Metals And Non-Metals 

Metals

Elements which are generally solid at room temperature and which forms positive ions by releasing electron from its outermost shell are called metals. Some examples of metal are- iron, copper, gold, mercury etc. 

Physical Properties of metals

1. Lustre: Metals, in their pure state, have a shining surface. This property is called metallic lustre.

2. Hardness: Metals are generally hard. The hardness varies from metal to metal.

3. Malleability:Metals can be beaten into thin sheets. This property is called malleability. Gold and silver are the most malleable metals.

4. Ductility: The ability of metals to be drawn into thin wires is called ductility. Gold is the most ductile metal.

5. Thermal and electrical conductivity:  Metals are good conductors of heat & electricity. The best conductors of heat are silver and copper. Lead and mercury are comparatively poor conductors of heat.

6. Melting Point: Metals have high melting points.

7. Colour: Metals are generally black, silver, golden- yellow in colour.

8. Sonorous: Metals produce sonorous sound on striking a hard surface.

Non-Metals

The elements that accept electrons and form negative ions are called non-metals.

Physical properties of non-metals

1. Non-metals show different colors.

2. Non-metals can not conduct heat and electricity.

3. Non-metals have no lustre, not sonorous and do not  possess ductility and malleability.

4.  Non-metals have low melting point.

Metals and non-metals with exceptional properties

1. Mercury is the only metal found in liquid state at normal temperature.

2. Metals have high melting point, but the melting points of gallium and cesium are lower.

3. Iodine is a non-metal but it is lustrous. 

4. Carbon is a non-metal that can exist in different forms. Each form is called an allotrope. Diamond, an allotrope of carbon, is the hardest natural substance known and has a very high melting and boiling point. Graphite, another allotrope of carbon, is a conductor of electricity. (iv) Alkali metals (lithium, sodium, potassium) are so soft that they can be cut with a knife. They have low densities and low melting points.

5. Alkali metals like lithium, sodium, potassium etc. are soft and can be cut with a knife. They have low density and low melting point.

Chemical properties of Metals

1. Burning of metals in air:

    Almost all metals combine with oxygen to form metal oxides. 

                    Metal + Oxygen → Metal oxide 

 For example, when copper is heated in air, it combines with oxygen to form copper(II) oxide, a black oxide. 

                    2Cu + O2 → 2CuO 

                                (Copper)              (Copper(II) oxide) 

Similarly, aluminium forms aluminium oxide. 

                    4Al + 3O2 → 2Al2O3 

                           (Aluminium)             (Aluminium oxide)

2. Reaction of metals with water:

        Metals react with water and produce corresponding metal oxide and hydrogen gas. Metal oxides that are soluble in water dissolve in it to further form metal hydroxide. But all metals do not react with water. 

                        Metal + Water → Metal oxide + Hydrogen 

                        Metal oxide + Water → Metal hydroxide

       Metals like potassium and sodium react violently with cold water. In case of sodium and potassium, the reaction is so violent and exothermic that the evolved hydrogen immediately catches fire. 

                        2K(s) + 2H2O(l) → 2KOH(aq) + H2(g) + heat energy 

                        2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) + heat energy

3. Reaction of metals with acid:

        Metals react with acids to give a salt and hydrogen gas. 

                       

                     Metal + Dilute acid → Salt + Hydrogen

        More active metals can remove less active metals from its salts solution. Eg.

             Fe (s) + CuSO4 (aq) --- → FeSO4 (aq) + Cu (s)

        Iron  removed Cu because it is more active than copper.    

4. Reaction of metals with non-metals:

        When metals and non-metals react with each other, exchange of electrons takes place between them. The electrons released by the metal are taken by the non-metals and a new compound is formed between them through ionic bonding.

            For example, we can see the reaction between sodium and chlorine. Sodium atom has one electron in its outermost shell. If it loses the electron from its M shell then its L shell now becomes the outermost shell and that has 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 has seven electrons in its outermost shell and it requires one more electron to complete its octet. If sodium and chlorine were to react, the electron lost by sodium could be taken up by chlorine. After gaining an electron, the chlorine atom gets a unit negative charge, because its nucleus has 17 protons and there are 18 electrons in its K, L and M shells. This gives us a chloride anion C1–. 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).

