Metals and their Compounds

CHEMISTRY SSS3 First term

WEEK 1

Metals and their Compounds

Performance objectives

Students should be able to:

1. State the general characteristics of metals.
2. State the compounds of metals.
3. Explain the principle of extraction of metals.
4. Explain the extraction of sodium.
5. Give at least three uses of sodium.

Content

Metals and their compounds

General characteristics of metals

Metals are a class of substances characterized by high electrical and thermal conductivity as well as by malleability, ductility, and high reflectivity of light.

Approximately three-quarters of all known chemical elements are metals. The most abundant varieties in the Earth’s crust are aluminum, iron, calcium, sodium, potassium, and magnesium. The vast majority of metals are found in ores (mineral-bearing substances), but a few such as copper, gold, platinum, and silver frequently occur in the free state because they do not readily react with other elements.

Characteristics of metals

Metals are lustrous, malleable, ductile, good conductors of heat and electricity. However, some other characteristics include:

• State: Metals are solids at room temperature with the exception of mercury, which is liquid at room temperature (Gallium is liquid on hot days).
• Luster: Metals have the quality of reflecting light from their surface and can be polished e.g., gold, silver, and copper.
• Malleability: Metals have the ability to withstand hammering and can be made into thin sheets known as foils. For example, a sugar cube-sized chunk of gold can be pounded into a thin sheet that will cover a football field.
• Ductility: Metals can be drawn into wires. For example, 100 g of silver can be drawn into a thin wire about 200 meters long.
• Hardness: All metals are hard except sodium and potassium, which are soft and can be cut with a knife.
• Valency: Metals typically have 1 to 3 electrons in the outermost shell of their atoms.
• Conduction: Metals are good conductors because they have free electrons. Silver and copper are the two best conductors of heat and electricity. Lead is the poorest conductor of heat. Bismuth, mercury, and iron are also poor conductors
• Density: Metals have high density and are very heavy. Iridium and osmium have the highest densities whereas lithium has the lowest density.
• Melting and Boiling Points: Metals have high melting and boiling points. Tungsten has the highest melting and boiling points whereas mercury has the lowest. Sodium and potassium also have low melting points.

Compounds of metals

Most metals do not occur in their natural state. They are often found as compounds such as metal oxides, sulfides and halides.

• Aluminum oxide is the main metal compound present in the ore known as bauxite.
• Iron pyrites or ‘fool’s gold’ is mainly iron sulfide.
• The Lake Grassmere salt works in Marlborough produces the metal compound known as sodium chloride from ‘salty’ seawater.

Metals can be produced (smelted) from their ores by a variety of methods:

• Aluminum is produced from its ore (bauxite) by passing a very large electric current through a molten mixture of the ore and a compound called cryolite.
• Titanium is mostly produced from its ore (rutile) by the Kroll Process, where the ore is treated with chlorine gas followed by a reaction with magnesium metal.

General Principles of Extraction of Metals

In order to extract the metal from ores, several physical and chemical methods are used. The method used depends upon the nature of the ore, the properties of the metal, and the local conditions, Thus, it is not possible to have a universal method for the extraction of all the metals from their ores. 

                        

What are "ores"?

An ore is any naturally-occurring source of a metal that you can economically extract the metal from. Aluminum, for example, is the most common metal in the Earth's crust, occurring in all sorts of minerals. However, it isn't economically worthwhile to extract it from most of these minerals. Instead, the usual ore of aluminum is bauxite - which contains from 50 - 70% of aluminum oxide.

Copper is much rarer, but fortunately can be found in high-grade ores (ones containing a high percentage of copper) in particular places. Because copper is a valuable metal, it is also worth extracting it from low-grade ores as well. Ores are common oxides, for example:

• bauxite (Al₂O₃)
• haematite (Fe₂O₃)
• rutile (TiO₂ )

     . Or sulfides, for example:

• pyrite (FeS₂ )
• chalcopyrite (CuFeS₂ )

Concentrating the ore

This simply means getting rid of as much of the unwanted rocky material as possible before the ore is converted into the metal. In some cases, this is done chemically. For example, pure aluminum oxide is obtained from bauxite by a process involving a reaction with sodium hydroxide solution. This is described in detail on the aluminum page in this section. Some copper ores can be converted into copper (II) sulfate solution by leaving the crushed ore in contact with dilute sulphuric acid for a long time. Copper can then be extracted from the copper (II) sulfate solution. But, in many cases, it is possible to separate the metal compound from unwanted rocky material by physical means. A common example of this involves froth flotation.

