1. Physical Properties of Metals and Non-Metals

Metals: Physical properties of metals, chemical properties of metals and non-metal oxide.

Metals are the elements that conduct heat and electricity and are malleable and ductile. Examples are Iron (Fe), Aluminium (Al), Silver (Ag), Copper (Cu), Gold (Au), Platinum (Pt), Lead (Pb), Potassium (K), Sodium (Na), Calcium (Ca) and Magnesium (Mg) etc.

Metals are the elements which form positive ions by losing electrons. Thus, metals are known as Electropositive Elements.

How Do Metals and Nonmetals 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.
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).

Physical Properties of Metals

  • Hardness: Most of the metals are hard, except alkali metals, such as sodium, potassium, lithium, etc. are very soft metals. These can be cut by using a knife.
  • Conduction: Metals are a good conductor of heat and electricity. This is the cause that electric wires are made of metals like copper and aluminium.
  • Melting and Boiling Point: :Metals have generally high melting and boiling points. (Except sodium and potassium metals which have low melting and boiling point.)
  • Strength: Most of the metals are strong and have high tensile strength. Because of this, big structures are made using metals, such as copper (Cu) and iron (Fe). (Except Sodium (Na) and potassium (K) which are soft metals).
  • State: Metals are solid at room temperature except for mercury (Hg).
  • Malleability: Metals are malleable. This means metals can be beaten into a thin sheet. Because of this property, iron is used in making big ships.
  • Ductility: Metals are ductile. This means metals can be drawn into thin wire. Because of this property, a wire is made of metals.
  • Density: Most of the metals have a high density.
  • Colour: Most of the metals are grey in colour. But gold and copper are exceptions.

 Non-Metals: Physical Properties of non-metals, chemical properties of non-metals, non¬metal oxides, Reaction of metal and Non-metal, Ionic bonds and formation of an ionic bond. Non-metals are the elements that do not conduct electricity and are neither malleable nor ductile.
Examples: Carbon (C), Sulphur (S), Phosphorous (P), Silicon (Si), Hydrogen (H), Oxygen (O), Nitrogen (N), Chlorine (Cl), Bromine (Br), Neon (Ne) and Argon (Ar) etc.
Non-metals are the elements which form negative ions by gaining an electron. Thus, non¬metals are also known as Electronegative Elements.

Physical properties of non-metals

  • Hardness: Non-metals are not hard rather they are generally soft. But the diamond is an exception; it is the hardest naturally occurring substance.
  • State: Non-metals may be solid, liquid or gas.
  • Lustre: Non-metals have a dull appearance. Diamond and iodine are exceptions.
  • Sonority: Non-metals are not sonorous, i.e., they do not produce a typical sound on being hit.
  • Conduction: Non-metals are a bad conductor of heat and electricity. Graphite which is allotrope of carbon is a good conductor of electricity and is an exception.
  • Malleability and ductility: Non-metals are brittle.
  • Melting and boiling point: Non-metals have generally low melting and boiling points.
  • Density: Most of the non-metals have low density.
  • Colour: Non-metals are in many colours.

Carbon in the form of graphite is non-metal which conduct electricity.

Iodine is non-metal which is lustrous having a shining surface.

Carbon in the form of diamond is a non-metal which is extremely hard.

Diamond is a non-metal which has a very high melting point and boiling point.

1. Physical Properties of Metals and Non-Metals

CHAPTER  3

METALS AND NON- METALS

 Physical properties of metals and non- metals.

 1     LUSTRE  -

  Metals have shining surface.

    Non-metals do not have shinig surface

 * Except iodine   

 2 Hardness -

    Metals are generally hard.

  • Except sodium lithium and potassium which are soft and can be cut with knife.

 Non – metals are generally soft.

  • Except diamond a form of carbon which is the hardest natural substan

3 Malleability - 

 Metals can be beaten into thin sheets gold and silver are the most malleable metals.

Non- metals are non-mallerable.

4 Ductility -

 Metals can be drawn into thin wires.

 Non- metals can not be drawn into thin wires. They are non- ductile.

5 Conductor of heat and electricity -

Metals are good conductor of heat and electricity of heat. Lead and mercury are poor conductor of heat.

Metals are good conductor of heat and electricity.

Non- metals are poor conductor of heat EXCEPT graphite.

6 state -

 The metals exist as solids. EXCEPT mercury

The non- metals exist as solids or gaseous.EXEPT  bromine.

7 density -

Metals have high density and high melting point.