    Compounds formed between metals and non-metals are called ionic compounds or electrovalent compounds, as they are formed through electrostatic force of attraction between oppositely charged ions.

Properties of Ionic Compounds 

(1) Physical properties- Since ionic compounds are formed by strong electrostatic force of attraction between positive and negative ions, ionic compounds are solid or whole and somewhat strong.

(2) These compounds are brittle and can be easily broken into pieces by application of  pressure.

(3) Melting point and boiling point: Due to the strong force of inter-ionic electrostatic attraction, it needs a lot of heat energy to separate the oppositely charged ions, hence the melting point and boiling point of ionic compounds are very high.

(4) Solubility: Electrolytic compounds are soluble in water and insoluble in solvents such as kerosene and petrol.

5. Electrical Conductivity- The conduction of electricity through a solution involves the movement 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. Ionic compounds in the solid state do not conduct electricity because movement of ions in the solid is not possible due to their rigid structure. But ionic compounds conduct electricity in the molten state. This is possible in the molten state since the electrostatic forces of attraction between the oppositely charged ions are overcome due to the heat. Thus, the ions move freely and conduct electricity.

Chemical Properties of Non-Metals

1. Reaction of Non-metals with Oxygen: Non-metals form respective oxide when reacting with oxygen.
                    Non-metal + Oxygen → Non-metallic oxide
When carbon reacts with oxygen, carbon dioxide is formed along with the production of heat.

                    C (s) + O2 (g) → CO2 (g) + Heat ( in sufficient oxygen)

        When carbon is burnt in an insufficient supply of air, it forms carbon monoxide. 

                2C (g) + O2 (g) → 2CO(g)

2. Reaction of Non-metal with Chlorine: Non-metal gives respective chloride when they react with chlorine gas.

                Non-metal + Chlorine → Non-metal chloride

Hydrogen gives hydrogen chloride when reacts with chlorine.

                H2 + 2Cl → 2HCl

3. Reaction of Non-metals with Hydrogen: Non-metals react with hydrogen to form covalent hydrides.

                Non-metal + Hydrogen → Covalent Hydride

Sulphur combines with hydrogen to form a covalent hydride called Hydrogen sulphide.

                H2 + S → H2S

Occurrence of metals

            Metals are found in nature as  free elements or compounds. The earth's crust is the main source of metal. Seawater also contains some soluble salts, such as sodium chloride and magnesium chloride etc. Materials found in the earth's crust in the form of elements or compounds are called minerals. Some minerals contain high percentage of some metals and can be used to extract those metals profitably. This type of mineral is called ore.

Metal Extraction

        Metals are extracted from their ores applying various methods, depending on their place in the reactivity series. Some metals are found in the earth’s crust in the free state. Some are found in the form of their compounds. The metals at the bottom of the activity series are the least reactive. They are often found in a free state. For example, gold, silver, platinum and copper are found in the free state. Copper and silver are also found in the combined state as their sulphide or oxide ores. The metals at the top of the activity series (K, Na, Ca, Mg and Al) are 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. They are found in the earth’s crust mainly as oxides, sulphides or carbonates.  Depending on the activity of the metals, different methods are used to extract them.

        Several steps are involved in the extraction of pure metal from ores. These are-

Ore → Enrichment of ores → Conversion of ore into metal oxide → Reduction of metal→ Refining of metals

Enrichment of Ores:  Ores mined from the earth are usually contaminated with large amounts of impurities such as soil, sand, etc., called gangue. The impurities must be removed from the ore prior to the extraction of the metal. The processes Several steps are involved in the extraction of pure metal from ores. Methods used for removing the gangue from the ore are based on the differences between the physical or chemical properties of the gangue and the ore. Different separation techniques are accordingly employed. Some of them are-

 1. Chemical method

2. Magnetic Separation

3. Froth flotation method etc.

Conversion of enriched ore to metallic oxide: Ores of metals are generally found as carbonates and sulphates. These carbonates and sulfates are converted to respective oxides, as the oxides of the metal can be easily reduced to obtain the metal. There are generally two methods used to convert ores to oxide-

1. Roasting

2. Calcination

Reduction of metallic oxides to metals: Based on the reactivity of the metal, appropriate method and suitable reducing agents are used to reduce the metallic oxides into corresponding metal. There are generally two methods used in this field- 

1. Carbon reduction

2. Electrical reduction.

Refining of metals: The metals which are extracted from an alloy contains residues of other metals and some other impurities. The removal of these impurities is called refining of metal. There are generally two methods used for metal refining -

1. Distillation method

2. Electrolytic refining.

Extracting Metals Low in the Activity Series 

            Metals low in the activity series are very unreactive. The oxides of these metals can be reduced to metals by heating alone. For example, 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. 

