Experiment 5. The salts of hydrohalogen acids insoluble in water

Put 3-5 drops of solutions hydrochloric and hydroiodic acids salts (for example, NaCl and KI) into the test-tubes and add into each several drops of Pb(NO₃)₂. Determine the color of occurred precipitates, write the equations.

CHAPTER # 12. THE CHALCOGENS

1. Introduction in General, Organic and Biochemistry, 7^th Edition, by Morris Hein, Leo R. Best, Scott Pattison and Susan Arena, Brooks/Cole Publishing Co., 2001. (Chapter 10, pp. 289-294, Chapter 27, pp. 765-777);

2. http://www.chemsoc.org/viselements/pages/Visual Elements Group 16.html

3. http://www.hrw.com/science/mc/index/htm

1. General characteristics

The elements of Group 6A (or the main sub-group of VI group) are:

Often this group of chemical elements (except the last Po) is named the Chalcogens.

Appearance

The first element of this sub-group, Oxygen, is the only gas, and is colorless and odorless. Sulfur is a pale yellow, brittle solid. Selenium can have either an amorphous or a crystalline structure; the amorphous form can be red or black, and the crystalline form can be red or gray. Tellurium is a silvery-white color with a metallic luster. Polonium is a naturally radioactive element.

Selenium and Tellurium are rare elements with few uses, and along with Polonium will not be considered further here.

General Reactivity

Oxygen and Sulfur are highly electronegative elements - the electronegativity of Oxygen is second only to that of Fluorine. Their general reactivity is therefore dominated by their ability to gain electrons.

There is a transition down the sub-group from non-metallic to more metallic properties, so that Oxygen is a non-metal and Tellurium a metalloid. All the elements except Polonium form M^2- ions.

There is a marked difference between Oxygen and the other members of the Group. This arises from:

Ø the small size of the O atom which enables it to form double bonds;

Ø its inability to expand its valence shell like the other elements as it has no accessible d-orbitals;

Ø its high electronegativity, which enables it to participate in Hydrogen-bonding.

Occurrence and Extraction

Oxygen occurs widely as the free element in the form of O₂, comprising 21% of the air by volume. It also occurs as O₃, ozone, at high altitudes in the ozone layer. In the combined form it is found in very many minerals, and also in water. Oxygen is obtained industrially by the fractional distillation of liquid air. It is stored under pressure in cylinders.

Sulfur is found as the free element and also as metal Sulfide ores and a number of Sulfates. Native sulfur is brought to the surface from underground deposits by the Frasch Process, which uses superheated water to melt the Sulfur and force it upwards.

Physical Properties

The covalent and ionic radii increase going down the sub-group, as electrons occupy shells with higher quantum numbers.

Oxygen occurs as two gaseous allotropes, O₂ (dioxygen or more commonly Oxygen) and O₃ (trioxygen or ozone). Oxygen is more common. It condenses to a pale blue paramagnetic liquid at -183⁰C. Ozone is a pale blue, pungent gas, which condenses to an inky-blue liquid at -112⁰C. The ozone layer in the upper atmosphere is an important shield against harmful UV radiation from the Sun.

Sulfur has several allotropes, the two main ones being rhombic and monoclinic sulfur. These both consist of S₈ molecules.

Industrial Information

Sulfuric acid is of immense industrial importance. Because it has three chemical functions and is very cheap to produce, Sulfuric acid is used at some stage of the manufacture of most products. It is said that the economic prosperity of a country can be assessed by its consumption of sulfuric acid. It is manufactured by the Contact Process.

Hydrogen peroxide H₂O₂ is used to bleach hair and textiles, as a mild disinfectant and in pollution control.

Sulfur hexafluoride SF₆ cannot be ionized by electric fields and so is widely used as a gaseous insulator in transformers and electrical switchgear.

Oxygen

Oxygen is the element of the second period, its formula is 1s²2s²2p⁴.

On the assumption of external electron shell structure (graphic presentation) and electronegativity of Oxygen, it is characterized by oxidation number equal to 2-, moreover, in composition with Fluorine the oxidation number of Oxygen will be 2+ (for example, OF₂) but in peroxides it will be 1- (for example, H₂O₂) or, sometimes, -½ (for example, superoxide KO₂).

Oxygen, like Fluorine, forms compounds with all elements except Helium, Neon and Argon. Atomic Oxygen is extremely active.

In elemental state Oxygen is mostly characterized as the oxidizing agent in the reactions:

1) 4FeS + 7O₂ = 2Fe₂O₃ + 4SO₂

O20 + 4 ® 2O2-	7 (reduction);
S2- - 6 ® S+4 Fe2+ - 1 ® Fe3+	4 (oxidation).

2) 2 KI + O₃ + H₂O = I₂ + 2 KOH + O₂

O0 + 2 ® O2-	1 (reduction);
2I- - 2 ® I20	1 (oxidation).