Except sodium and potassium.

8 Oxides -

 Metallic oxides are basic in nature.

Non metallic oxides are acidic in nature.

9 Sonorous -

Metals produce a sound on strinking a hard surface.

Non-metals are not sonorous.

2. Chemical Properties of Metals

Chemical Properties of Metals

1. Reaction with oxygen: Most of the metals form respective metal oxides when reacting with oxygen.

Metal + Oxygen → Metal Oxide

Examples:

Reaction of Potassium with Oxygen: Potassium metal forms potassium oxide when reacts with oxygen.

Reaction of Sodium with Oxygen: Sodium metal forms sodium oxide when reacts with oxygen.


 Reaction of Copper metal with Oxygen: Copper does not react with oxygen at room temperature but when burnt in air, it gives oxide.

Silver, gold and platinum do not combine with the oxygen of air even at high temperature. They are the least reactive.

2. Reaction of metals with water: Metals form respective hydroxide and hydrogen gas when reacting with water.

Metal + Water → Metal hydroxide + Hydrogen

Reaction of Sodium metal with Water: Sodium metal forms sodium hydroxide and liberates hydrogen gas along with lot of heat when reacting with water.

Reaction of Iron with Water: Reaction of iron with cold water is very slow and comes into notice after a long time. Iron forms rust (iron oxide) when reacts with moisture present in the atmosphere. Iron oxide and hydrogen gas are formed by passing of steam over iron metal.

3. Reaction of metals with dilute acid: Metals form respective salts when reacting with dilute acid. Reaction of Sodium metal with dilute hydrochloric acid: Sodium metal gives sodium chloride and hydrogen gas when react with dilute hydrochloric acid.

Hydrogen gas is not when metal is treated with nitric acid (HNO3):

Nitric acid is strong oxidising agent and it oxidises the hydrogen gas (H2) liberated into water (H2O) and itself get reduced to some oxide of nitrogen like nitrous oxide (N2O)3 nitric oxide (NO) and nitrogen dioxide (NO2).

Copper, gold, silver are known as noble metals. These do not react with water or dilute acids.

The order of reactivity of metal towards dilute hydrochloric acid or sulphuric acid is in the order;

K > Na > Ca > Mg > Al > Zn > Fe > Cu > Hg > Ag

Metal Oxides

Chemical Properties: Metal oxides are basic in nature. The aqueous solution of metal oxides turns red litmus blue.

Reaction of Metal oxides with Water: Most of the metal oxides are insoluble in water. Alkali metal oxides are soluble in water. Alkali metal oxides give strong base when dissolved in water.


Reaction of Sodium oxide with Water: Sodium oxide gives sodium hydroxide when reacts with water.

Reaction of Zinc oxide and Aluminium oxide: Aluminum oxide and zinc oxide are insoluble in water. Aluminium oxide and zinc oxide are amphoteric in nature. An amphoteric substance shows both acidic and basic characters. It reacts with base like acid and reacts with an acid like a base.

When zinc oxide reacts with sodium hydroxide, it behaves like an acid. In this reaction, sodium zincate and water are formed

Reactivity Series of Metals: The order of intensity or reactivity of metal is known as Reactivity Series. Reactivity of elements decreases on moving from top to bottom in the given reactivity series. In the reactivity series, copper, gold, and silver are at the bottom and hence, least reactive. These metals are known as Noble metals.

Reactivity of some metals are given in descending order :

K > Na > Ca > Mg > Al > Zn > Fe > Pb > Cu

 

2. Chemical Properties of Metals

Chemical properties of Metals

(i)  Reaction with air :-

All metals combine with oxygen to form metal oxide .

Metal + O2           Metal oxide

For example,

2Cu  +  O2          2CUO    copperoxide (black)

 4Al + 3O2        2Al2O3     Aliminium oxide

  • Sodium and potassium react so vigorously that they catch fire in open so they are kept immersed in kerosence
  • Surfaces of Mg , Al, Zn pb are covered with a thin layer of oxide whish prevent them from further oxidation. Anodizing is a process of forming a tick oxide layer of aluminium.
  • Iron does not burn on heating but iron filling burn vigorously.
  •  Copper does not burn but the hot metal is coated with a black coloured layer of copper (ii) oxide
  • Silver and gold do not react with oxygen even at high temperatures.

Amphoteric oxide    Metal oxides which react with both acids as well as bases to produce salts and water are called amphoteric oxides.

