Oxygen and Sulfur - Department of Chemistry

Chemistry 1051 Descriptive Chemistry Material on Midterm #3
Intersession

Note: This handout is meant to highlight important facts and reactions. Exam
questions can be based on anything in the assigned text pages. You must read
each section carefully.

Oxygen and Sulfur

PHHM (9th Edition): Section 22-3
PHH (8th Edition): Section 8-3 and 23-3

Oxygen: approx 20% of atmosphere by volume

most abundant element in earth’s crust where it is combined with other
elements in compounds e.g. Fe2O3.

Preparation of oxygen

1. by fractional distillation of liquid air (most common method)

2. in small quantities for breathing in air-tight places like spacecrafts, submarines

4 KO2(s) + 2 CO2(g) → 2 K2CO3(s) + 3 O2(g)

NOTE: O2⎯ disproportionates to form O2 and O2⎯ in the reaction above. O2⎯
ions react with CO2(g) to form CO32⎯ ions.

Question: Write half-reactions for the oxidation of O2⎯ to O2 and its reduction to O2⎯.
Use these to write the balance disproportionation reaction for O2⎯ forming
O2 and O2⎯.

Answer: [O2⎯ → O2 + e⎯] x 3
oxidation: [O2⎯ + 3 e⎯ → 2 O2⎯]
reduction: 4 O2⎯ → 3 O2 + 2 O2⎯
overall:

3. by electrolysis of dilute solutions of strong electrolytes (acids, bases, unreactive
salts like Na2SO4). The ions from the strong electrolyte increase the conductivity
of the H2O.

2 H2O(l) → 2 H2(g) + O2(g)

oxidation of anode: 2 H2O(l) → O2(g) + 4 H+(aq) + 4 e⎯
reduction of cathode: 4 H2O(l) + 4 e⎯ → 2 H2(g) + 4 OH⎯(aq)

Allotropy: the existence of two or more forms of an element that differ in their
bonding and molecular structure.

Allotropes of oxygen: dioxygen O2 and trioxygen O3 (ozone)

Properties of ozone

1. third most powerful oxidizing agent after F2 and OF2
2. used to purify drinking water in place of Cl2
3. unhealthy to breathe
4. found mostly in upper atmosphere and in smog
5. ozone in upper atmosphere absorbs harmful UV radiation from the sun.

Uses of oxygen: combustion of organic compounds (very exothermic)

e.g. 2 C2H6(g) + 7 O2(g) → 4 CO2(g) + 6 H2O(l)
C2H6O(l) + 3 O2(g) → 2 CO2(g) + 3 H2O(l)

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Sulfur: exists as a yellow solid at RT and consists of S8 molecules

Comparison of oxygen and sulfur

Similarities: both form ionic compounds with active metals
both form similar covalent compounds e.g. H2O and H2S, CO2 and CS2

Differences: oxygen is a lot smaller, considerably more electronegative, and cannot
form an expanded octet.

H2O vs. H2S: H2O is an extensively hydrogen bonded liquid, a highly polar solvent
and a good ligand.
H2S is a poisonous gas, only weakly polar and is a poor solvent.

Production of sulfur

1. Frasch process: super heated steam is forced into sulfur containing rock.
The melted sulfur is forced to the surface by compressed air.

2. S is recovered from H2S, a major impurity in oil and natural gas.
H2S is split into two streams. Half is oxidized to SO2 which is then allowed to
react with H2S to form S(s).

2 H2S(g) + 3 O2(g) → 2 SO2(g) + 2 H2O(g) combustion
SO2(g) + 2 H2S(g) → 3 S(s) + 2 H2O(g) reverse disproportionation

Industrial production of concentrated sulfuric acid

This occurs in a series of steps

1. S(s) + O2(g) → SO2(g)
2. SO2(g) + O2(g) → 2 SO3(g) needs catalyst, mainly V2O5(s)
3. SO3(g) + H2SO4(l) → H2S2O7(l), disulfuric acid
4. H2S2O7(l) + H2O(l) → 2 H2SO4(l) i.e. pure sulfuric acid

5. H2SO4(l) ⎯H⎯⎯2O→ H2SO4(conc., aq)

i.e. concentrated sulfuric acid (95−98% H2SO4 by mass)

Disulfuric acid: It can be formed by heating pure liquid H2SO4 in which H2O is lost.
One molecule acts as a Lewis base donating a lone pair from oxygen
(OH) to the sulfur of the second molecule which acts as a Lewis acid.

2 H2SO4(l) → H2S2O7(l) + H2O(g)

The structure of disulfuric acid:

Properties of concentrated sulfuric acid.

1. It has a great affinity for H2O and can dehydrate some organic compounds.

conc. H 2SO 4
⎯⎯⎯⎯⎯→e.g. C12H22O11(s) 12 C(s) + 11 H2O(l)

2. It is a moderately good oxidizing agent.

Cu(s) + 2 H2SO4(conc.) → CuSO4(aq) + 2 H2O(l) + SO2(g)

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Formation of sulfurous acid (contains sulfite ion)

SO2(g) + H2O(l) → H2SO3(aq)

Formation of thiosulfate ion

SO32⎯(aq) + S(s) → S2O32⎯(aq) reverse disproportionation
Note: S2O32⎯ is the most common titrant for I2 and is used in photographic processing.

