" Metals and Non- metals"

Metals:-

• Elements which have the ability to donate electrons.
•  Example:- Lithium, Sodium, Potassium,Iron, Aluminium, Magnesium

Physical properties:- 

• Have shining surface . This property is called metallic lustre.
• Generally hard . Hardness vary from metal. Exception :- lithium, sodium and potassium are so soft that can be cut with a knife.
• All metals are solid at room temperature. Exception:- Mercury (Hg) - liquid at room temperature.
• Malleability:- Metals can be beaten into thin sheets . This property is called malleability.
• Ductility:- Metals can be beaten into thin wires . This property is called ductility.
• Good conductors of heat . Exception:- lead  (Pb)
• Have high melting points . Exception:- Gallium and Caesium have very low mp.
• Sonorous:- Produce a sound on striking a hard surface.
• Note:- Silver and copper are best conductors of heat.

PVC - Polyvinyl chloride.

• It is used to coat current carrying wires.

Non - metals:- 

• Elements which  don't have  the ability to donate electrons (or gain electron)
•  Example: carbon , nitrogen , oxygen , fluorine , sulphur , Iodine hydrogen etc.

Physical properties:- 

• Non metals are non lustrous . Exception :- Iodine
• Non metals are brittle .
• Exception :- Diamond ( Allotrope of carbon and hardest known substance)
• Non metals are either solids or gases at room temperature. Exception :- Bromine (liquid at room temperature)
• Non malleable
• Non ductile
• Bad conductor of heat and electricity. Exception:- Graphite (Allotrope of carbon)
• Low melting and boiling point. Exception:- Diamond.
• Non sonorous.

Chemical properties of metals:- 

• When metal burn in air:- 
• Almost all metals combine with oxygen to form metal oxides. Exception:- silver and gold.

Metal + Oxygen --> metal oxide
Cu + O2 -- > 2CuO (black )

4Al + 3O2 -- > 2Al2O3

• Metal oxides are either basic oxides or amphoteric oxides.
• Basic oxide:- Which react with acids only . Example:- CuO
• Amphoteric oxides:- Which react with base as well as acid to produce salt and water.
• Example:- Al2O3 , ZnO

Example:- 

Cu + O2 --> CuO

CuO + HCl --> CuCl2 + H2O

Al2O3 + 6HCl --> 2AlCl3 + H2O

Al2O3 + 2NaOH --> 2 NaAlO2 + H2O

• Most metal oxides are insoluble in water.
• Some dissolve in water to form alkali.
• Example:- Na2O and K2O dissolve in water to produce alkali.

Na2O (s) + H2O (l) --> 2NaOH (aq)

K2O (s) + H2O (l) --> 2KOH (aq)

• All metals do not react with oxygen at same rate 
• Different metals show different reactivities towards oxygen.
• Metals such as K and Na react so vigorously that they catch fire if kept in open. Hence to protect them and to prevent accidental fires, they are kept immersed in kerosine oil.
• At ordinary temperature, the surface of metals such as magnesium, aluminium, zinc and lead etc. Are covered with a thin layer of oxide. The protective oxide layer prevents the metal from further oxidation.
• Cu does not burn but hot metal is coated with a black coloured layer of copper (II) Oxide.
• Silver and gold don't react with oxygen even at high temperature.

Anodising :- process of forming a thick oxide layer of aluminium.

What happens when metals react with water:-

Metal + water -- > Metal oxide + hydrogen 

2K + 2H2O (cold) --> 2KOH + H2 + heat energy

2Na + 2H2O (cold) --> 2NaOH + H2 + heat energy.

Ca + 2H2O (hot) --> Ca(OH)2 + H2

Mg + 2H2O (hot) --> Mg(OH)2 +H2

Al/ Fe/ Zn + H2O (steam)--> Al2O3 /                                        Fe2O3/ZnO + H2

All metals don't react with water example:- Pb , Cu , Ag and Au

What happens when metals react with Acids?

Metal + dilute acid --> salt + hydrogen

Example:- 

Mg/Al/Zn/Fe + HCl (aq) --> MgCl2/AlCl3/ZnCl2/FeCl2 + H2(g)

Copper does not  react with Dil HCl

Note:- 

• H2 (g) is not evolved when metal reacts with nitric acid.
• Reason:- Because HNO3 is a strong oxidising agent .
• It oxidised H2 to H2O and itself get reduced to any of nitrogen oxides (N2O and NO2)
• But Mg and Mn react with very dilute HNO3 to evolve H2 gas.