Example ;   

Al2O3   +    6HCl      2AlCl3 +   H2O

 Al2O3  +  2NaOH    2NaAlO2    +   H2O

                                (SODIUM ALUMINATE )

Most metal oxides are insoluble in water but some of these dissolve in water to form alkalis. Sodium oxide and potassium oxide dissolved in water to produce alkali.

Na2O(S) +    H2O(l)        2 NaOH (aq)

K2O(S)    +   H2O(l)         2 KOH (aq)

(ii)  Reactions of metals with water:-

 Metal+ water        Metal oxide  +  hydrogen

Metal oxide  +  water      metal hydroxide

Metals like potassium and sodium react violenty with cold water.

Na  +   2 H2O          NaOH + H2  + heat energy.

The reaction of calcium of water is less violent

Ca + 2H2O    →   Ca(oH)2 + H2

Magnesium react with hot water to form magnesium hydroxide and hydrogen.

Mg +  2H2O     →  Mg(OH)   H2

Metals like aluminium iron and zinc do not react with cold or hot water. But they react with steam to form metal oxide and hydrogen.

2 Al + 3 H20         Al2O3  +   3H2

3 Fe  + 4 H2O     →  Fe3O4   +   4H2

Metal such as lead ,copper, silver and gold do not react with water at all.

(ii) Reaction of metals with acids.

Metal +  Dilute acid     salt + hydrogen

Copper and silver do not react wit dil acids.

 For example 

 Fe +  2HCl          FeCl2  +  H2

Mg  +  2HCl          MgCl2  +   H2

Zn  +   2HCl           ZnCl2  +  H2

(iii)   Reaction of metals with solution of other Metal salts ; 

Metal A+  Salts solution B      Salt solution A + Metal B   

Reaction of metals with salt solutions.

More reactive metals can displace less reactive metals from their compounds in solution form.

Fe +  CuSO4          FeSO4  +    Cu       

Fe displaces Cu because Fe is more reactive metals than Cu .

REACTIVITY SERIELS ;

The reactivity series is a list of metals arranged in the order of their decreasing activities.

3. Metals and Non-Metals Reactions

Ionic Compounds

The electrostatic attractions between the opposite charged ions hold the compound together.
Example: MgCl2, CaO, MgO, NaCl etc.

Properties of Ionic Compound

Ionic compounds

  1. Are usually crystalline solids (made of ions).
  2. Have high melting and boiling points.
  3. Conduct electricity when in aqueous solution and when melted.

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 attractions and are not free to move; hence do not conduct electricity.

For example, ionic compounds such as NaCl does not conduct electricity when solid but when dissolved in water or in a molten state, it will conduct electricity.

Salt solution 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 attractions. When a crystal of NaCl comes into contact with water, the partial positively 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 assist in the breaking of the strong electrostatic forces of attractions within the crystal and ultimately in the solubility of the crystal.

.Properties of Ionic compound

  • Ionic compounds are solid. Ionic bond has a greater force of attraction because of which ions attract each other strongly. This makes ionic compounds solid.
  • Ionic compounds are brittle.
  • Ionic compounds have high melting and boiling points because force of attraction between ions of ionic compounds is very strong.
  • Ionic compounds generally dissolve in water.
  • Ionic compounds are generally insoluble in organic solvents; like kerosene, petrol, etc.
  • Ionic compounds do not conduct electricity in the solid state.

3. Metals and Non-Metals Reactions

METALS AND NON- METALS REACTIONS

  • Atoms of the metals lose electron from their valence shell to form cation.
  • Atoms of the non- metals gain electrons in the valence to form anion.

Formation of sodium chloride.

Formation of magnesium chloride

Properties of ionic compound

  1. PHYSICAL NATURE  :-  They are solid and hard (because of the strong force of attraction between the positive and negative ions) . They are brittle.
  2. Melting and boiling point:_  They have high melting and boiling pont.
  3. Solubility :- soluble in water and insoluble in solvents such as kerosene , petrol etc.
  4. Conductor of electricity :- Ionic compound conduct electricity in molter (ions move to the opposite electrodes  when electricity is passed )
  • They do not conduct electricity in solid state as movements of ions is not possible in solid They conduct electricity in molten state.

4. Occurence of Metals

Occurrence of Metals

Most of the elements, especially metals occur in nature in the 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 – exist in the native or free state.

Extraction of Metals

Metals of high reactivity – Na, K, Mg, Al.

Metals of medium reactivity – Fe, Zn, Pb, Sn.

