each case. Halogens are electronegative thus they do no tolerate large positive formal charges. As the formal charge increases with the number of oxygen atoms, they increase in oxidizing capacity. Rank the oxoanions above by strength as an oxidizing agent. Oxoanions of the halogens

Oxidation State The oxidation state of an element is its formal charge assuming only ionic bonding Electrons are not shared they are placed on the more electronegative element Example: Determine the Oxidation states of positive charges atoms in: 1) LiCl 2) BeCl 2 3) BCl 3 4) CCl 4 Think of it as determining the charge of the ions in an ionic compound Example: Consider the following carbon compounds. Determine the Oxidation State of C. 1) CH 4 2) CH 3 OH 3) CH 2 O 4) CHOOH 5) CO 2

Rules for Assigning Oxidation States 1) Pure elements have oxidation states of 0 2) Ions have oxidation states that add up to the charge of the ion 3) Hydrogen has an oxidation state of +1 unless bonded to a less electronegative atom. When bound to metals or boron it has an oxidation state of -1. 4) Fluorine has an oxidation state of -1 5) Oxygen has an oxidation state of -2 unless bonded to fluorine or another oxygen. 6) Halogens other than fluorine have oxidation states of -1 unless bonded to oxygen or a more electronegative halogen 7) The rest are determined by the process of elimination, where the oxidation states must add up to the total charge of the molecule or ion.

Oxidation State Determine the oxidation state of all the element in the following molecules: a) H 2 H = 0 b) CO 2 O = -20 = C + 2(-2) c) BF 3 F= -1 0 = B + 3(-1) d) H 2 SO 4 H = 1O = -20 = S +(2(+1) + 4(-2)) = S - 6 e) SO 2 ClF O = -2 F = -1Cl = -1 0 = S +(2(-2) -1 -1) = S - 6 f) IO 2 F 2 _ O = -2 F = -1-1 = I + (2(-2) + 2(-1)) = I - 6 C = 4 B = 3 S = 6 g) HPO 4 2- O = -2H = 1 I = 5 -2 = P + (4(-2) + 1) = P - 7P = 5

Oxidation States of Poly Atomic Ions PO 4 3- SO 4 2- NO 3 - ClO 3 - CO 3 2- NO 2 - PO 3 3- SO 3 2- ClO

Oxoanions of the halogens The name of an oxoanion depends on how many oxygen atoms surround the central atom To name an oxoanion which has had H + added, add hydrogen before The ‘old’ name. Remember that the added proton changes the charge by +1! e.g. HCO 3 - is “hydrogen carbonate” ClNCSP Ohypochlorite (ClO - ) O2O2 chlorite (ClO 2 - ) nitrite (NO 2 - ) O3O3 chlorate (ClO 3 - ) nitrate (NO 3 - ) carbonate (CO 3 2- ) sulfite (SO 3 2- ) phosphite (PO 3 3- ) O4O4 perchlorate (ClO 4 - ) sulfate (SO 4 2- ) phosphate (PO 4 3- )

Oxoacids of the Halogens Note that if enough H + have been added to render the oxoanion neutral, it is no longer an oxoanion. It is an oxoacid! Note that the hydrogens are attached to oxygen in most oxoacids. * ClNCSP Hydrochloric acid (HCl) OHypochlorous acid (HOCl) O2O2 Chlorous acid (HClO 2 ) Nitrous acid (HNO 2 ) O3O3 Chloric acid (HClO 3 ) Nitric acid (HNO 3 ) Carbonic acid (H 2 CO 3 ) Sulfurous acid (H 2 SO 3 ) Phosphorous acid * (H 3 PO 3 ) O4O4 Perchloric acid (HClO 4 ) Sulfuric acid (H 2 SO 4 ) Phosphoric acid (H 3 PO 4 ) * Phosphorous acid is an exception to this rule. Only two of its hydrogens are attached to oxygen.

Oxoacids of the Halogens Draw the Lewis structures of sulfurous acid and sulfuric acid, and indicate the molecular geometry of each. Which of these two acids is stronger? As a general rule, the strength of an oxoacid increases with the number of oxygen atoms. The Pauling’s rules quntify this trend for the strength of oxoacids: For the general formula O p E(OH) q. For multiple protons (i.e. q>1), pK a increases by ~5 every time H + is removed. Exercise) Estimate pK a values for sulfurous acid, sulfuric acid, hydrogen sulfite and hydrogen sulfate.

Isotopic Enrichment of Uranium It requires a centrifuge that can spin at 1,500 Hz (90,000 RPM). This corresponds to a speed of greater than 2 Mach. Ex. Washing machines operate at only about 12 to 25 Hz during the spin cycle. Very strong, light materials requires that are very precisely machined. Zippe-type centrifuge 235 U is 0.72 % abundant with 238 U being 99.28% 1)Uranium ore,U 3 O 8 or "yellowcake“ is dissolved in nitric acid, producing uranyl nitrate UO 2 (NO 3 ) 2. 2) Adding NH 4 OH gives ammonium diuranate, "ADU", (NH 4 ) 2 U 2 O 7. 3) Reduction with hydrogen gives UO 2, 4) UO 2 is converted with hydrofluoric acid (HF) to uranium tetrafluoride, UF 4. 5)Oxidation with fluorine yields UF 6. Heated to 300 o C At bottom to produce convection currents UF 6