APHciHdMra(i9nthaEndditSioOn)2 emissions
page 258
PHH (8th Edition) page 700-701

SO2 in the atmosphere comes mainly from the combustion of sulfur-containing coal and
from the roasting of sulfide ores. Roasting involves the high temperature reaction of
sulfide ores with oxygen in order to convert them to the corresponding oxides. The
oxides can then be reduced to the metal by reaction with carbon or other suitable
reducing agent at high temperature.

S(s) from coal + O2(g) ⎯⎯∆→ SO2(g)
2 ZnS(s) + 3 O2(g) ⎯⎯∆→ 2 ZnO(s) + 2 SO2(g)

ZnO(s) + CO(g) ⎯⎯∆→ Zn(s) + CO2(g)

SO2 can be oxidized to SO3 and by reaction with H2O produces sulfuric acid which
accounts for well over half of the acidity of acid rain.

Questions:

PHHM (9th Edition): Chapter 22: 17, 21, 23, 27, 29, 32, 33, 35, 37
PHH(8th Edition): Chapter 23: 33, 35, 36, 38, 39, 41, 43.

1. Write balanced equations for
(a) the reaction of KO2(s) with CO2(g) to yield O2(g) and K2CO3(s)
(b) the sequence of reactions by which H2SO4(aq) is formed from S(s)
(c) the oxidation of Zn(s) by H2SO4(conc.) yielding ZnSO4, H2O(l) and SO2(g)
(d) the dehydration of glucose C6H12O6 by conc. H2SO4
(d) the roasting of CuS(s) with O2(g) in which SO2(g) is formed.

2. Briefly describe the Frasch process for producing S(s).

3. Compare the properties of H2O and H2S.

4. Complete the following reactions leading to the production of sulfur from H2S:

___ H2S(g) + ___ O2(g) →

___ H2S(g) + ___ SO2(g) → (reverse disproportionation)

5. Explain the role of SO2 in acid rain. What are the main sources of SO2 in the
atmosphere?

6. Distinguish between allotropes and isotopes of an element. Give examples of
both. (A good example of isotopes of an element are “simple” hydrogen and
deuterium.)

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Nitrogen and phosphorous

PHHM (9th Edition): Chapter 22, pp 939-944 and 938-948
PHH (8th Edition): Section: 8-2 page 270-275

23-4 page 925-932

Nitrogen: N2: stable, diatomic gas found mostly in the atmosphere

Ammonia and related compounds

formation: The Haber process

3 H2(g) + N2(g) ⇌ 2 NH3(g) T > 400°C, Fe catalyst, Ptotal > 200 atm

NH3 as a base: neutralization by sulfuric acid to form ammonium sulfate, an
important fertilizer

∆,Pt −Rh
catalyst
⎯⎯⎯⎯→2 NH3(aq) + H2SO4(aq) (NH4)2SO4(aq)

Formation of N2O: laughing gas

NH4NO3(s) ⎯⎯∆→ N2O(g) + 2 H2O(g) reverse disproportionation

Commercial synthesis of nitric acid: oxidation of NH3 to HNO3:
1. 4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g) Pt catalyst, T = 850°C
2. 2 NO(g) + O2(g) → 2 NO2(g)
3. 3 NO2(g) + H2O(l) → 2 HNO3(aq) + NO(g) disproportionation

Phosphorus

Production: reduction of phosphate rock (apatite) in an electric furnace

⎯⎯→2 Ca3(PO4)2(s) + 10 C(s) + 6 SiO2(s) ∆ 6 CaSiO3(l) + 10 CO(g) + P4(g)

Allotropic forms: tetrahedral P4 molecules called white phosphorous. When heated
one P-P bond breaks and the fragments join together into long
chains (P4)n called red phosphorus.

Oxidation of P4

P4(s) + 3 O2(g) → P4O6(s) limited O2
P4(s) + 5 O2(g) → P4O10(s) excess O2

Formation of oxoacids from anhydrides (oxides) phosphorous acid
phosphoric acid
P4O6(s) + 6 H2O(l) → 4 H3PO3(aq)
P4O10(s) + 6 H2O(l) → 4 H3PO4(aq)

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Formation of polyphosphoric acids:

Two phosphoric acid molecules combine to lose H2O (H from one molecule and OH
from another) forming P-O-P bonding to form diphosphoric acid. This is a similar
reaction to the formation of disulfuric acid from sulfuric acid. Additional phosphoric acid
molecules combine losing one H2O for each one added.

e.g. H3PO4(l) + H3PO4(l) → H4P2O7(l) (diphosphoric acid) + H2O

Questions:

PHHM (9th Edition): Chapter 22 43, 45, 49, 51
PHH (8th Edition): Chapter 8 10(b), 11, 12, 15, 25.
Chapter 23 49, 51

Additional questions:

1. Draw a Lewis structure for N2. Use this to explain why nitrogen gas is generally
unreactive.

2. Write equations to describe the commercial synthesis of HNO3(aq) i.e. the
oxidation of NH3 by O2 to form NO and H2O(g), the oxidation of NO by O2 and the
disproportionation of NO2 in water forming HNO3(aq) and NO(g).

3. Write equations for (a) oxidation of P4 to P4O6 (b) the formation of H3PO4 from
P4O10 and H2O.

4. Distinguish between red and white phosphorus.
5. Write a balanced equation for the formation of triphosphoric acid H5P3O10 from

three molecules of phosphoric acid.

Developed by Dr. Chris Flinn

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