Metals



Properties of Metals


Physical Properties

  • shiny
  • dense
  • malleable and ductile
  • generally high m.p. and b.p.
  • excellent conductors of heat and electricity
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Chemical Properties

  • form positive ions (cations)
  • metal oxides tend to be basic (also amphoteric)
  • metal hydroxide also follow above
  • react with acid to form hydrogen gas + a salt

Reactivity Series

Chemically some metals are more reactive than others. This can be deduced using the reactivity series of metals, which is an arrangement of metals (and carbon and hydrogen) from the most reactive to the least reactive.

The extent of reactivity of a metal is considered by its ability to form a positive ion. The reactivity series of metals was determined by experiments in laboratories involving the observation of different metals reacting with water and steam, dilute acid and oxygen.

1. Reactions with Air (Oxygen)

Most metals will react with air to form a metal oxide (these are either basic or amphoteric). The most reactive metals like potassium, sodium and magnesium will burn with a very bright flame and will tarnish quickly in open air. Moderately reactive metals such as magnesium, aluminium, zinc and iron burn in strong flame but don't tarnish rapidly in open air. Iron as we know rusts to give a red-brownish compound (Fe2O3). The least reactive metals such as lead and copper show oxidisation in strong heat but show little tarnishing in open air. Metals below lead and copper in the reactive series are considered to not react with oxygen, e.g. gold and silver.

2. Reactions with Water

We can divide reactions between metals and water into two categories - hot and cold water.

Some metals such as potassium and sodium are so reactive they will react vigorously with cold water to form metal hydroxides and hydrogen gas. For the moderately reactive metals they will only react with steam to produce the appropriate oxide and hydrogen gas. The less reactive metals do not react with water or steam.


2K + 2H2O = 2KOH + H2

Mg + H2O = MgO + H2

3. Reactions with Dilute Acids

Most reactive and moderately reactive metals react with dilute acid with decreasing. An exception is that of aluminium which takes longer to ‘eat through’ the protective oxide layer. Very unreactive metals show no or very slight activity. Lead reacts very slowly with acid and stops eventually while copper does not react at all. It is this property of copper and metals below it in the reactivity series that enable people to clean coins and jewellery with dilute acids.

Note that aluminium appears to be very unreactive. This is because the surface of a piece of aluminium oxidises very quickly and becomes a protective cover preventing further reaction. Other metals oxidise like this, such as iron(III) oxide (rust) but don't protect the metal underneath as the oxide layer is usually porous.  

Displacement of Metal Ions

The order of reactivity can be used to predict what will happen when a metal is place in a solution of another metal. For example, when magnesium ribbon is place in a solution of copper(II) sulfate, the blue colour of the copper ion fades fast and a reddish-brown coating appears on the surface of the zinc copper. Displacement occurs.

CuSO4(aq) + Mg(s) = MgSO4 + Cu(s)
pale blue           colourless     red-brown

Observations like these with different chemical leads to a general rule:

·      Any metal will displace another metal lower on the reactivity series from a solution of its ions.
This does not apply to any metals that react with water, so can only be applied to magnesium and below.

Using the displacement theory above we can also include carbon and hydrogen in the reactivity series. Hydrogen ions displace lead but not copper so it must be in the middle of the two. Carbon (placed between aluminium and zinc) is able to reduce metal oxides with heat.

We can now finally give a full complete reactivity series:

                                                                          most reactive
                                                                                                                                                                 K
Na
Ca       
Mg
Al
(C)
Zn
Fe 
Pb
(H)
Cu
Ag

                                                                        least reactive

                                                                                                                       

Effect of Heat on Metal Compounds

Applying heat to some metal compounds causes the compound to break apart. This is known as thermal decomposition.

For the IGCSE CIE course the following needs to be memorised:

Ions of highly reactive metals hold on very strongly to their anions and therefore do not decompose very easily (require lots of heat). In general the less reactive the metal cation, the more it decomposes in heat.











Extraction of Metals

The ease of obtaining metals from their ores depends on the metals position in the reactivity series. An ore is a mixture of mineral containing a metal compound within it.

Less reactive metals are easily removed from their ores by strong heating, slightly more reactive metals can be removed by reduction using heated carbon. Reactive metals are difficult to obtain, as they require huge amounts of energy to decompose, instead we use electrolysis. The very unreactive metals such as gold and silver are found in pure elemental form. Copper is a small exception, it is heated to decompose it and then electrolysed to ensure maximum purity for use as a conductor.

A summary of above is below:
















Extraction of Iron

Iron extraction is carried out with the help of a blast furnace within which are the following raw materials:







The following is a step by step method to extract iron.

1. The raw materials are fed into the top of the blast furnace

2. Hot air is blown through the bottom of the furnace and oxygen reacts with coke to produce carbon dioxide

C + O2 = CO2

3. The CO2 produced rises into the upper chamber of the furnace and here reacts with hot coke again to produce carbon monoxide

C + CO2 = 2CO

4. The carbon monoxide (CO) rises and acts as a reducing agent and reduces the iron ore to iron.

Fe2O3 + 3CO = 2Fe + 3CO2

Since this part of the extraction takes place at around 700oC the iron produced is molten and ‘trickles’ to the bottom of the furnace.

