Chemical Properties
Chemical Properties of Metals
● Alkali metals (Li, Na, K, etc.) react vigorously with water and oxygen or air.
● Mg reacts with hot water.
● Al, Fe
and Zn react with steam.
● Cu, Ag, Pt, and Au do not react with water or
dilute acids.
Reaction of Metals with Oxygen (Burnt in Air)
A metal oxide is formed when metals are burned in the air and react with oxygen in the air. Metal oxides are a type of basic material found in nature. They change the colour of red litmus to blue. To avoid reactions with oxygen, moisture, and carbon dioxide in the air, sodium and potassium metals are kept in kerosene oil.
Metal + Oxygen→ Metal oxide (basic)
● Na and K are kept immersed in
kerosene oil as they react vigorously with air and catch fire.
4K(s)+O2(g)→2K2O(s)
(vigorous reaction)
● Mg, Al, Zn, and Pb react slowly with air and form a
protective layer that prevents corrosion.
2Mg(s)+O2(g)→2MgO(s)
(Mg burns with white dazzling light)
4Al(s)+3O2(g)→2Al2O3(s)
● Silver, platinum and gold don’t burn or react with air.
Basic Oxides of Metals
Metal oxides are crystalline solids that contain a metal cation and an oxide anion. They typically react with water to form bases or with acids to form salts. MO + H2O → M(OH)2 (where M = group 2 metal). Thus, these compounds are often called basic oxides.
Some metallic oxides get dissolved in water and form alkalis. Their aqueous solution turns red litmus blue.
Na2O(s)+H2O(l)→2NaOH(aq)
K2O(s)+H2O(l)→2KOH(aq)
Amphoteric Oxides of Metals
Amphoteric oxides are metal oxides which react with both acids as well as bases to form salt and water.
For example – Al2O3,
ZnO, PbO, SnO
Al2O3(s)+6HCl(aq)→2AlCl3(aq)+3H2O(l)
Al2O3(s)+2NaOH(aq)→2NaAlO2(aq)+H2O(l)
ZnO(s)+2HCl(aq)→ZnCl2(aq)+H2O(l)
ZnO(s)+2NaOH(aq)→Na2ZnO2(aq)+H2O(l)
Reactivity Series
The reactivity series of metals, also known as the activity series, refers to the arrangement of metals in the descending order of their reactivities.
The below table illustrates the reactivity of metals from high order to low order.
| Symbol | Element |
| K | Potassium ( Highly Active Metal) |
| Ba | Barium |
| Ca | Calcium |
| Na | Sodium |
| Mg | Magnesium |
| Al | Aluminium |
| Zn | Zinc |
| Fe | Iron |
| Ni | Nickel |
| Sn | Tin |
| Pb | Lead |
| H | Hydrogen |
| Cu | Copper |
| Hg | Mercury |
| Ag | Silver |
| Au | Gold |
| Pt | Platinum |
Roasting
Converts sulphide
ores into oxides on heating strongly in the presence of excess air.
It
also removes volatile impurities.
2ZnS(s)+3O2(g)+Heat→2ZnO(s)+2SO2(g)
Calcination
Converts carbonate and hydrated ores into oxides on heating strongly in the presence of limited air. It also removes volatile impurities.
ZnCO3(s)+heat→ZnO(s)+CO2(g)
CaCO3(s)+heat→CaO(s)+CO2(g)
Al2O3.2H2O(s)+heat→2Al2O3(s)+2H2O(l)
2Fe2O3.3H2O(s)+heat→2Fe2O3(s)+3H2O(l)
Reaction of Metals with Water or Steam
Aluminium, iron, and zinc are metals that do not react with either cold or hot water. However, when they come into contact with steam, they produce metal oxide and hydrogen. Lead, copper, silver, and gold are metals that do not react with water.
Metal+Water→Metalhydroxide or Metaloxide+Hydrogen
2Na+2H2O(cold)→2NaOH+H2+heat
Ca+2H2O(cold)→Ca(OH)2+H2
Mg+2H2O(hot)→Mg(OH)2+H2
2Al+3H2O(steam)→Al2O3+3H2
Zn+H2O(steam)→ZnO+H2
3Fe+4H2O(steam)→Fe3O4+4H2
Reaction of Metals with Acid
Metal+diluteacid→Salt+Hydrogengas
2Na(s)+2HCl(dilute)→2NaCl(aq)+H2(g)
2K(s)+H2SO4(dilute)→K2SO4(aq)+H2(g)
Only Mg and Mn,
react with very dilute nitric acid to liberate hydrogen gas.
Mg(s)+2HNO3(dilute)→Mg(NO3)2(aq)+H2(g)
Mn(s)+2HNO3(dilute)→Mn(NO3)2(aq)+H2(g)
Displacement Reaction
A more reactive element displaces a less reactive element from its compound or solution.
How Do Metals React with the Solution of Other Metal Salts
A more reactive metal can displace a less reactive metal from its salt solution in a displacement reaction. Metal displacement reaction is a common name for this reaction. The reactivity of certain regularly used metals has been ordered in decreasing order. This is referred to as the reactivity or activity series.
