SAT Chemistry Some Representative Groups and Families - Sulfur Family

SAT Chemistry Some Representative Groups and Families - Sulfur Family

Since we discussed oxygen in Chapter 5, the most important element in this family left to discuss is sulfur.

Sulfur is found free in the volcanic regions of Japan, Mexico, and Sicily. It is removed from the rock mixtures by heating in retorts or furnaces.

Sulfuric Acid

IMPORTANT PROPERTIES OF SULFURIC ACID. Sulfuric acid ionizes in two steps:
H₂SO₄(l) + H₂O(l) ⇌ H₃O⁺(aq) + HSO₄^-(aq)                        K_a1 is very large
HSO₄^-(aq) + H₂O(l) ⇌ H₃0⁺(aq) + SO₄^2-(aq)                     K_a2 is very small
to form a strong acid solution. The ionization is more extensive in a dilute solution. Most hydronium ions are formed in the first step. Salts formed with the HSO₄^- (bisulfate ion) are called acid salts; the SO₄^2- (sulfate ion) forms normal salts.

Sulfuric acid reacts like other acids, as shown below:

With active metals:     Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)
(for dilute H₂SO₄)
With bases:                  2NaOH(aq) + H₂SO₄(aq) → Na₂SO₄(aq) + 2H₂O(l)
With metal oxides:      MgO(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂O(l)
With carbonates:         CaCO₃(s) + H₂SO₄(aq) → CaSO₄(aq) + H₂O(l) + CO₂(g)
Sulfuric acid has other particular characteristics.

As an oxidizing agent:
Cu(s) + 2H₂SO₄(aq) (concentrated) → CuSO₄(aq) + SO₂(g) + 2H₂O(l)
As a dehydrating agent with carbohydrates:

Other Important Compounds of Sulfur
Hydrogen sulfide is a colorless gas having an odor of rotten eggs. It is fairly soluble in water and is poisonous in rather small concentrations. It can be prepared by reacting ferrous sulfide with an acid, such as dilute HC1:
FeS(s) + 2HCl(aq) → FeCl₂(aq) + H₂S(g)

Hydrogen sulfide burns in excess oxygen to form compounds of water and sulfur dioxide. If insufficient oxygen is available, some free sulfur will form. It is only a weak acid in a water solution. Hydrogen sulfide is used widely in qualitative laboratory tests since many metallic sulfides precipitate with recognizable colors. These sulfides are sometimes used as paint pig­ments. Some common sulfides and their colors are:

ZnS—White
CdS—Bright yellow
As₂S₃—Lemon yellow
Sb₂S₃—Orange
CuS—Black
HgS—Black
PbS—Brown-black

Another important compound of sulfur is sulfur dioxide. It is a colorless gas with a suf­focating odor.
The structure of sulfur dioxide is a good example of resonance structures. Its molecule is depicted in Figure
39.

You will notice in Figure 35 that the covalent bonds between sulfur and oxygen are shown in one drawing as single bonds and in the other as double bonds. This signifies that the bonds between the sulfur and oxygens have been shown by experimentation to be neither single nor double bonds, but “hybrids” of the two. Sulfur trioxide, shown below, also has resonance structures.

HALOGEN FAMILY
The common members of the halogen family are shown in Table 13 with some important facts concerning them.

Because each halogen lacks one electron in its outer principal energy level, these elements usually are acceptors of electrons (oxidizing agents). Fluorine is the most active nonmetal in the periodic chart.
Some Important Halides and Their Uses
Hydrochloric acid —common acid prepared in the laboratory by reacting sodium chloride with concentrated sulfuric acid. It is used in many important industrial processes.
Silver bromide and silver iodide —halides used on photographic film. Light intensity is recorded by developing as black metallic silver those portions of the film upon which the light fell during exposure.
Hydrofluoric acid —acid used to etch glass by reacting with SiO₂ to release silicon fluoride gas. Also used to frost lightbulbs.
Fluorides —used in drinking water and toothpaste to reduce tooth decay.

NITROGEN FAMILY
The most common member of this family is nitrogen itself. It is a colorless, odorless, tasteless, and rather inactive gas that makes up about four-fifths of the air in our atmosphere. The inactivity of N₂ gas can be explained by the fact that the two atoms of nitrogen are bonded

must be “pushed” into combining with other elements, many of its compounds tend to decompose violently with a release of the energy that went into forming them.
Nature “fixes” nitrogen, or makes nitrogen combine, by means of a nitrogen-fixing bacteria found in the roots of beans, peas, clover, and other leguminous plants. Discharges of lightning also cause some nitrogen fixation with oxygen to form nitrogen oxides.