Aqua regia :-

• A freshly prepared mixture of conc. HCl and conc. HNO3 in ratio of 3:1
• Can dissolve gold even t neither of the acids can do so alone.
• Properties:- Highly corrosive
• Fuming liquid

How do metals react with solutions of other metal salts?

Metal react with solutions of other metal salt by displacing them

Only more reactive metal (A) can displace less reactive metal (B) from its salt solution.

Metal A + salts Solution of B --> salt solution of A + metal B

A + BC --> AC + B

Reactivity series :-

• A series is a list of metals  arranged in order of their decreasing activities.

How do metals and Non-metals react?

The reactivity of elements is due to a tendency to attain a completely filled Valence shell.

They react either by gaining electrons (non- metals) or losing electrons (metals)

For example:-

 Na(11) 

Electronic configuration:- 2,8,1

Na --> Na+  +. e- 

          2,8 (cation formation)

Also :- 

Cl    +   e   -->.   Cl-

(2,8,7)               (2,8,8) anion 

                            Formation

• In NaCl , sodium give one electron to Chlorine so that both of them attain noble gas configuration for stability.
• So the compound formed by transfer of electron from a metal to a non metal are known as ionic compound / electrovalent compound
• In NaCl , Na+ and Cl- are held together by strong electrostatic forces of attraction.

Properties of ionic compounds/ electrostatic compound

:-

(i) Physical properties:-

• Solid
• Hard:- Due to electrostatic forces of attraction between cation and anion
• Brittle - break into pieces when pressure is applied.

(ii) Melting and boiling points:-

• High melting and boiling point.
• Reason:- Large amount of energy is required to break strong interionic attraction

(iii) Solubility :-

• Soluble in water
• Insoluble in other solvents such as kerosine, petrol etc.

(iv) Conduction of electricity:- 

• Ionic compounds conduct electricity in aqueous state or in molten state.
• This is because in aqueous and molten state ions move freely and conducting electricity.
• In solid states ions donor move freely . So don't conduct electricity.

Occurrence of metals:-

• Earth crust is the major source of metals.
• Sea water contain Soluble salts such as NaCl , MgCl2
• Minerals:- The element or compound which occur naturally in earth crust.

Ores:- 

• Minerals which contain at high % of particular metal and metal can be profitably extracted from it  
• Such minerals are called ores.

Gangue:-

•  Ore mined from earth are usually contaminated with large amount of impurities such as sand , soil etc.
• Such impurity is called gangue.

Extraction of metals:- 

• Metal at the bottom of reactivity series are of least reactivity.
• They are found in free state.
• Ex:- Gold (Au) , silver (Ag) , Platenum(Pt) and copper (Cu)
• Copper and silver are also found in combined state as their sulphide and oxide ores.
• Metals at the top of  activity series (K, Na, Ca, Mg and Al) are so reactive that are never found in nature as free elements
• Metals in middle of activity series (Zn,Fe,Pb etc) are moderately reactive.
• They are found in earth crust mainly as oxides , sulphides or carbonates)

Classification of metals:- 

On the basis of reactivity:- Three types

• Metals of high reactivity
• Metals of medium reactivity
• Metals of low reactivity

Extracting metals low in the activity series:- 

• Metals low in the reactivity series are very interactive.
• Oxides of these metals can be reduced to metal on heating.
• Example :- 2HgS + 3O2 + heat--> 2HgO + 2SO2(g)
• 2HgO + heat -- > 2Hg (l) + O2(g)
• (Mercuric oxide reduced to metal on further heating)
• Example :- 
• 2Cu2S + 3O2 + heat --> 2Cu2O (s) + 2SO2(g)
• 2Cu2O + Cu2S + heat --> 6 Cu (s) + SO2(g)

Extracting metals in middle of activity series (Fe, Zn ,Pb)

Roasting :- When sulphide ores are converted into oxides by heating strongly in the presence of excess air . This process is known as roasting.

Calcination:- When carbonate ores converted into oxides by heating strongly in limited air . This process is called calcination .

• Metals in middle of activity series are moderately reactive.
• Usually present as sulphide or carbonate in nature .
• It is easier to obtain metal from its oxide as compare to its sulphides and carbonates.
• So before reduction, metal sulphides and carbonates must be converted into metal oxides.
• Metal sulphide converted into metal oxide by roasting.
• Metal carbonate converted into metal oxides by calcination.
• For example:- By roasting :-
• 2ZnS(s) + 3O2 (g) + heat --> 2ZnO (s) + 2SO2(g)

Calcination:- 

ZnCO3 (s) + heat --> ZnO (s) + CO2(g)

Metal oxides are then reduced to metal by using suitable reducing agents such as carbon.