Metal A+Salt of metal B → Salt of metal A + Metal B
Fe(s)+CuSO4(aq)→FeSO4(aq)+Cu(s)
Cu(s)+2AgNO3(aq)→Cu(NO3)(aq)+2Ag(s)
- It’s a component of thermite welding. Aluminium displaces iron from its oxide in this process.
- It is used in the production of steel. In which iron is displaced from its oxide by carbon.
- It is mostly utilised in metal extraction.
Reaction of Metals with Bases
The base has a bitter taste and a slippery texture. A base dissolved in water is called an alkali. When chemically reacting with acids, such compounds produce salts. Bases are known to turn blue on red litmus paper.
Base+metal →
salt+hydrogen
2NaOH(aq)+Zn(s) → Na2ZnO2(aq)+H2(g)
2NaOH(aq)+2Al(s)+2H2O(l)
→ 2NaAlO2(aq)+2H2(g)
Extraction of Metals and Non-Metals
Applications of Displacement Reaction
Uses of displacement reaction
- Extraction of metals
- Manufacturing of steel
- Thermite reaction: Al(s)+Fe2O3(s) → Al2O3+Fe(molten)
The thermite reaction is used in the welding of railway tracks, cracked machine parts, etc.
Occurrence of Metals
Most elements,
especially metals, occur in nature in a combined state with other elements. All these compounds of metals
are known as minerals. But out of them, only a few are
viable sources of that metal. Such sources are called ores.
Au, Pt – exists in
the native or free state.
Extraction of Metals
The process of extracting metal ores buried deep underground is called Mining. The metal ores are found in the earth’s crust in varying abundance. The extraction of metals from ores is what allows us to use the minerals in the ground! The ores are very different from the finished metals that we see in buildings and bridges. Ores consist of the desired metal compound and the impurities and earthly substances called Gangue.
Metals of high
reactivity – Na, K, Mg, Al.
Metals of medium reactivity – Fe, Zn, Pb,
Sn.
Metals of low reactivity – Cu, Ag, Hg
Enrichment of Ores
It means the removal of impurities or gangue from ore through various physical and chemical processes. The technique used for a particular ore depends on the difference in the properties of the ore and the gangue.
In chemistry, a gangue is an undesirable substance or impurity that surrounds the mineral in an ore deposit, such as sand, rock, or any other material. When it comes to mining, this mineral is very frequent.
Extracting Metals Low in Reactivity Series
By self-reduction- when the sulphide ores of less electropositive metals like Hg, Pb, Cu etc., are heated in air, a part of the ore gets converted to oxide, which then reacts with the remaining sulphide ore to give the crude metal and sulphur dioxide. In this process, no external reducing agent is used.
1. 2HgS(Cinnabar)+3O2(g)+heat→2HgO(crude
metal)+2SO2(g)
2HgO(s)+heat→2Hg(l)+O2(g)
2. Cu2S(Copper
pyrite)+3O2(g)+heat→2Cu2O(s)+2SO2(g)
2Cu2O(s)+Cu2S(s)+heat→6Cu(crude
metal)+SO2(g)
3.
2PbS(Galena)+3O2(g)+heat→2PbO(s)+2SO2(g)
PbS(s)+2PbO(s)→2Pb(crudemetal)+SO2(g)
Extracting Metals in the Middle of Reactivity Series
Calcination is a process in which ore is heated in the absence of air, or air might be supplied in limited quantity. Roasting involves heating ore lower than its melting point in the presence of air or oxygen. Calcination involves the thermal decomposition of carbonate ores.
Smelting – it
involves heating the roasted or calcined ore (metal oxide) to a high temperature with a suitable reducing
agent. The crude metal is obtained in its molten state.
Fe2O3+3C(coke)→2Fe+3CO2
Aluminothermic
reaction – also known as the Goldschmidt reaction, is a highly exothermic reaction in which metal oxides,
usually of Fe and Cr, are heated to a high temperature with aluminium.
Fe2O3+2Al→Al2O3+2Fe+heat
Cr2O3+2Al→Al2O3+2Cr+heat
Extraction of Metals Towards the Top of the Reactivity Series
Electrolytic reduction:
1. Down’s process: Molten NaCl is electrolysed in a special apparatus.
At the cathode (reduction):
Na+(molten)+e−→Na(s)
Metal is deposited.
At the anode (oxidation):
2Cl−(molten)→Cl2(g)+2e–
Chlorine gas is liberated.
2. Hall’s process: A mixture of molten alumina and a fluoride solvent, usually cryolite (Na3AlF6), is electrolysed.
At the cathode (reduction):
2Al3++6e–→
2Al(s)
Metal is deposited.
At the anode (oxidation):
6O2– →
3O2(g)+12e–
Oxygen gas is liberated.
The metals at the top of the reactivity series are highly reactive. They cannot be obtained from their compounds by heating with carbon, because these metals have more affinity for oxygen than carbon. Hence, for the extraction of such metals electrolytic reduction method is used.
Refining of Metals
Refining of metals – removing impurities or gangue from crude metal. It is the last step in metallurgy and is based on the difference between the properties of metal and gangue.