Froth flotation

The ore is first crushed and then treated with something which will bind to the particles of the metal compound that you want and make those particles hydrophobic. "Hydrophobic" literally means "water fearing". In concentrating copper ores, for example, pine oil is often used. The pine oil binds to the copper compounds, but not to the unwanted rocky material.

The treated ore is then put in a large bath of water containing a foaming agent (a soap or detergent of some kind), and air is blown through the mixture to make a lot of bubbles. Because they are water-repellent, the coated particles of the metal compound tend to be picked up by the air bubbles, float to the top of the bath, and are allowed to flow out over the sides. The rest of the rocky material stays in the bath.

Reducing the metal compound to the metal

Why is this reduction?

At its simplest, where you are starting from metal oxides, the ore is being reduced because oxygen is being removed.

 

Fe₂O₃ -------à Fe

Removal of oxygen = Reduction

 Al₂O₃ --------à Al

Removal of oxygen = Reduction

However, if you are starting with a sulfide ore, for example, that's not a lot of help! It is much more helpful to use the definition of reduction in terms of addition of electrons. To a reasonable approximation, you can think of these ores as containing positive metal ions. To convert them to the metal, you need to add electrons - reduction.

Al³⁺ + 3e^- -----à Al     (addition of electron = reduction)

Fe³⁺ + 3e^- -----à Fe   (addition of electrons = reduction)

Choosing a method of reduction

There are various economic factors you need to think about in choosing a method of reduction for a particular ore. These are all covered in detail on other pages in this section under the extractions of particular metals. What follows is a quick summary.

You need to consider:

• The cost of the reducing agent;
• energy costs;
• The desired purity of the metal.

There may be various environmental considerations as well - some of which will have economic costs.

Chemical Reduction

Carbon (as coke or charcoal) is cheap. It not only acts as a reducing agent, but it also acts as the fuel to provide heat for the process. However, in some cases (for example with aluminum) the temperature needed for carbon reduction is too high to be economic - so a different method has to be used. Carbon may also be left in the metal as an impurity. Sometimes this can be removed afterward (for example, in the extraction of iron); sometimes it can't (for example in producing titanium), and a different method would have to be used in cases like this.

Other more reactive metals can be used to reduce the ore. Titanium is produced by reducing titanium (IV) chloride using a more reactive metal such as sodium or magnesium. As you will see if you read the page about titanium extraction, this is the only way of producing high purity metal.

TiCl₄ + 4Na → Ti + 4NaCl (1)

The more reactive metal sodium releases electrons easily as it forms its ions:

4Na → 4Na⁺ + 4e⁻ (2)

These electrons are used to reduce the titanium (IV) chloride:

TiCl₄ + 4e⁻ → Ti + 4Cl⁻ (3)

The downside of this is expense. You have first to extract (or to buy) the sodium or magnesium. The more reactive the metal is the more difficult and expensive the extraction becomes. That means that you have to use a very expensive reducing agent to extract the titanium. As you will see if you read the page about titanium extraction, there are other problems in its extraction which also add to the cost.

 

 

Reduction by electrolysis

This is a common extraction process for the more reactive metals - for example, for aluminum and metals above it in the electrochemical series. You may also come across it in other cases such as one method of extracting copper and in the purification of copper. During electrolysis, electrons are being added directly to the metal ions at the cathode (the negative electrode). The downside (particularly in the aluminum case) is the cost of the electricity. An advantage is that it can produce very pure metals.

 General properties of Alkali Metals

 Alkali metals are metals present in group 1 elements and they include sodium, lithium, potassium, caesium, rubidium, and francium. They are univalent elements because they have one valence electron. They ionize readily by donating one electron.

