metals and non-metals
TRANSCRIPT
Standard/ Class/ Grade - X SSC, CBSE; - VIII ICSE
Metals and Non-metals
Gurudatta K Wagh
Elements
• Metals
• Non-metals
• Metalloids
Metalloids possess both the properties of metals as well as non-metals e.g. silicon Si, germanium Ge, antimony/ stibium Sb, etc.
Overview of the Chapter
Metals - Physical propertiesNon-metals - Physical propertiesMetals - Chemical properties
Reaction of metals with oxygenReaction of metals with waterReaction of metals with acidsReaction of metals with solutions of other metal saltsReaction of metals with non-metalsProperties of ionic compounds
Occurrence of metalsActivity seriesExtraction of metals of high reactivityExtraction of metals of medium reactivityExtraction of metals of low reactivityRefining of metals
Corrosion of metalsPrevention of corrosion
Metals - Physical properties
1. Solids: At room temperature. Exceptions: Mercury/ hydrargyrum Hg and gallium Ga are in liquid state at room temperature
2. Lustre: Shine in their pure state, polished to give a highly reflective surface
3. Malleability: Can be made into thin sheets
Metals - Physical properties
4. Ductility: Drawn into thin wires. Most ductile: Gold/ aurum Au, silver/ argentum Ag. One gram of gold can be drawn into a wire of 2 km length
5. Conduction of heat: Good conductors of heat; have high melting points. Silver Ag and copper/ cuprum Cu best conductors of heat. Lead/ plumbum Pb and mercury poor conductors of heat
6. Conduction of electricity: Good conductors of electricity. Electric wires are made of Cu having a outer insulation of PVC (polyvinyl chloride)
Metals - Physical properties
7. Hardness: Generally hard, differs from metal to metal. Alkali metals sodium Na and potassium K are soft metals and can be cut easily with a knife
8. Melting point and Boiling point: High. Tungsten/ wolfram W highest melting point; sodium Na and potassium K low melting points
9. Sonorous: Produce a sound on striking a hard surface e.g. bells
Non-metals - Physical properties
1. Either in solid or gaseous state. Exception: Bromine Br is in liquid state
2. Usually do not have lustre. Exception: Iodine I
3. Do not possess the property of hardness. Exception: Carbon C in the form of diamond. It is the hardest substance known which also has a high melting and boiling point
4. Do not conduct electricity. Exception graphite (allotrope of carbon)
Since there are exceptions in each of the physical
properties, elements can be classified more clearly as
metals and non-metals on the basis of their chemical
properties.
Metals - Chemical properties
• Reaction of metals with oxygen
Almost all metals react with oxygen to form metal oxide, but the reactivity differs for different metals
Sodium and Potassium are the most reactive metals. Sodium reacts with oxygen in air at room temperature to form sodium oxide. Hence, sodium is stored under kerosene oil to prevent its reaction with oxygen, moisture and carbon dioxide 4Na + O2 → 2Na2O
Most metal oxides are insoluble in water but
some of these dissolve in water to form alkalis, and are basic in nature
e.g. sodium oxide Na2O and potassium oxide K2O dissolve in water to produce alkalis Na2O + H20 → 2NaOH K20 + H20 → 2KOH
Magnesium does not react with oxygen at room temperature
but on heating, magnesium burns in air with intense light and heat to form
magnesium oxide MgO
2Mg + O2 heat → 2MgO
Zinc burns in air only on strong heating to form zinc oxide2Zn + O2 heat → 2ZnO
Iron does not burn even on strong heating, but iron filings burn vigorously when sprinkled in the flame of the burner3Fe + 202 heat → Fe304
Copper is least reactive. It does not burn, but on heating, the hot metal is coated with a black
coloured layer of copper oxide
2Cu + O2 heat → 2CuO
Aluminium develops a thin oxide layer when exposed to air
4Al + 302 heat → 2Al2O3
Anodising
Anodising is a process of forming a thick oxide layer of aluminium. Aluminium develops a thin oxide layer when exposed to air. This aluminium oxide coat makes it resistant to further corrosion. The resistance can be improved by making the oxide layer thicker
In this technique aluminium article is used as an anode. Electrolyte used is dilute sulphuric acid. The anode reaction results in formation of a black coloured thin film of aluminium oxide on the surface of anode
