How Do Metals And Non-metals React?

The chemical reactivity of elements is related to their tendency to achieve a stable electronic configuration, usually by attaining a completely filled outermost electron shell (like noble gases).

Metals tend to lose electrons from their outermost shell to form positively charged ions (cations). Non-metals tend to gain electrons in their outermost shell to achieve a complete octet and form negatively charged ions (anions).

When a metal reacts with a non-metal, electrons are transferred from the metal atom(s) to the non-metal atom(s).

Formation of Ionic Compounds (Electrovalent Compounds):

• Example: Formation of Sodium Chloride ($\text{NaCl}$): A sodium atom (Na) has 1 electron in its outermost shell (electronic configuration 2, 8, 1). It readily loses this electron to form a stable sodium ion ($\text{Na}^+$, electronic configuration 2, 8). A chlorine atom (Cl) has 7 electrons in its outermost shell (electronic configuration 2, 8, 7). It readily gains one electron to complete its octet and form a stable chloride ion ($\text{Cl}^-$, electronic configuration 2, 8, 8). The electron lost by sodium is gained by chlorine.

  $\text{Na} \longrightarrow \text{Na}^+ + \text{e}^-$

  $\text{Cl} + \text{e}^- \longrightarrow \text{Cl}^-$

  $\text{Na}^+ + \text{Cl}^- \longrightarrow \text{NaCl}$ (ionic compound)

• Example: Formation of Magnesium Chloride ($\text{MgCl}_2$): A magnesium atom (Mg) has 2 electrons in its outermost shell (electronic configuration 2, 8, 2). It needs to lose these two electrons to become a stable $\text{Mg}^{2+}$ ion (electronic configuration 2, 8). Each chlorine atom (Cl) needs one electron to complete its octet. Therefore, one Mg atom transfers one electron to each of the two Cl atoms, forming one $\text{Mg}^{2+}$ ion and two $\text{Cl}^-$ ions.

  $\text{Mg} \longrightarrow \text{Mg}^{2+} + 2\text{e}^-$

  $2\text{Cl} + 2\text{e}^- \longrightarrow 2\text{Cl}^-$

  $\text{Mg}^{2+} + 2\text{Cl}^- \longrightarrow \text{MgCl}_2$ (ionic compound)

Compounds formed by the transfer of electrons from a metal to a non-metal are called ionic compounds or electrovalent compounds. They are held together by strong electrostatic forces of attraction between the oppositely charged ions.

Properties Of Ionic Compounds

Ionic compounds exhibit distinct physical properties due to the strong electrostatic forces between their ions:

• Physical Nature: Ionic compounds are generally hard, brittle solids. The strong attractive forces between ions make them solid and hard. When pressure is applied, the ordered arrangement of ions can be disrupted, causing the crystal to break (brittleness).
• Melting and Boiling Points: Ionic compounds have high melting and boiling points. A large amount of thermal energy is required to overcome the strong electrostatic forces holding the ions together in the crystal lattice.
• Solubility: Ionic compounds are generally soluble in water (a polar solvent) because water molecules can effectively separate and hydrate the ions. They are generally insoluble in non-polar solvents like kerosene and petrol.
• Conduction of Electricity: The conduction of electricity requires the movement of charged particles (ions or electrons).
  • In the solid state, ionic compounds do not conduct electricity because the ions are held in fixed positions in the crystal lattice and cannot move freely.
  • In the molten state or in aqueous solution, ionic compounds conduct electricity because the ions are free to move. In the molten state, heat overcomes the electrostatic forces, allowing ions mobility. In solution, water molecules separate the ions, allowing them to move and carry electric charge towards the electrodes.

Occurrence Of Metals

The majority of metals are found in the Earth's crust. Some are found in seawater as soluble salts. Metals occur either in the free state (native state) or in the form of their compounds.

Elements or compounds found naturally in the Earth's crust are called minerals. Ores are minerals from which metals can be profitably extracted.

The occurrence of metals in nature is related to their reactivity:

• Least Reactive Metals: Metals low in the activity series (e.g., gold, silver, platinum) are very unreactive and are often found in the free state (native state). Copper and silver can also be found in combined forms as sulphides or oxides.
• Moderately Reactive Metals: Metals in the middle of the activity series (e.g., zinc, iron, lead, copper) are moderately reactive. They are usually found in the Earth's crust in the form of their oxides, sulphides, or carbonates. Oxide ores are common because oxygen is very reactive and abundant.
• Most Reactive Metals: Metals at the top of the activity series (e.g., potassium, sodium, calcium, magnesium, aluminium) are highly reactive and are never found in nature as free elements. They exist only in the form of their compounds.

The process of extracting pure metals from their ores and then refining them for use is called metallurgy.

Extraction Of Metals

The process of extracting metals from their ores involves several steps, which vary depending on the metal's reactivity. A general overview of the steps is shown in the diagram.

Enrichment Of Ores

Ores mined from the Earth often contain unwanted impurities like soil, sand, and rocky material. These impurities are collectively called gangue. The gangue must be removed from the ore before metal extraction.

The processes used to remove gangue are called enrichment of ore or concentration of ore. Different separation techniques are employed based on the physical or chemical properties of the ore and gangue (e.g., hydraulic washing, magnetic separation, froth flotation).

Extracting Metals Low In The Activity Series

Metals at the bottom of the reactivity series are the least reactive. Their oxides can often be reduced to metal by simply heating the ore in air.

