As we know that elements on the basis of their properties are classified into metals and non-metals.
We can differentiate between metals and non-metals only by differentiating its physical and chemical properties. Let’s start with physical properties.
Physical properties of Metals and Non-Metals:
|Lusture||Have shining surface.
|Do not have
Except: Sodium(Na) and
can be cut by a knife
are commonly soft.
The hardest natural
Except: Mercury(Hg) – Liquid
|Solid or Gaseous
|Malleability||Metals can be beaten in
Gold(Au) and silver(Ag)
– Most Malleable metals
|Ductility||Metals can be drawn to
|They are non-ductile.|
|Conductivity||Metals are good conductors
of Heat and Electricity
Silver (Ag) and Copper (Cu)
are best conductors.
Lead (Pb) and Mercury (Hg)
are poor conductors of Heat.
Except: Na and k
|Sonorous||Metals when stricken produce
Chemical properties of metals and non-metals are listed below:
Chemical properties of metals:
1. Reaction of metal with oxygen (air):
Metal + Oxygen → Metal Oxide
Metal when reacted with oxygen forms metal oxide.
Example: Copper combines with air when heated forms copper oxide, a black oxide.
2Cu + O2 → 2CuO
- The tendency of different metals to react with oxygen is different.
- When Sodium and potassium are kept in open they catch fire instantly. So they are kept in kerosene oil.
- Some of the metals such as Al, Mg, Zn, Pb are coated with a thin layer of oxide which prevents them from further oxidation.
- Copper on heating doesn’t burn but produces copper oxide on its surface.
- Gold(Au) and Silver (Ag) don’t show any reaction with oxygen.
- Iron doesn’t burn on simple heating but iron powder do.
Amphoteric Oxide: Those metal oxides which on reaction with acid as well as base produces salt and water are amphoteric oxides.
As we see in case of Aluminium oxide
- Al2O3 + 6HCl → 2AlCl3 + H2O
- Al2O3 + 2NaOH → 2NaAlO2 + H2O
2. Reaction of Metal with Water:
Metal + Water → Metal Oxide + Hydrogen
The Metal reacts with water to form metal oxide and release hydrogen gas. Metal oxide (soluble in water) further reacts with water to form metal hydroxide.
Metal Oxide + Water → Metal Hydroxide
- 2Al(s) + 3H2O (g) → Al2O3(s) + 3H2(g)
- 2K(s) + 2H2O(l) → 2KOH(aq) + H2(g) + ∆
- 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) + ∆
Low reactive metals such as lead, copper, silver and gold do not react with water at all.
3. Reaction of metal with acid:
Metal + Dilute acid → Metal Salt + Hydrogen
Metals react with acids to give a salt and hydrogen gas.
- Mg + 2Hcl → MgCl2 + H2
- Fe + 2HCl → FeCl2 + H2
- Hydrogen gas is not evolved when a metal reacts with nitric acid.
- Copper, Silver and Mercury do not react with dilute acids.
4. Reaction of metal with salt solution:
Metal A + Salt solution of B → Salt solution of A + Metal B
If and only if Metal “A” is more reactive then Metal “B”.
- CuSO4 + Zn → ZnSO4 + Cu
- Fe + CuSO4 → FeSO4 + Cu
- When metal reacts with salt solution than more reactive metal displaces the less reactive metal from its salt solution.
- It is a type of displacement reaction.
5. Reaction of metal with chlorine:
Metal + Chlorine → Metal Chloride
Metal reacts with chlorine to form metal chloride thus forms ionic bond. Therefore, we get ionic compound.
- 2Na + Cl2 → 2NaCl
- 2K + Cl2 → 2KCl
6. Reaction of metal with Hydrogen:
Metal + Hydrogen → Metal Hydride
Some Metals react with Hydrogen to form Metal Hydrides. Only most reactive metals take part in the above-stated reaction.
Example: 2Na + H2 → 2NaH
The list of metals and non-metals arranged orderly about their decreasing activity is said to be reactivity series.
As we know that noble gases have completely filled valence electrons. Therefore it shows less chemical activity.
By this property, we can define reactivity as the tendency of any element to attain noble gas configuration.
Electronic Configuration of some elements (metals and non-metals):
|Types of Element||Element||Atomic Number||Electronic
|Noble Gases||· Helium(He)
2, 8, 8
|2, 8, 1
2, 8, 2
2, 8, 3
2, 8, 8, 1
2, 8, 8, 2
2, 8, 5
2, 8, 6
2, 8, 7
Reaction of Metals with Non-Metals:
Sodium is a metal and has 1 electron in its outermost shell. When it loses 1 electron from its outermost shell the previous shell has the complete octet. Thus, It attains a noble gas configuration and is stable. It has now 11 protons and only 10 electrons thus carry a positive charge i.e Na+(Sodium cation).
In case of Chlorine, It has 7 electrons in its outermost shell. To attain noble gas configuration it has either to gain 1 electron or to lose 7 electrons. Now, losing 7 electrons simultaneously is the tougher task then gaining 1 electron. Thus, Chlorine gains 1 electron and attains noble gas Configuration. Now it has 18 electrons and 17 protons. Which makes it carries negative charge Cl–.