Oxygen compounds in which the oxidation number is equal to 1- are called peroxides. For example, Hydrogen peroxide H₂O₂, where covalent and non-polar bonds in molecule connect two atoms of Oxygen.

Water solution of Hydrogen peroxide (perhydrol) presents weak dibasic acid:

H₂O₂ ↔ H⁺ + HO₂^-;

HO₂^- ↔ H⁺ + O₂^2-.

Medium of oxidation number of Oxygen in peroxides causes its oxidation-reduction duality in the reactions. H₂O₂ is the reducing agent:

H₂O₂ + 2KMnO₄ + 2KOH = 2K₂MnO₄ + 2H₂O + O₂

Mn7+ + ® Mn6+	2 (reduction);
2O- - 2 ® O20	1 (oxidation).

but in presence of strong reducing agent it is oxidizing agent:

H₂O₂ + 2KI + H₂SO₄ = I₂ + K₂SO₄ + 2H₂O

2O- + 2 ® 2O2-	1 (reduction);
2I- - 2 ® I20	1 (oxidation).

Ozone is also a highly reactive and powerful oxidizing agent which can cleave the C=C double bond.

Sulfur

In contrast to Oxygen, Sulfur atom has free d-sublevel in the external electronic shell: 1s²2s²2p⁶3s²3p⁴3d⁰.

Accordingly, Sulfur can form the compounds due to two, four and six unpaired electrons, showing such oxidation levels: 2-, 0, 2+, 4+, 6+:

Ground (nonexcited) state ( 2-, 0, 2+ oxidation number, for example, in H₂S, S, SO compounds):

The first excited state (4+ oxidation number, for example, in SO₂ compound):

The second excited state ( 6+ oxidation number, for example, in H₂SO₄ compound):

Characteristic of the most important sulfur compounds is shown in Table 16.

Table 16. Characteristic of the most important Sulfur compounds

Oxidation numbers	2-	1-		2+	4+	6+
Compounds	H2S, CuS	H2S2; FeS2 (pyrite)	S	SO	SO2, H2SO3	SO3, H2SO4
Name of corresponding acids and salts	Hydrosulfuric acid, sulfides	Persulfides	Free Sulfur	Sulfur (II) Oxide	Sulfur acid, sulphites	Sulfuric acid, sulphates
Redox properties	Only reducing agent	Properties of oxidizing and reducing agent	Properties of oxidizing and reducing agent	-	Reducing properties surpasses oxidizing ones	Only oxidizing agent
Acid-base properties	H2S - weak acid	-	-	Unstable non-Salifiable Oxide	H2SO3 - acid of medium strength, unstable in water solutions	H2SO4 - strong acid, stable in water solutions
	Strengthening of acid properties

Sulfur is the typical non-metal, which is inferior to Halogens, Oxygen and Nitrogen in electroactivity and thus is oxidized by them:

S + O₂ = SO₂;

S + 2Cl₂ = SCl₄.

As oxidizing agent, Sulfur reacts with the metals and Hydrogen at high temperature. For example,

Zn + S = ZnS;

S + H₂ = H₂S.

Hydrogen sulphide (H₂S) is the gas and occurs in nature as the result of putrescent bacteria action to proteins containing sulfur so it can be detected by its characteristic odor (“bad eggs gas”). Inhalation of pure Hydrogen sulphide can bring about to immediate death; even 0,01% of its concentration in the air is extremely dangerous for human being.

Aqueous solution of H₂S is the weak dibasic acid; therefore there are acid and neutral salts (sulphides and hydrosulphides):

H₂S ↔ H⁺ + HS^-;

HS^- ↔ H⁺ + S^2-.

H₂S is strong reducing agent. In air it burns with the generation of SO₂ or S (in case of Oxygen deficit):

2 H₂S + 3 O₂ = 2 H₂O + 2 SO₃

O20 + 4 ® 2O2-	3 (reduction);
S2- - 6 ® S4+	2 (oxidation).

or 2 H₂S + O₂ (in deficit) = 2 H₂O + 2 S

O20 + 4 ® 2O2-	1 (reduction);
S2- - 2 ® S0	2 (oxidation).

Sulfur dioxide (SO₂) is sulfurous gas (pungent, choking gas) formed when Sulfur or Sulfides are burnt in air or Oxygen, and as all fossil fuels contain Sulfur it is formed when they burn and contributes to the problem of “acid rain” :

S + O₂ SO₂ ↑;

2CuS + 3O₂ 2 CuO + 2 SO₂ ↑;

4FeS₂ + 11O₂ = 2Fe₂O₃ + 8SO₂↑.

SO₂ may be prepared by Sulfuric acid reduction:

2H₂SO₄ (conc) + Cu = CuSO₄ + SO₂ ↑ + 2H₂O.

Aqueous solution of SO₂ is called Sulfurous acid, which is unknown in free state. H₂SO₃ is of medium strength and can produce acid and neutral salts. Sulfurous acid salts decay in case of strong acids action with SO₂ isolation. This is the base for getting of Sulfur dioxide in laboratory conditions:

Na₂SO₃ + 2HCl = 2NaCl + SO₂ ↑ + H₂O.