By putting appropriate dyes in the electrolytic bath, coloured surface with decorative finish can be achieved
Kitchen articles like anodised pressure cookers, anodised pans and also frames of sliding windows are the applications of anodising techniques
Amphoteric oxides
Metal oxides are usually basic in nature
but
some metal oxides such as aluminium oxide Al2O3 and zinc
oxide ZnO react both with acids as well as bases to
produce salt and water
Al2O3 + 6HCl → 2AlCl3 + 3H2O
Al2O3 + 2NaOH → 2NaAlO2 (sodium aluminate) + H20
• Reaction of metals with water
Metals like potassium and sodium react vigorously with cold water
Sodium reacts with water to evolve hydrogen which immediately catches fire producing a lot of heat (Exothermic reaction) 2K + 2H2O → 2KOH + H2 + heat energy
2Na + 2H2O → 2NaOH + H2 + heat energy
Reaction of Na with water
Calcium reacts with water less vigorously. The heat evolved is not sufficient for hydrogen to catch fire
Instead, calcium starts floating because the bubbles of hydrogen gas formed stick to the surface of the metal
Ca + 2H2O → Ca(OH)2 + H2
Magnesium reacts with hot water to form magnesium hydroxide Mg(OH)2 and hydrogen H2. Magnesium also starts floating since the bubbles of hydrogen gas stick to its surface.
Metals like aluminium, iron and zinc do not react either with cold or hot water but they react with steam to form metal oxide and hydrogen (g)
2Al + 3H2O → Al2O3 + 3H2
3Fe + 4H2O → Fe3O4 + 4H2
Metals like gold, silver, and copper do not react with water at all
• Reaction of metals with acids
Metals react with acids to give salt and hydrogen gas
Dilute hydrochloric acid HCl
Formation of metal chloride and hydrogen gas
The reactivity decreases in the order, Mg > Al > Zn > FeMg + 2HCl → MgCl2 + H2↑2Al + 6 HCl → 2AlCl3 + 3H2↑Zn + 2HCl → ZnCl2 + H2↑ Fe +2HCl → FeCl2 + H2↑
No bubbles are seen in case of copper. This shows that copper does not react with dilute hydrochloric acid
Sulphuric acid H2SO4
Formation of metal sulphate and hydrogen gas
Fe + H2SO4 → FeSO4(aq) + H2(g)Iron Fe reacts vigorously and exothermically with sulfuric acid to produce iron(II) sulfate
Zn + H2SO4 → ZnSO4(aq) + H2(g)
Mg + H2SO4 → MgSO4(aq) + H2(g)
Nitric acid HNO3
Hydrogen gas is not evolved when a metal reacts with nitric acid, as it is a strong oxidizing agent
It oxidizes the hydrogen to water and itself gets reduced to any of the nitrogen oxides (N20 nitrous oxide, NO nitric oxide, NO2 nitrogen dioxide)
But magnesium and manganese react with dilute HNO3 to evolve hydrogen gas
Mg(s) + 2HNO3(aq) → Mg(NO3)2(aq) + H2(g)magnesium + nitric acid → magnesium nitrate + hydrogen
Element Reaction with Air Reaction with Water Reaction with Acid Potassium K Sodium Na
Burns vigorously to form oxides
With cold water forms hydrogen gas and alkaline hydroxide solution. React with decreasing vigour down the series till Ca
Violent reaction to give hydrogen gas and salt solution
Calcium Ca Magnesium MgAluminium AlZinc ZnIron Fe
Burn with decreasing vigour down the series
Mg reacts with hot water to fomi magnesium hydroxide and hydrogen gas. For Al, Zn and Fe, no reaction with cold and hot water. With steam forms metal oxide and hydrogen gas
React to form hydrogen gas and salt solution with decreasing vigour down the series
Reactions of metals with air, water, and acids
Aquaregia
Aquaregia is a highly corrosive as well as fuming liquid and is one of the few reagents that is able to dissolve gold and platinum
It is a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in the ratio of 3:1
Aqua regia (Latin, lit. "royal water" or "king's water") was so named by alchemists because it can dissolve the noble metals gold and platinum
• Reactions of metals with solutions of other metal salts
Put an iron nail in the solution of copper sulphate
Fe(s) + CuSO4(aq) → Cu(s) + FeSO4(aq)
metal A + salt solution B → salt solution A + metal B
The iron nail gets coated with a reddish brown colour copper and the blue colour of copper sulphate solution fades out
In this reaction more reactive iron has displaced copper which is less reactive from the copper sulphate solution
This reaction is known as displacement reaction. The brown coating on the iron nail shows that copper is deposited on the iron nail by displacing iron