Na (2.8.1) → Na+ (2.8) + 1e–
Cl (2.8.7) + 1e– → Cl– (2.8.8)
Thus by this way Sodium chloride (NaCl) is formed.
Properties of Non-metals:
Non-metals have distinct properties when compared to metals. Halogen group and noble gases are usually non-metals.
Properties of non-metals are :
- They don’t have a shiny surface (Non-lustrous) except iodine.
- Solid non-metals are generally soft except diamond.
- Always in solid or gaseous state except for bromine (liquid).
- They can’t be drawn to thin wires (non-ductile).
- Non-metals can’t be beaten into thin sheets (non-malleable).
- They are generally bad conductor of electricity (except graphite).
- when stricken hard don’t produce sound like metals.
- It has high ionization energies.
- Also, it has high electronegativities.
- Non-metals gain electrons easily.
- These are poor conductors of heat.
- Have lower melting and boiling points when compared to metals.
The compound which is formed by transfer of electrons from metals to non-metals is said to be ionic compounds or electrovalent compounds.
Properties of Ionic Compounds:
- Physical Appearance:
Ionic compounds are solid.
Strong Inter-ionic force of attraction.
Generally brittle thus broken into pieces when pressure applied.
- Melting and Boiling Points:
Ionic compounds have high melting and boiling points.
Electrovalent compounds are commonly water soluble.
Ionic compounds conduct electricity in molten state (Solution in water) but not in solid state.
Extraction of Metals:
- Minerals: Minerals are the elements or compounds, which occur naturally in the earth’s crust.
- Ores: Some minerals contains high percentage of certain metals. When from these minerals if metals can be profitably extracted then these minerals are called ores.
Properties of ores of different metals:
- Metals which lies in the bottom of reactivity series mostly occur in free-state thus doesn’t require extra effort to extract. Ex: Silver and gold.
- (K, Na, Ca, Mg and Al) are the metals at the top of the activity series and is so reactive that they are never found in nature as free elements.
- The metals in the middle of the activity series (Zn, Fe, Pb, etc.) are moderately reactive. Their oxides, sulphides or carbonates are found.
Extraction of metals from ores on basis of reactivity:
Ores, when mined from the earth, are contaminated with large amounts of impurities such as sand, soil etc., called gangue. To extract metal we first remove these impurities.
- Extraction of low reactive metals:
When these metals are in oxide/sulphide form, it is reduced simply by Heating.
Example: (Mercury) 2HgO + Heat → 2Hg + O2
- Extraction of medium reactive metals:
Oxide ores are easy to reduce, thus carbonate and sulphide ores are converted into oxide ores then reduced further.
- Roasting: when sulphide ores are converted into oxides by heating strongly in the presence of excess air. This process is known as roasting.
Example: 2ZnS + 3O2 + Heat → 2ZnO + 2SO2
- Calcination: when carbonate ores are changed into oxides by heating in limited supply air. This process is known as calcination.
Example: ZnCO3 + Heat → ZnO + CO
- Reduction: The metal oxides thus formed, then reduced to the corresponding metals by using suitable reducing agents such as carbon. This process is known as reduction.
Example: ZnO + C → Zn + CO
3. Extraction of High reactive metals:
We obtain these metals from its ores by electrolytic reduction. Electrolysis is the only way to reduce such ores to metals. Respective ions settle at cathodes and anodes.
Example: sodium, magnesium, calcium, aluminium, etc., we obtain these metals by electrolytic method of reduction.
Refining of metals:
After reduction itself, metals obtained are not fully pure. These are impure metals and we carry refining process to make it pure.
Mostly we use Electrolytic Refining.
- Impure metal at anode.
- Pure metal attached with cathode.
- Acidified sulphate solution as electrolyte
Mechanism of refining:
- When current is passed through electrolyte, The impure metal falls and dissolves in electrolyte.
- Then equal amount of pure metal from the electrolyte is deposited at cathode.
- The insoluble impurities Settles in the bottom of the anode strip thus, we say it anode mud.
When metals are exposed to air, it forms a coat of layer on its surface which is mainly oxide (sometimes sulphides) is known as corrosion.
- Generally, we see that iron acquires a brown coating when it is kept in open moist air. This is rust.
- Silver jwelleries becomes black as it forms sulphide layer on its surface in open air.
- Copper utencil’s suface becomes green as it forms copper carbonates on its surface.
Prevention of Corrosion:
We can prevent Corrosion by:
- Painting the surface of metals
- Galvanising (applying zinc coat)
- Chrome Plating
- Anodising or making alloys
- Corrosion of iron is prevented by painting, oiling, greasing.
- Galvanization: We apply a thin coat of zinc on steel surface to save from corrosion.
- Alloys: Different metals (with metals or non-metals) when homogenously mixed improves the property of metal further also prevents rusting.
Some common alloys:
- Steel (Fe + Ni & Cr)
- Brass (Cu + Zn)
- Amalgam (Hg + any metal)
- Bronze (Cu + Sn)
- Solder (Pb + Sn)
- Iron is much soft so we mix it with little carbon (0.05%) to make it hard.