Nitric Acid
An important compound of nitrogen is nitric acid. This acid is useful in making dyes, celluloid film, and many of the lacquers on cars.
Its physical properties are: it is a colorless liquid (when pure), it is one and one-half times as dense as water, it has a boiling point of 86°C, the commercial form is about 68% pure, and it is miscible with water in all proportions.
Its outstanding chemical properties are: the dilute acid shows the usual properties of an acid except that it rarely produces hydrogen when it reacts with metals, and it is quite unsta­ble and decomposes as follows:
4HNO₃(aq) → 2H₂O(l) + 4NO₂(g) + O₂(g)

Because of this ease of decomposition, nitric acid is a good oxidizing agent. When it reacts with metals, the nitrogen product formed will depend on the conditions of the reaction, especially the concentration of the acid, the activity of the metal, and the temperature. If the nitric acid is concentrated and the metal is copper, the principal reduction product will be nitrogen dioxide (N0₂), a heavy, red-brown gas with a pungent odor.
Cu(s) + 4HNO₃(aq)  → Cu(NO₃)₂(aq) + 2NO₂(g) + 2H₂O(l)
With dilute nitric acid, this reaction is:
3Cu + 8HNO₃ → 3Cu(NO₃)₂ + 4H₂O+ 2NO(g)
The product NO, called nitric oxide, is colorless and is immediately oxidized in air to NO₂ gas.
With still more dilute nitric acid, considerable quantities of nitrous oxide (N₂O) are formed; with an active metal like zinc, the product may be the ammonium ion (NH₄⁺).
When nitric acid is mixed with hydrochloric acid, the mixture is called aqua regia because of its ability to dissolve gold.

METALS
Properties of Metals
Some physical properties of metals are: they have metallic luster, they can conduct heat and electricity, they can be pounded into sheets (are malleable), they can be drawn into wires (are ductile), most have a silvery color, and none is soluble in any ordinary solvent without a chemical change.
The general chemical properties of metals are: they are electropositive, and the more active metallic oxides form bases, although some metals form amphoteric hydroxides that can react as both acids and bases.

Some Important Reduction Methods of Iron Ore
Iron ore is refined by reduction in a blast furnace, that is, a large, cylinder-shaped furnace charged with iron ore (usually hematite, Fe₂0₃), limestone, and coke. A hot air blast, often enriched with oxygen, is blown into the lower part of the furnace through a series of pipes called tuyeres. The chemical reactions that occur can be summarized as follows:

The molten iron from the blast furnace is called pig iron.
From pig iron, the molten metal may undergo one of three steel-making processes that burn out impurities and set the contents of carbon, manganese, sulfur, phosphorus, and silicon. Often nickel and chromium are alloyed in steel to give the particular properties of hardness needed for tool parts. The three most important means of making steel involve the basic oxygen, the open-hearth, and the electric furnaces. The first two methods are the most common.
The basic oxygen furnace uses a lined “pot” into which the molten pig iron is poured. Then a high-speed jet of oxygen is blown from a water-cooled lance into the top of the pot. This “burns out” impurities to make a batch of steel rapidly and cheaply.

The open-hearth furnace is a large oven containing a dish-shaped area to hold the molten iron, scrap steel, and other additives with which it is charged. Alternating blasts of flame are directed across the surface of the melted metal until the proper proportions of additives are established for that “heat” so that the steel will have the particular properties needed by the customer. The tapping of one of these furnaces holding 50 to 400 tons of steel is a truly beau­tiful sight.
The final method of making steel involves the electric arc furnace. This method uses enormous amounts of electricity through graphite cathodes that are lowered into the molten iron to purify it and produce a high grade of steel.
Alloys
An alloy is a mixture of two or more metals. In a mixture certain properties of the metals involved are affected. Three of these are:

1. Melting point             The melting point of an alloy is lower than that of its components.
2. Hardness                    An alloy is usually harder than the metals that compose it.
3. Crystal structure      The size of the crystalline particles in the alloy determines many of the physical properties. The size of these particles can be controlled by heat treatment. If the alloy cools slowly, the crystalline particles tend to be larger. Thus, by heating and cooling an alloy, its properties can be altered considerably.

Common alloys are:

1. Brass, which is made up of copper and zinc.
2. Bronze, which is made up of copper and tin.
3. Steel, which has controlled amounts of carbon, manganese, sulfur, phosphorus, and silicon, is alloyed with nickel and chromium.
4. Sterling silver, which is alloyed with copper.

Metalloids
In the preceding sections, representative metals and nonmetals have been reviewed, along with the properties of each. Some elements, however, are difficult to classify as one or the other. One example is carbon. The diamond form of carbon is a poor conductor, yet the graphite form conducts fairly well. Neither form looks metallic, so carbon is classified as a nonmetal.
Silicon looks like a metal. However, its conductivity properties are closer to those of carbon.
Since some elements are neither distincdy metallic nor clearly nonmetallic, a third class, called the metalloids, is recognized.

The properties of metalloids are intermediate between those of metals and those of non­metals. Although most metals form ionic compounds, metalloids as a group may form ionic or covalent bonds. Under certain conditions pure metalloids conduct electricity, but do so poorly, and are thus termed semiconductors. This property makes the metalloids important in microcircuitry.
The metalloids are located in the periodic table along the heavy dark line that starts along­side boron and drops down in steplike fashion between the elements found lower in the table (see Figure 36).