Examples:

• Mercury: Cinnabar (HgS), an ore of mercury, is heated in air. It first converts to mercuric oxide (HgO), which then further decomposes to mercury upon continued heating.
  • $2\text{HgS(s)} + 3\text{O}_2\text{(g)} \xrightarrow{\text{Heat}} 2\text{HgO(s)} + 2\text{SO}_2\text{(g)}$
  • $2\text{HgO(s)} \xrightarrow{\text{Heat}} 2\text{Hg(l)} + \text{O}_2\text{(g)}$
• Copper: Copper sulphide (Cu₂S) ore can be heated in air to obtain copper metal.
  • $2\text{Cu}_2\text{S} + 3\text{O}_2\text{(g)} \xrightarrow{\text{Heat}} 2\text{Cu}_2\text{O(s)} + 2\text{SO}_2\text{(g)}$
  • $2\text{Cu}_2\text{O} + \text{Cu}_2\text{S} \xrightarrow{\text{Heat}} 6\text{Cu(s)} + \text{SO}_2\text{(g)}$ (Autoreduction)

Extracting Metals In The Middle Of The Activity Series

Metals in the middle of the reactivity series (e.g., Zn, Fe, Pb, Cu) are moderately reactive and are typically found as sulphide or carbonate ores.

It is easier to obtain a metal from its oxide compared to its sulphide or carbonate. Therefore, sulphide and carbonate ores are first converted into metal oxides by heating:

• Roasting: Sulphide ores are strongly heated in the presence of excess air to convert them into oxides.

  Example: Roasting of zinc sulphide ore:

  $2\text{ZnS(s)} + 3\text{O}_2\text{(g)} \xrightarrow{\text{Heat}} 2\text{ZnO(s)} + 2\text{SO}_2\text{(g)}$

• Calcination: Carbonate ores are strongly heated in limited or no air to convert them into oxides.

  Example: Calcination of zinc carbonate ore:

  $\text{ZnCO}_3\text{(s)} \xrightarrow{\text{Heat}} \text{ZnO(s)} + \text{CO}_2\text{(g)}$

Once converted to oxides, the metal oxides are reduced to the corresponding metals using suitable reducing agents. Carbon (coke) is a commonly used reducing agent.

Example: Reduction of zinc oxide with carbon:

$\text{ZnO(s)} + \text{C(s)} \xrightarrow{\text{Heat}} \text{Zn(s)} + \text{CO(g)}$

Obtaining metals from their compounds by removing oxygen (or adding hydrogen) is a reduction process.

Highly reactive metals (like Na, Ca, Al) can also be used as reducing agents to displace less reactive metals from their oxides, especially when carbon is not effective. These displacement reactions are often highly exothermic, producing the metal in a molten state.

Example: Reaction of manganese dioxide with aluminium powder:

$3\text{MnO}_2\text{(s)} + 4\text{Al(s)} \longrightarrow 3\text{Mn(l)} + 2\text{Al}_2\text{O}_3\text{(s)} + \text{Heat}$

The reaction of iron(III) oxide ($\text{Fe}_2\text{O}_3$) with aluminium is known as the thermit reaction and is used for joining railway tracks or cracked machine parts. It produces molten iron.

$\text{Fe}_2\text{O}_3\text{(s)} + 2\text{Al(s)} \longrightarrow 2\text{Fe(l)} + \text{Al}_2\text{O}_3\text{(s)} + \text{Heat}$

Extracting Metals Towards The Top Of The Activity Series

Metals high in the reactivity series (e.g., K, Na, Ca, Mg, Al) are very reactive and have a strong affinity for oxygen. They cannot be reduced from their oxides using carbon or other common reducing agents, as these metals are more reactive than carbon.

These highly reactive metals are extracted by electrolytic reduction of their molten salts (chlorides or oxides). The metal is deposited at the cathode (negative electrode), and non-metals (like chlorine or oxygen) are liberated at the anode (positive electrode).

Example: Extraction of sodium from molten sodium chloride:

• At the cathode (negative electrode): $\text{Na}^+ + \text{e}^- \longrightarrow \text{Na}$ (Sodium ions gain electrons to form sodium metal)
• At the anode (positive electrode): $2\text{Cl}^- \longrightarrow \text{Cl}_2 + 2\text{e}^-$ (Chloride ions lose electrons to form chlorine gas)

Aluminium is obtained by the electrolytic reduction of aluminium oxide ($\text{Al}_2\text{O}_3$) dissolved in molten cryolite (to lower the melting point).

Refining Of Metals

Metals obtained from various extraction processes are often impure and contain impurities. These impurities must be removed to obtain pure metals. The process of purifying impure metals is called refining.

Electrolytic refining is the most widely used method for refining many common metals (like copper, zinc, tin, nickel, silver, gold).

Process of Electrolytic Refining:

• The impure metal is made the anode (positive electrode).
• A thin strip of pure metal is made the cathode (negative electrode).
• A solution of the metal salt (an electrolyte) is used.
• When electric current is passed, metal atoms from the impure anode lose electrons and dissolve into the electrolyte as metal ions ($\text{M} \longrightarrow \text{M}^{n+} + n\text{e}^-$).
• Metal ions from the electrolyte gain electrons at the cathode and are deposited as pure metal ($\text{M}^{n+} + n\text{e}^- \longrightarrow \text{M}$).
• Pure metal is transferred from the anode through the electrolyte and deposited on the cathode.
• Soluble impurities in the anode metal dissolve in the electrolyte.
• Insoluble impurities from the anode settle at the bottom as anode mud.