Electrolytic Refining
Metals like copper,
zinc, nickel, silver, tin, gold etc., are refined electrolytically.
Anode: impure or
crude metal
Cathode: a thin strip of pure metal
Electrolyte: aqueous solution of metal salt
From anode (oxidation): metal ions are released into the
solution
At
cathode (reduction): the equivalent amount of metal from the solution is deposited
Impurities deposit at the bottom of the anode.
The Why Questions
Electronic Configuration
Group 1 elements – Alkali metals
| Element | Electronic Configuration |
| Lithium(Li) | 2,1 |
| Sodium(Na) | 2,8,1 |
| Potassium(K) | 2,8,8,1 |
| Rubidium(Rb) | 2,8,18,8,1 |
Group 2 elements – Alkaline earth metals
| Element | Electronic Configuration |
| Beryllium(Be) | 2,2 |
| Magnesium(Mg) | 2,8,2 |
| Calcium(Ca) | 2,8,8,2 |
| Stronium(Sr) | 2,8,18,8,2 |
How Do Metals and Non-Metals React?
Metals lose valence
electron(s) and form cations.
Non-metals gain those electrons in their
valence shell and form anions.
The cation and the anion are attracted
to each other by strong electrostatic force, thus forming an ionic bond.
For example: In calcium chloride, the ionic bond is formed by
opposite-charged calcium and chloride ions.
The calcium atom loses 2
electrons and attains the electronic configuration of the nearest noble gas (Ar). By doing so, it gains a
net charge of +2.
The two Chlorine atoms take one electron each, thus gaining a charge of -1 (each) and attain the electronic configuration of the nearest noble gas (Ar).
Ionic Compounds
Ionic compounds are neutral compounds that are made up of positively charged cations and negatively charged anions. Binary ionic compounds (ionic compounds containing only two types of elements) are named by first writing the name of the cation, then the name of the anion.
The electrostatic
attractions between the opposite-charged ions hold the compound together.
Example: MgCl2,
CaO, MgO, NaCl etc.
Properties of Ionic Compound
Ionic compounds
- Are usually crystalline solids (made of ions).
- Have high melting and boiling points.
- Conduct electricity when in aqueous solution and when melted.
- Are mostly soluble in water and polar solvents.
Physical Nature
Ionic solids usually exist in regular, well-defined crystal structures.
Electric Conduction of Ionic Compounds
Ionic compounds
conduct electricity in the molten or aqueous state when ions become free and act as charge carriers.
In solid form, ions are strongly held by electrostatic forces of
attraction and are not free to move; hence do not conduct electricity.
For example, ionic compounds such as NaCl do not conduct electricity
when solid, but when dissolved in water or in a molten state, they will conduct electricity.
Melting and Boiling Points of Ionic Compounds
In ionic compounds, the strong electrostatic forces between ions require a high amount of energy to break. Thus, the melting point and boiling point of an ionic compound are usually very high.
Solubility of Ionic Compounds
Most ionic
compounds are soluble in water due to the separation of ions by water. This occurs due to the polar nature
of water.
For example, NaCl is a 3-D salt crystal composed of Na+ and
Cl− ions
bound together through electrostatic forces of attraction. When a crystal of NaCl comes into contact with
water, the partially positive-charged ends of water molecules interact with the Cl− ions,
while the negatively charged end of the water molecules interacts with the Na+ ions.
This ion-dipole interaction between ions and water molecules assists in the breaking of the strong
electrostatic forces of attraction within the crystal and, ultimately, in the solubility of the crystal.
Corrosion
Alloys
Alloys are homogeneous mixtures of a metal with other metals or nonmetals. Alloy formation enhances the desirable properties of the material, such as hardness, tensile strength and resistance to corrosion.
Examples of a few
alloys:
Brass: copper and zinc
Bronze: copper and tin
Solder:
lead and tin
Amalgam: mercury and other metal
Corrosion
Gradual
deterioration of a material, usually a metal, by the action of moisture, air or chemicals in the
surrounding environment.
Rusting:
4Fe(s)+3O2(from air)+xH2O(moisture)→2Fe2O3. xH2O(rust)
Corrosion of
copper:
Cu(s)+H2O(moisture)+CO2(from
air)→CuCO3.Cu(OH)2(green)
Corrosion of
silver:
Ag(s)+H2S(from air)→Ag2S(black)+H2(g)
Prevention of Corrosion
Prevention:
1. Coating
with paints or oil or grease: The application of paint or oil or grease on metal surfaces keep out air and
moisture.
2. Alloying: Alloyed metal is more resistant to corrosion. Example: stainless steel.
3. Galvanization: This is a process of coating molten zinc on iron articles. Zinc forms a protective layer and prevents corrosion.
4. Electroplating: It is a method of coating one metal with another by the use of an electric current. This
method not only lends protection but also enhances the metallic appearance.
Examples: silver plating, and nickel plating.
5. Sacrificial protection: Magnesium is more reactive than iron. When it is coated on articles made of iron or steel, it acts as the cathode undergoes a reaction (sacrifice) instead of iron and protects the articles.