Na   ----à Na⁺ + e^-

They form positive ions. Hence, they are good reducing agents.

They form electrovalent compounds. They are good conductors of electricity and are very electropositive.

They react vigorously with cold water to alkalis.

Na₂O + H₂O ----à 2NaOH _(aq)

Sodium

Sodium (Na), chemical element of the alkali metal group (Group 1 [Ia]) of the periodic table. Sodi­um is an el­e­ment in the 1ˢᵗ group of the 3ʳᵈ pe­ri­od of the Pe­ri­od­ic Ta­ble. It is in the sub­group of al­ka­line met­als; Na₂O sodi­um ox­ide and NaOH sodi­um hy­drox­ide dis­play typ­i­cal base prop­er­ties. The most im­por­tant nat­u­ral com­pounds of sodi­um are ta­ble salt NaCl, Glauber’s salt Na₂­SO₄·10H₂O and syl­van­ite NaCl·KCl. Sodium is a very soft silvery-white metal. Sodium is the most common alkali metal and the sixth most abundant element on Earth, comprising 2.8 percent of Earth’s crust. It occurs abundantly in nature in compounds, especially common salt—sodium chloride (NaCl)—which forms the mineral halite and constitutes about 80 percent of the dissolved constituents of seawater.

Extraction of sodium

On industrial scale, sodium metal is extracted by "Down's Process. Principle
Down's Process is based on the electrolysis of fused NaCl. Construction of down cell
Down's cell consists of a rectangular container of steel. Inside of the tank is lined with firebricks. anode is a graphite rod which projects centrally up through the base of the cell. Cathode is a ring of iron, which surrounds the anode. The anode and cathode are separated from each other by a cylindrical steel gauze diaphragm so that Na and Cl2 are kept apart. A bell like hood is submerged over the anode.                                            Down’s cell
Draw backs on down’s cell   The melting point of NaCl is 801C. At this temperature, molten NaCl and Na form a metallic fog in the container which is impossible to separate.
Steps to overcome this difficulty   In order to overcome this difficulty instead of only NaCl, a mixture of NaCl and CaCl2 is electrolyzed in down's cell. The melting point of this mixture is 600C. At 600C no metallic fog is formed.
COMPOSITION OF CHARGE:         NaCl = 42%         CaCl2 = 58%
The process   When an electric current is passed through the molten mixture of NaCl and CaCl2, NaCl decomposes into Na+ and Cl- ion. Na+ ions migrate towards the cathode while Cl- ions towards the anode. The molten sodium collects in the cathode compartment where it rises to the top and is tapped off by a pipe. Chlorine is collected at the anode.
The chemistry of reaction Fused NaCl contains sodium and chloride ions.
2NaCl ---à 2Na+ + 2Cl-
Electrochemical changes
At cathode Na+-ions migrate to the cathode where they are reduced to Na.
2Na+ + 2e- ----à 2Na (Reduction) At anode
Cl--ions migrate to the anode and oxidized to form chlorine gas.
2Cl- -----à Cl2 + 2e- (Oxidation)

Overall reaction

2Na+ + 2e- -----à 2Na 2Cl- ------à Cl2 + 2e- __________________ 2Na+ + 2Cl-  -----à 2Na + Cl2

 

Properties of Sodium

Phys­i­cal prop­er­ties

Like oth­er al­ka­line met­als, sodi­um is a sil­very-white met­al. It is quite mal­leable and soft (metal­lic sodi­um can eas­i­ly be cut with a knife or scalpel, and a fresh cut shines in air). It has a high electroconductivity.

Chem­i­cal prop­er­ties of metal­lic sodi­um

Metal­lic sodi­um is usu­al­ly stored un­der a lay­er of kerosene, as it has high chem­i­cal re­ac­tiv­i­ty (Na can re­act vi­o­lent­ly with oxy­gen and air mois­ture even at room tem­per­a­ture). The pres­ence of sodi­um ions in salt can be de­ter­mined by the col­or of a flame – Na⁺ turns flame yel­low.