The greenish colour of the solution in the test tube shows that Fe2+ ions are present in the solution. This shows that iron is more reactive than copper, as Fe2+ ions have displaced Cu2+ ions from copper sulphate solution
This is a single displacement reaction in which copper has been displaced from iron from copper sulphate solution
A more reactive metal can displace a less active metal from its compound in a solution
• Reaction of metals with non-metals
Let us look at the electronic configuration of some metals and non-metals. It will help us to understand the reactivity of different elements and formation of different compounds
We have learnt that noble gases have a completely filled valance shell, hence are chemically inactive
Electronic configuration of some metals and non-metalsType of element
Element Atomic number
Electronic configuration
K L M N
Metals Sodium (Na)Magnesium (Mg)Aluminium (Al)Potassium (K) Calcium (Ca)
l l 12 13 19 20
22222
88888
12388
12
Non-metals Nitrogen (N) Oxygen (O) Fluorine (F) Phosphorus (P) Sulphur (S) Chlorine (Cl)
78 9 15 16 l7
222222
567888
567
Sodium, a metal, is a silver coloured metal that reacts so violently with water that flames are produced due to formation of hydrogen gas
Chlorine, a non-metal, is a greenish coloured gas which is so poisonous that it was used as a weapon in World War l When chemically bonded together, these two dangerous substances form a compound sodium chloride so safe that we eat it every day (common table salt)
Sodium (2, 8, 1) atom has one electron in its outermost shell. If it loses one electron from its "M" shell then its "L" shell becomes the outermost shell to acquire a stable octet. The nucleus of this atom still has 11 protons but the number of electrons has become 10, so there is a net positive charge giving us a sodium cation (Na+)
On the other hand chlorine (2, 8, 7) has 7 electrons in its outermost shell and requires one more electron to complete its octet
The electron lost by sodium is taken up by chlorine. After gaining one electron, its K, L and M shells have altogether 18 electrons, but the nucleus still has 17 protons
This leads to the formation of chloride anion (Cl-). Both these elements have a give and take relation between them
Ionic or electrovalent bond/ compound
Sodium and chloride ions being oppositely charged attract each other and are held by strong electrostatic forces of attraction to exist as sodium chloride (NaCl), resulting in formation of an electrovalent bond or an ionic bond
Sodium chloride exists as aggregates of oppositely charged ions in definite geometrical shape
Such a bond formed by the give and take of electrons is called as ionic or electrovalent bond. Compounds formed in this manner by the transfer of electrons from a metal Na to a non-metal Cl are known as ionic compounds or electrovalent compounds
Covalent compound
The chemical compound that is formed by mutual sharing of one or more pairs of electrons between the two combining atoms is called a covalent compound e. g. water, ammonia, etc.
Properties of ionic compounds
1. Ionic compounds are solids and hard due to strong force of attraction between positive and negative ions
2. They are generally brittle and break into pieces when pressure is applied
3. Ionic compounds have high melting and boiling points, as a considerable amount of energy is required to break the strong inter molecular attraction
4. They are generally soluble in water and insoluble in solvents such as kerosene, petrol, etc.
5. Ionic compounds in the solid state do not conduct electricity because the movement of ions in the solid state is not possible due to their rigid structure, but they conduct electricity in the molten state. The conduction of electricity through a solution involves the movements of charged particles. A solution of an ionic compound in water contains ions, which move to the opposite electrodes when electricity is passed through the solution
Occurrence of metals
The most unreactive metals (not affected by air and water) e. g. silver, gold and platinum are generally found in free or native state
Most metals however are found in combined state in the form of their oxide ores, carbonate ores, or sulphide ores, etc.