5. The limestone (calcium carbonate) decomposes to form a calcium oxide and carbon dioxide. The calcium oxide is useful as it is a base and reacts with the acidic impurities such as silicon oxides to form a slag (mainly calcium silicate).

CaCO3 = CaO + CO2
CaO + SiO2 = CaSiO3 (calcium silicate)

The slag trickles down to the bottom of the furnace and float on top of the molten iron since it is less dense. It is used to make roads surfaces.

Note:

The molten iron as well as molten slag may be tapped off at regular intervals
the waste gases (CO2, etc.) escape from the top of the furnace. Since they are very hot they are used in a heat exchange process to heat incoming air, therefore making the process more economical. 

Conversion of Iron into Steel

Iron produced using the blast furnace method is called pig iron. It is still impure as it consists of about 5% of carbon, silicon, sulfur and phosphorus. The impurities make it brittle and quite undesirable.

This can be fixed by decreasing the percentage of impurities in it and adding other metals it with other elements. When this is done we have the product steel, which is an alloy (studied next). There are a huge number of different steels each one varying in properties due to difference in elemental composition.

Making steel can be done using another furnace called the oxygen furnace:

1. Molten pig from the blast furnace is poured into the furnace

2. Air (oxygen is introduced), which reacts with the impurities to form gaseous oxides which escape

Carbon oxidises into carbon monoxide and carbon dioxide
Sulfur oxidises into sulfur dioxide
Silicon and phosphorus oxidise into solids and are further reacted with calcium oxide (from decomposition of limestone) and skimmed off the top of the molten iron.

3. Certain metals responsible for certain properties are appropriately added to the molten iron. The oxygen bubbled in the molten also serves as a mixer ensuring an even mix.

4. Constant checking on the percentage of various elements is carried out and when the desired levels are reached the blast of oxygen is halted and the molten mixture is then poured out.

Extraction of Zinc

Zinc exists as zinc sulphide (ZnS) in the ore zinc blende. It is extracted from its ore by the following process:

1. The zinc ore, zinc blende is concentrate by froth flotation. That is, the crushed ore is and dumped into tanks of water containing special froth chemicals. Air is then blown through the water to create froth. The rock particles (i.e. useless particles) become soaked with water and sink to the bottom of the tank while the zinc sulphide floats and is carried to the top. Here it is skimmed off.

2. The zinc sulphide is now fed into a furnace where it is roasted with oxygen so it converts into zinc oxide (ZnO). This step removes the sulfur.

2ZnS + 3O2 = 2ZnO + 2SO2

3. The zinc oxide is taken and mixed with powdered coke at temperatures of around 1400°C. The carbon in the coke reduces the zinc oxide to zinc.

ZnO + C = Zn + CO

The carbon monoxide escapes from the furnace but since it is hot it can be used to heat incoming air through a heat exchange system to make the process more economical. 

Carbon monoxide can also be oxidised further to release energy.

Extraction of Aluminium

See the full process in chapter 5, but for summary:
  • bauxite (main ore of aluminium oxide) dissolved in cryolite and place in electrolytic cell consisting of carbon electrodes

  • mixture heated with current, (cyrolite decreases this temperature)
  • mobile ions now present and at the anode oxygen is liberated while at the cathode aluminium is reduced and accumulates in the bottom of the cell to be tapped off.
  • oxygen produced by oxidation of oxygen ions reacts with carbon cathode to produce carbon dioxide, therefore the anode has to be replaced regularly

Note that this process requires HUGE amounts of energy and this is normally the reason why aluminium smelters are normally located near hydroelectric power plants – providers of cheap electricity. 


Alloys

The majority of the metallic substances that we use today are alloys. Alloys are mixtures of two or more metals as well as other elements. They are so useful because of their properties coming from the combination of their components. The properties of alloys can be altered by changing their percentage composition. Mixing molten metals thoroughly forms alloys.

Steel

Out of all the alloys we use, steel probably has to be the most important. There are many variations of steel. One thing in common is that all have varying amounts of iron and carbon as well as other elements. We know from previously how the different variations of steel are produced. There are two main variations we need to know:

i) Mild Steel

Mild steel is 99.5% iron and 0.5% carbon. Mild steel is relatively easily worked and loses most of its brittleness. It therefore is used in car body manufacturing and machinery.

ii) Stainless Steel

Stainless steel is probably the variation that most people are aware of. It is composed of 74% iron, 18% carbon and 8% nickel. It is very resistant to corrosion, hard and attractive. The chromium prevents the steel from corroding, while the nickel makes it very hard. Stainless steel have a wide application from cutlery to equipment in chemical plants.

Brass

Brass is an alloy of zinc (5-40%) and copper (60-95%). It is used for:
heat exchangers as brass is very good conductor of heat, instruments & jewellery as brass is very workable and has acoustic abilities, zippers, taps, etc. as it is corrosion resistant