Re­ac­tion of sodi­um with oxy­gen

When sodi­um burns in oxy­gen, the main prod­uct of re­ac­tion will not be sodi­um ox­ide, but sodi­um per­ox­ide:

2Na + O₂ = Na₂O₂ (sodi­um ox­ide Na₂O also forms in this re­ac­tion, but in trace quan­ti­ties).

In air, sodi­um can also swift­ly ox­i­dize to sodi­um ox­ide (the ox­ide will not be pure, as a con­sid­er­able (some­times greater) quan­ti­ty of per­ox­ide Na₂O₂ will be present among the prod­ucts):

4Na + O₂ = 2Na₂O.

The re­ac­tion of sodi­um with oxy­gen may be writ­ten as fol­lows:

6Na + 2O₂ = Na₂O₂ + 2Na₂O.

Re­ac­tion of sodi­um with non-met­als

Sodi­um, be­sides com­bus­tion, is ca­pa­ble of tak­ing part in many oth­er re­ac­tions with non-met­als:

Re­ac­tion with hy­dro­gen with for­ma­tion of hy­dride:

2Na + H₂ = 2NaH (re­ac­tion takes place with heat­ing to 250-400 ᵒC (482-752 ᵒF) and at in­creased pres­sure);

Re­ac­tion with sul­fur with for­ma­tion of sul­fides:

2Na + S = Na₂S (sodi­um sul­fide);

Re­ac­tion with halo­gens:

2Na + Cl₂ = 2Na­Cl (sodi­um chlo­ride);

2Na + Br₂ = 2NaBr (sodi­um bro­mide);

Metal­lic sodi­um re­acts poor­ly with ni­tro­gen (the re­ac­tion may be con­duct­ed in a glow dis­charge – a burn­ing dis­charge formed in a low cur­rent and with low gas pres­sure):

6Na + N₂ = 2Na₃N.

 

Re­ac­tion of sodi­um with met­als

When the sur­face of metal­lic sodi­um con­tacts mer­cury, an amal­gam is formed – an al­loy of met­al with mer­cury.

With potas­si­um, the met­al forms an al­loy with the for­mu­la NaK. It is quite ag­gres­sive – it may com­bust in air. If the con­tent of potas­si­um in the al­loy varies from 40% to 90% in re­la­tion to sodi­um, the al­loy re­mains liq­uid at room tem­per­a­ture. The sodi­um-potas­si­um al­loy NaK is ob­tained by al­loy­ing liq­uid potas­si­um hy­drox­ide KOH with melt­ed metal­lic sodi­um. This re­ac­tion is car­ried out at a tem­per­a­ture of 400-450 ᵒC (752-842 ᵒF).

Sodi­um re­acts with wa­ter

Ow­ing to its chem­i­cal re­ac­tiv­i­ty, sodi­um may en­ter into re­ac­tions with var­i­ous com­pounds, for ex­am­ple:

Re­ac­tion of sodi­um with acids (di­lut­ed):

2Na + 2HCl = 2Na­Cl + H₂ (sodi­um chlo­ride NaCl and hy­dro­gen gas H₂ form).

Re­ac­tion with ni­tric and sul­fu­ric acids:

8Na + 8H₂­SO₄ = NaHS + 7NaH­SO₄ + 4H₂O (with heat­ing, con­cen­trat­ed acid);

8Na + 10H­NO₃ = 8NaNO₃ + NH₄NO₃ + 3H₂O (di­lut­ed acid – 3-5%);

11Na + 14H­NO₃ = 11­NaNO₃ + NO + N₂O + 7H₂O (acid of 20% con­cen­tra­tion).

Re­ac­tion with wa­ter (re­ac­tion takes place vig­or­ous­ly):

2Na + H₂O = 2NaOH + H₂.

Uses of Sodium

• It also used in improving the structure of certain alloys; soaps, purification of molten metals, and in sodium vapor lamps.
• Sodium is a component of sodium chloride, which is a very important compound found in the living environment.
• Sodium is important in the manufacturing of organic compounds and in making esters.
• Solid sodium carbonate is required in making glass.