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Group 17 -The Halogens -The only group containing solids, liquidsand 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

Cl2

Br2

I2

Melting Point

Boiling Point

State (at 20 °C)

Density (at 20 °C)

Fluorine -220 °C -188 °C Gas 0.0017 g/cm3

Chlorine -101 °C -34 °C Gas 0.0032 g/cm3

Bromine -7.25 °C 58.8 °C Liquid 3.123 g/cm3

Iodine 114 °C 185 °C Solid 4.93 g/cm3

Electronic Properties– - Form monoatomic anions with -1 charge. – - 7 valence electrons: [N.G.] ns2 np5

– - 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 1681 -328.0 +2.866

Chlorine 1251 -349.0 +1.358

Bromine 1140 -324.6 +1.065

Iodine 1008 -295.2 +0.535

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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, somoist chlorine gas is a much stronger oxidizing agent thananhydrous chlorine – as evidenced by the fact that it can bleach dyeslike indigo, litmus and malachite green while anhydrous chlorinecannot.

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Reactions with Metals

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Reactions with P and S– Halogens react with phosphorus (P4) to give PX3 or PX5.

– Halogens react with sulfur (S8) to give S2X2 or SX6 (depending on the

– ratio of reactants). Many other sulfur-halogen compounds can be– made (including, but not limited to, SX2, SX4 and S2X4).

*See http://boyles.sdsmt.edu/pwithcl/reaction_of_white_phosphorus_and.htm for a video 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:– Cl2 + 2 OH- → ClO- + Cl- + H2O

– These are disproportionation reactions because one oxidation state of

– chlorine (Cl2) reacts to give two different oxidation states of chlorine (Cl– and an oxoanion). – Other such disproportionation reactions continue:–

• 2 ClO− → Cl− + ClO2

− • • ClO− + ClO2

− → Cl− + ClO3

−

• Sodium chlorate crystals can be heated to undergo yet another reaction:

• 4 ClO3− → Cl− +

3 ClO4−

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, buttreating every bond as 100% ionic, so that both electrons in the bondgo 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 StateThe 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) LiCl2) BeCl23) BCl34) CCl4

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) CH4

2) CH3OH3) CH2O4) CHOOH5) CO2

Rules for Assigning Oxidation States1) 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 StateDetermine the oxidation state of all the element in the following molecules:

a) H2 H = 0

b) CO2O = -2 0 = C + 2(-2)

c) BF3 F= -1 0 = B + 3(-1)

d) H2SO4 H = 1 O = -2 0 = S +(2(+1) + 4(-2)) = S - 6

e) SO2ClF O = -2 F = -1 Cl = -1 0 = S +(2(-2) -1 -1) = S - 6

f) IO2F2_ O = -2 F = -1 -1 = I + (2(-2) + 2(-1)) = I - 6

C = 4

B = 3

S = 6

S = 6

g) HPO42- O = -2 H = 1

I = 5

-2 = P + (4(-2) + 1) = P - 7 P = 5

Oxidation States of Poly Atomic Ions

PO43- SO4

2-

NO3-

ClO3-

CO32-

NO2-

PO33- SO3

2- ClO2-

4 5

3

5

3

6

4

5

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Oxoanions of the halogensThe name of an oxoanion depends on how many oxygen atoms surroundthe 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. HCO3

- is “hydrogen carbonate”

Cl N C S PO hypochlorite

(ClO-)

O2 chlorite(ClO2

-)nitrite(NO2

-)

O3 chlorate(ClO3

-)nitrate(NO3

-)carbonate

(CO32-)

sulfite(SO3

2-)phosphite

(PO33-)

O4 perchlorate(ClO4

-)sulfate(SO4

2-)phosphate

(PO43-)

Oxoacids of the HalogensNote that if enough H+ have been added to render the oxoanion neutral, it isno longer an oxoanion. It is an oxoacid! Note that the hydrogens areattached to oxygen in most oxoacids.*

Cl N C S PHydrochloric acid

(HCl)

O Hypochlorous acid(HOCl)

O2 Chlorous acid(HClO2)

Nitrous acid(HNO2)

O3 Chloric acid(HClO3)

Nitric acid(HNO3)

Carbonic acid

(H2CO3)

Sulfurous acid(H2SO3)

Phosphorous acid*

(H3PO3)

O4 Perchloric acid(HClO4)

Sulfuric acid(H2SO4)

Phosphoric acid(H3PO4)

*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 OpE(OH)q.

For multiple protons (i.e. q>1), pKa increases by ~5 every time H+ is removed.

Exercise) Estimate pKa values for sulfurous acid, sulfuric acid, hydrogensulfite and hydrogen sulfate.

ppKa 58

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

235U is 0.72 % abundant with 238U being 99.28%1) Uranium ore,U3O8 or "yellowcake“ is dissolved in nitric acid, producing uranyl nitrate UO2(NO3)2. 2) Adding NH4OH gives ammonium diuranate, "ADU", (NH4)2U2O7. 3) Reduction with hydrogen gives UO2, 4) UO2 is converted with hydrofluoric acid (HF) to uranium tetrafluoride, UF4.5)Oxidation with fluorine yields UF6.

Heated to 300 oCAt bottom to produce convection currents

UF6