Minerals The naturally occurring compounds of metals along with other impurities are known as minerals
Ores The minerals from which metals are extracted profitably and conveniently are called as ores
Gangue Ores contain metal compounds with some of the impurities like soil, sand, rocky material, etc. These impurities are called as gangue
Metallurgy Metals can be extracted from the ores by employing different separation techniques. The process used for extraction of metals in their pure form from their ores is called metallurgy
Activity series and related metallurgy
K Na Ca Mg Al Zn Fe Pb Cu Hg Ag Au
Decreasing order
of reactivity
Top
↓
↓
↓
↓
↓
Bottom
The arrangement of metals in the decreasing order of their reactivity in the form of series is called the reactivity series of the metals
Thus the most reactive metal is potassium and is placed in the top of the list and least reactive metal is gold which is placed at the bottom of the list
Three categories of metals
*Metals of high reactivity
*Metals of medium reactivity
*Metals of low reactivity
• Extraction of metals of high reactivity
Metals high up in the reactivity series are very reactive e.g. sodium, potassium, calcium, aluminium, etc. These metals are obtained by electrolytic reduction
Sodium, magnesium and calcium are obtained by electrolysis of their molten chlorides
The metals are deposited at the cathode (-vely charged electrode), whereas, chlorine is liberated at the anode (+vely charged electrode). The reaction of sodium is as follows: At cathode : Na+ + e- → Na At anode : 2Cl- → Cl2 + 2e-
Extraction of aluminium Al (silvery white, atomic no. 13; electronic configuration 2,8,3; valency 3)
Aluminium is a reactive metal. Al is extracted from its main ore bauxite (Al2O3.H2O). It contains 30 % - 70 % aluminium oxide Al203
Aluminum is obtained by the electrolytic reduction of aluminium oxide
The remaining portion is gangue made of sand, silica (SiO2), iron oxide (Fe2O3) etc.
Extraction of Al involves two steps
Step 1 Concentration of ore i.e. conversion of bauxite into alumina (by Bayer's process)
* Crude bauxite contains impurities like iron oxide Fe2O3
and silica SiO2. These impurities are removed by Bayer's process. In this process, the ore is first crushed and then treated with hot concentrated caustic soda (NaOH) solution under high pressure for 2-8 hours at 140 °C to 150 °C in a tank called digester. Aluminium oxide being amphoteric in nature dissolves in aqueous sodium hydroxide to form water soluble sodium aluminate Al2O3 + 2NaOH → 2NaAlO2 + H2O
The iron oxide in the gangue does not dissolve in aqueous sodium hydroxide and is removed by filtration. However silica from the gangue dissolves in aqueous sodium hydroxide forming water soluble sodium silicate. Diluting sodium aluminate with water and then cooling to 50 °C, it is hydrolysed to give aluminium hydroxide as a precipitate NaAlO2 + 2H2O → NaOH + Al(OH)3↓
The precipitate is filtered, washed, dried and ignited at 1000 °C to get Alumina (Al2O3) 2Al(OH)3 heat→ Al2O3 + 3H2O
Step 2 Electrolytic reduction of alumina
In this process a molten mixture of pure alumina (MP > 2000 °C) is electrolysed in a steel tank. This tank is lined inside with carbon ( graphite) which acts as a cathode, and a set of carbon (graphite) rods dipped in the molten electrolyte act as anode. Cryolite (AlF3-3NaF) and fluorspar (CaF2) is also added to the mixture to reduce the melting point to about 1000 °C
On passing the current, aluminium is formed at the cathode. The molten aluminium being heavier than the electrolyte used, sinks to the bottom of the tank from where it is removed periodically. On the other hand oxygen is liberated at the anode
The electrode reactions are:Anode: 2O2- - 4e- → O2
Cathode: Al3+ + 3e- → Al
The oxygen gas liberated, reacts with carbon anode and forms carbon dioxide. As the anode gets oxidized during the electrolysis of alumina, it has to be replaced from time to time
•Extraction of metals of medium reactivityThe metals in the middle of activity series such as iron, zinc, lead, copper, etc. are moderately reactive. These are usually present as sulphides or carbonates in nature. It is easier to obtain a metal from its oxide, as compared to its sulphides and carbonates. The sulphide ores are first converted into oxides by heating strongly in excess of air. This process is known as roasting. The carbonate ores are changed into oxides by heating strongly in limited air. This process is known as calcination. The chemical reaction that takes place during roasting and calcination of zinc ores is: Roasting 2ZnS + 302 → 2ZnO + 2SO2 Calcination ZnCO3 → ZnO + C02
This zinc oxide is then reduced to zinc by using a suitable reducing agent such as carbon ZnO + C → Zn + CO