                        2HgS(s) + 3O2 (g) ^Heat→  2HgO(s) + 2SO2 (g)

                        2HgO (s)  ^Heat→ 2Hg (l) + O₂ (g)

            Similarly, copper which is found as Cu2S in nature can be obtained from its ore by just heating in air. 

                    2Cu₂S (s) + 30₂ (g) ^Heat→ 2Cu₂O (s) + 2SO₂ (g)

                     2Cu₂O (s) + Cu₂S (s) ^Heat→ 6 Cu (s) + SO₂ (g)

Extracting Metals in the Middle of the Activity Series 

            The metals in the middle of the 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. Therefore, prior to reduction, the metal sulphides and carbonates must be converted into metal oxides. There are two methods used for this purpose- 

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

                        ZnS (s) + 3O₂ (g) ^Heat→ 2ZnO (s) + 2SO₂ (g) 

2. Calcination- The carbonate ores are changed into oxides by heating strongly in limited air. This process is known as calcination. 

                           

                         ZnCO₃ (s) ^Heat→ ZnO (s) + CO₂ (g)

        The metal oxides are then reduced to the corresponding metals by using suitable reducing agents such as carbon. For example, when zinc oxide is heated with carbon, it is reduced to metallic zinc. 

                        ZnO(s) + C(s) → Zn(s) + CO(g)

Extracting Metals towards the Top of the Activity Series 

        The metals high up in the reactivity series are very reactive. They cannot be obtained from their compounds by heating with carbon. For example, carbon cannot reduce the oxides of sodium, magnesium, calcium, aluminium, etc., to the respective metals. This is because these metals have more affinity for oxygen than carbon. These metals are obtained by electrolytic reduction. For example, sodium, magnesium and calcium are obtained by the electrolysis of their molten chlorides. The metals are deposited at the cathode (the negatively charged electrode), whereas, chlorine is liberated at the anode (the positively charged electrode). The reactions are – 

At cathode----        Na⁺ + e^-  →Na 

At anode ----        2Cl^- → Cl₂ + 2e^-

        Similarly, aluminium is obtained by the electrolytic reduction of aluminium oxide.

Refining of Metals

        The 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 is electrolytic refining. 

Electrolytic Refining: Many metals, such as copper, zinc, tin, nickel, silver, gold, etc., are refined electrolytically. In this process, the impure metal is made the anode and a thin strip of pure metal is made the cathode. A solution of the metal salt is used as an electrolyte. The apparatus is set up as shown in figure below. On passing the current through the electrolyte, the pure metal from the anode dissolves into the electrolyte. An equivalent amount of pure metal from the electrolyte is deposited on the cathode. The soluble impurities go into the solution, whereas, the insoluble impurities settle down at the bottom of the anode and are known as anode mud.

Corrosion

        When a metal is exposed to various substances around it such as water vapor, acids, etc., it is seen with different colors on them. This process is called metal corrosion. Examples of corrosion are the reddish brown layer (rusting) on the iron, the green layer on the copper and the reddish brown coating on the silver.

 Prevention of corrosion: The following methods can be applied to prevent corrosion.

1. Painting: The surface of metal is painted to prevent corrosion so that the metal is not exposed to water vapor.

2. Oiling: It is also possible to protect the metal from corrosion by applying oil on it.

3. Greasing: Grease can be applied at the junction of metal devices to prevent corrosion.

4. Galvanising: Galvanizing is a method of applying zinc coating to steel and stainless steel. The metal does not affect from corrosion even after the zinc coating has been removed.

5. Chrome plating: In this method a thin layer of chromium metal is electro-chemically coated on the metal surface.