Metals of low reactivity – Cu, Ag, Hg

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)

2Fe2O3.3H2O(s)+heat→2Fe2O3(s)+3H2O(l)

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(Copperpyrite)+3O2(g)+heat→2Cu2O(s)+2SO2(g)

2Cu2O(s)+Cu2S(s)+heat→6Cu(crude metal)+SO2(g)

Extracting Metals in the Middle of Reactivity Series

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: 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.

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.

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 the 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 solution is deposited

Impurities deposit at the bottom of the anode.

4. Occurence of Metals

OCCURANCE OF METALS

MINERALS :- The elements which occur naturally in the earth’s crust are called minerals.

ORES:- Minerals that contain very high percentage of particular metal andthe metal can be profitably extracted from it, such minerals are called ores.

  • Metals at the bottom of the activity series are least reactive they are often found in free state . For Example –  Ag,Au, Cu.
  • Metals at the ttop of the acitivity series (k,Na, Co, Mg, and Al) are so reactive that they never found in free state.
  • Metals in the middle of the activity series ( Zn, Fe,Pb etc are moderately reactive . they occur as sulphates ,oxides or carbonates.
  • They ore of many metals are oxide because oxygen is very reactive and is abundant on the earth.
  • Steps involved in the extraction of pure metals from ores.

 Step involved in the extraction of metals from ores.

 EXTRACTION OF METALS FROM ORES :-

 1 ENCRICHMENT OF ORES    Ores are usually contaminated with large amounts of impurities such as soil, sand etc called gangue these impuriestes are removed from the ore prior to the extraction of mertal.

 2  Extraction of metals   Metals low in the activity series are very anreactive the oxides of these metals are reduced to metals by heating.

 For example           

 2HgO         2Hg  +  O2

Mercury oxide is reduced to mercury on heating

The metals in the middle of the activity series ( Zn, Fe, Pb, Cu ) are moderately active. The metal sulphides  and carbontes  are converted into metal oxide. The sulphate ores are converted into oxides by heating strongly in the presence of excess air this process is known as roasting .

2 ZnS + 3O2 →      2ZnO + CO2

  • The carbonate ores are changed into oxides by heating strongly in limited air this process is called calcination .

 Zn CO3          ZnO + CO2

  • Then metal oxides are reduced to corresponding metals by using reducing agent like carbon.

 ZnO  +  C       Zn   +  CO

  • This reaction of iron (iii) oxide ( Fe2O3) with aluminium is used to join railway tracks or cracked machine parts . this reaction is known as thermit reaction.

 Fe2O3    +  2Al      2 Fe +  Al2O3 + heat

 Metals high up in the reactivity series are very reactive. The metalo are obtained by electrolytic reduction. The metals are deposit at the cathode and chlore is deposited at anode.

At cathode     Na+    +  e-         Na
At anode        2Cl-       Cl2   +  2e-

Refining of metals

 The most widely used method for refining impure metals is electronic refining.

Electrolytic refining of copper

Electrolytic refining :-   Metals  ( Cu, Zn , Ag, Au etc ) are refined electrolytically . the impure metal is made the anode and athin strip of pure metal is made the cathode . a solution a metal salt is used as an electrolyte.

Electrolytic refining of copper.

Anode :  Impure copper

Cathode: Strip of pure copper

The insoluble impurities settle at the bottom of the anode and is called anode mud.  

5. Corrosion

Corrosion

Alloys

Alloys are homogeneous mixtures of 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 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: 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 electric current. This method not only lends protection but also enhances the metallic appearance.

5. Corrosion

CORROSION :-

The surface of metals is corroded when they are exposed to moist air for a long period of time. This is called corrosion

For example 

  • Silver becomes black when exposed to air and form a coating of silver sulphide.
  • Copper react with  moist Co2 and form a green coat of copper carbonate.
  • Iron acquires a coating of brown floky substance called rust.

 Prevention of corrosion. 

The rusting of iron can be prevented by painting , oiling, greasing, galvanizing, chrome plating anodizing and making alloys.

  • Galvanisation:-  It is a method of protecting steel and iron from ressting by coating them a thin layer of zinc.
  • Alloying is a very good method of improving the properties of a metal an alloy is a homogenous mixture of two or more metals or a metal and a non- metal
  • Iron      it is mixed with small amount of carbon
  • Steel    iron + nickel and chromium.
  • Brass   copper  + tin
  • Solder   →  lead  +  tin ( used for welding electric wire together )