The medium of oxidation number in Sulfur (IV) compounds causes their oxidation-reduction duality:

Na₂SO₃ + Cl₂ + H₂O = NaHSO₄ + HCl + NaCl

Cl20 + 2 ® 2Cl-	1 (reduction);
S4+ - 2 ® S6+	1 (oxidation).

SO₂ + 2 H₂S = 3 S + 2H₂O

S4+ + 4 ® S0	1 (reduction);
S2- - 2 ® S0	2 (oxidation).

SO₂ is used for fumigation of basements, cellars in order to avoid funguses. Calcium hydroSulfite Ca(HSO₃)₂ with sulfurous acid is used for sulfitation - one of the methods for preserving of tender fruits and vegetables.

Sulfur trioxide (SO₃) or Sulfur (VI) oxide is the Sulfur acid anhydride. In industry it is produced by SO₂ oxidation in presence of catalyst (Pt) (in so-called Contact Process):

2 SO₂ + O₂ = 2 SO₃.

Sulfuric acid is manufactured as the result of SO₂ interaction with water or dilute sulfuric acid:

SO₃ + H₂O = H₂SO₄.

Strong acidic properties of H₂SO₄ solutions are the basis for their use in fertilizer production or mineral premixes for animals (simple and triple superphosphates, Ammonium Sulfate):

Ca₃(PO₄)₂ +2H₂SO₄ + 5H₂O = Ca(H₂PO₄)₂·H₂O + 2[CaSO₄·2H₂O];

Simple superphosphate

Ca₃(PO₄)₂ + 3H₂SO₄ = 2H₃PO₄ + 3CaSO₄;

Ca₃(PO₄)₂ + 4H₃PO₄ + 3H₂O = 3 Ca(H₂PO₄)₂·H₂O.

Triple superphosphate

H₂SO₄ + 2NH₃ = (NH₄)₂SO₄.

Concentrated H₂SO₄ is a fairly strong oxidizing agent. Non-metals are oxidized by it up to oxides and H₂SO₄ itself reduces to SO₂:

C + 2 H₂SO₄ (conc) = CO₂ + SO₂ + 2 H₂O

S6+ + 2 ® S4+	2 (reduction);
C0 - 4 ® C4+	1 (oxidation).

PRACTICE PROBLEMS

1. Detect Oxygen and Sulfur oxidation numbers in the following compounds: O₂, O₃, CaO, BaO, OF₂, S₈, FeS, S₂, K₂SO₄, Na₂SO₃, SO₃, Al₂(SO₄)₃.

2. At which numbers of oxidation does Sulfur present oxidation-reduction duality?

3. In which cases does Oxygen present the properties of reducing agent?

4. Why Hydrogen Sulfide cannot act as the oxidizing agent?

5. Set up the reaction of Sulfuric acid getting taking into account the following series of conversions:

FeS₂ ® SO₂ ® H₂SO₃ ® H₂SO₄.

6. Set up the reaction according to series of conversions:

SO₃ ® H₂SO₄ ® (NH₄)₂SO₄ ® CaSO₄.

7. Complete the equations of redox reactions:

H₂S + H₂O₂ ®

HClO + H₂O₂ ®

KMnO₄ K₂MnO₄ + MnO₂ + O₂

S + NaOH ® Na₂S + Na₂SO₃ + …

NaHSO₃ + Cl₂ + H₂O ® NaHSO₄ + …

LABORATORY TRAINING

Experiment 1. Oxidizing properties of H₂O₂

Add 2-3 drops of H₂SO₄ and 3-4 drops of Hydrogen Peroxide to 2-3 drops of Potassium or Sodium iodide. Prove free Iodine release in given reaction by adding starch solution. Set up the reaction.

Experiment 2. Reducing properties of H₂O₂

Drop H₂O₂ up to discoloration of the mixture to the test-tube to 2-3 drops of KMnO₄, acidified by 1-2 drops of dilute H₂SO₄. What gas is released? Set up the reaction.

Experiment 3. Reducing properties of S^2-

Fill test-tube with 5-7 drops of hydroSulfide water and add some drops of chloric water up to mixture turbidity.

Fill another test-tube with 5-7 drops of Sodium or Ammonium Sulfide. Add 3-4 drops of concentrated nitric acid and watch solution turbidity as the result of sulfur generation. Set up the reactions.

Experiment 4. Reducing properties of S⁴⁺

Fill test-tube with 5-7 drops of Na₂SO₃, acidify by 3-4 drops of H₂SO₄ and add 3-5 drops of iodic water (acting component I₂). Why does the color of I₂ disappear? Write the reaction.

Experiment 5. Oxidizing properties of S⁴⁺

Acidify 5-6 drops of Na₂SO₃ by 2-3 drops of diluted H₂SO₄ and add some drops of hydrosulfuric water up to mixture turbidity. Write the reaction.

Date: 2015-01-12; view: 846