• ZnO (s) + C(s) --> Zn(s) + CO(g)

Beside using carbon to reduce metal oxides to metals, displacement reaction can also be used.

The highly reactive metal (Na,Ca,Al etc) are used as reducing agents because they can used as reducing agents because they can displace metal of lower reactivity from their compounds.

Example:- 3MnO2 (s) + 4Al --> 3Mn + 2Al2O3 + heat.

Displacement reactions are highly exothermic.

The amount of heat evolved is so large that the metal are produced in molten state.

Practical application :- 

Reaction of iron oxide (Fe2O3) with aluminium is used to join railway tracks or cracked machine parts.

This reaction is known as thermit reaction.

Fe2O3 (s) + 2Al (s) --> 2Fe (l) + Al2O3 (s) + heat

Extracting metals towards the top of activity series:-

• The metals high up in reactivity series very reactive.
• These metals are obtained by electrolytic reduction.
• In this , metals are deposited at cathode (negatively charged electrode) 
• Non metal are deposited at anode (positively charged electrode)
• To understand let us consider the example of NaCl
• At cathode :- Na+ + e- --> Na
• At anode :- 2Cl- --> Cl2 + 2e-
• Example 2:-
• At cathode :- 2Al3+ + 6e- --> 2Al
• At anode :- 6O2- --> 3O2 + 6e-

Refining of metals:- 

• Metals produced by various reduction processes are not very pure.
• They contain impurities which must be removed to obtain pure metals.
• The most widely used method for refining impure metals us electrolytic refining.

Electrolytic refining:-

• Many metals such as copper , zinc , tin , silver , nickel, gold etc. Are also refined electrically.
• In this process , the impure metal is made the anode and thin strip of pure metal is made the cathode.
• A Solution of metal salt is used as an electrolyte.
• On passing the current through the electrolyte, pure metal from anode dissolves into electrolyte.
• An equivalent amount of pure metal from electrolyte is deposited on cathode.
• Soluble impurity go into Solution.
• Insoluble impurity settle down at the bottom of anode. (anode mud)

Corrosion:- 

Example :- silver articles become black after some times when exposed to air.

Reason:- silver + sulphur in air --> black coat of silver sulphide.

Copper reacts with moist CO2 in air and slowly losses it's shiny brown surface and gain a green coat .

Reason:- copper + carbon dioxide --> copper carbonate

Iron react with moist air for a long time acquire a coat of a brown flaky substance called rust.

Reason:- Fe + moist air --> ferric oxide (rust)

Prevention of corrosion:- by painting, oiling , greasing , galvanising, chrome plating , anodising or making alloys.

Galvanisation:- 

• Method of protecting steel and iron from rusting. 
• This is done by coating then with a thin layer of zinc.
• Reason:- 
• Zinc is more reactive than iron and loses electrons more easily. 
• The outer layer of any  galvanised material react with atmospheric oxygen to form ZnO (zinc oxide) which is stronger than zinc.
• Thus even if the outer layer of zinc undergoes corrosion, the material is getting coated with a strong substance ZnO .
• Thus it is better able to resist corrosion.

Alloy formation (Alloying):-

• Good method to improve the property of metal.
• Can get desired properties by this method.
• Example:- Iron (widely used metal) never used in pure form.
• Because it is very soft and stretches easily when hot.
• On mixing carbon (0.05%) , it become hard and strong.

Stainless steel:-

•  When iron is mixed with nickel and chromium , we get stainless steel.
• It is hard and does not rust.
• Properties of any metal can be changed if it is mixed with some other substances.

Alloy:-

•  mixture of two or more metals or a metal and a non metal.
• Prepared by first melting the primary metal and then dissolving the other elements into it in definite proportions.
• It is then cooled to room temperature.

Note:-

•  If one of the metal us mercury, then the alloy is known as amalgam.

Properties of alloys:- 

• Electrical conductivity and melting point of alloy is less than that of pure metal.
• Example:- brass and bronze are not good conductors of electricity where as copper is used for making electric circuits.
• Solder has a low mp and is used for welding electrical wires together.

Alloys:-

Brass :- copper + zinc

Bronze :- copper + tin

Solder :- lead + tin

Fact about gold :- 

• Pure gold(24 carats) is very soft .
• Not suitable for making jewellery.
• For making it hard, it is Alloying with silver or copper.
• In India, 22 carat gold is used for making jewellery.
• It means 22 parts of pure gold is alloyed with 2 parts t either copper or silver.

Thankyou :-)