Besides using carbon to reduce metal oxides to metals, highly reactive metals like sodium, calcium, aluminium, etc. are also used as reducing agents because they can displace the metal of lower reactivity from their compounds. e.g. when manganese dioxide is heated with aluminium powder, the following reaction takes place 3MnO2 + 4Al → 3Mn + 2Al203 + heat
The amount of heat evolved in such reactions is so large that the metals are produced in the molten state. Another example of this type of reaction is the thermit reaction in which iron oxide reacts with aluminium to give iron and aluminium oxide evolving lot of heat Fe2O3 + 2Al → 2Fe + Al203 + heat
•Extraction of metals of low reactivity
The metals at the bottom of the activity series are least reactive. They are often found in free state e.g. gold, silver and copper. But copper and silver are also found in combined state as their sulphide or oxide ores. e. g. copper which is found as Cu2S in nature can be obtained from its ore by just heating in air 2Cu2S + 3O2 heat→ 2Cu2O +2SO2
2Cu2O + Cu2S heat→ 6Cu + SO2
Similarly, cinnabar (HgS) is an ore of mercury. When it is heated in air, it is first converted into mercuric oxide (HgO). Mercuric oxide is then reduced to mercury on further heating
heat 2HgS + 3O2 heat→ 2HgO + 2SO2 heat 2HgO heat→ 2Hg + O2
Refining of metalsThe metals produced by various reduction processes described above are not very pure. They contain impurities, which should be removed, to obtain pure metals. The most widely used method for refining impure metals is electrolytic refining
Corrosion of metalsCorrosion is degradation of materials due to reaction with its environment. The major problem of corrosion occurs with iron, since it is used as a structural material in constructions, bridges, ship building, automobile etc.
Have you ever observed*Old iron grills in the buildings?*Copper vessels which are not cleaned for a long time? *Silver ornament or idols which are kept exposed to air for a long time?*Old worn out cars?*Iron reacts with moist air to acquire a coating of brown flaky substance called rust. *Copper reacts with moist carbon dioxide with air and slowly loses its shine to gain a green coat of copper carbonate.
*Silver articles become black after some time when exposed to air because it reacts with hydrogen sulphide in the air to form a coating of silver sulphide
*Aluminum undergoes an oxidation reaction forming a thin layer of aluminum oxide, which acts as a barrier to oxygen and water preventing the further oxidation and resistant to further corrosion. This resistance can be improved by making the oxide layer thicker. example anodizing
Prevention of corrosion
Corrosion of metals can be prevented if the contact between metal and air is cut off. This is done in a number of ways. Some of the methods are given below:
(1) Corrosion can be prevented if the metal is coated with something which does not allow moisture and oxygen to react with it(2) Coating of metals with paint, oil, grease or varnish prevents the corrosion of metals e. g rusting of iron can be prevented by this method
Coating of corrosive metals with non-corrosive metals also prevents corrosion. Some of the methods by which metals can be coated with non-corrosive metals are:
Galvanizing: It is process of giving a thin coating of zinc on iron or steel to protect them from corrosion. e.g. shiny iron nails, pins, etc.
Tinning: It is the process of giving a coating of tin, i.e., molten tin over other metal. e.g. Cooking vessels made of copper and brass get a greenish coating due to corrosion. This greenish coating is poisonous. Therefore they are given a coating of tin to prevent corrosion (Kalhai)
Electroplating: In this method a metal is covered with another metal using electrolysis. Silver-plated spoons, gold-plated jewellery, etc. are electroplated
Anodizing: In this method metals like copper and aluminum are electrically coated with a thin strong film of their oxides. This film protects the metals from corrosion
Alloying: An alloy is an homogenous mixture of two or more metals or a metal and a non-metal in definite proportion. The resultant metals called alloys do not corrode easily, e.g. stainless steel. Examples, Brass (copper and zinc) Bronze (copper and tin) Stainless steel (iron, nickel and chromium, carbon )
If one of the metals is mercury, then the alloy is known as an amalgam
Pure gold, known as 24 carat gold, is very soft. It is, therefore, not suitable for making jewellery. It is alloyed with either silver or copper to make it hard. Generally, in India 22 carat gold is used for making ornaments. It means that 22 parts of pure gold is alloyed with 2 parts of either copper or silver
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