Metals and Non-metalsClass 10 Chemistry Notes

Metals and Non-metals

Elements can be broadly classified into two main categories based on their properties: metals and non-metals. These properties determine how we use them in our daily lives, from the copper wires that carry electricity to the carbon that forms the basis of life.

PHYSICAL PROPERTIES

Metals

Metals can be identified and grouped based on their distinct physical characteristics.

• Metallic Lustre: In their pure state, metals have a shiny surface. This property is known as metallic lustre.
• Hardness: Metals are generally hard, although the degree of hardness varies from one metal to another.
• Malleability: This is the property that allows metals to be beaten into thin sheets. Gold and silver are the most malleable metals.
• Ductility: This is the ability of metals to be drawn into thin wires. Gold is the most ductile metal; a single gram of gold can be drawn into a wire about 2 km long. Malleability and ductility are the reasons metals can be shaped for various uses.
• Conductors of Heat and Electricity: Metals are excellent conductors of heat and electricity. This is why cooking vessels are made of metal. The best conductors of heat are silver and copper, while lead and mercury are relatively poor conductors. Their ability to conduct electricity is why they are used for electrical wiring.
• High Melting Points: Metals typically have high melting points, which allows them to remain solid even when heated to high temperatures, a useful property for cooking vessels.
• Sonorous: Metals produce a distinct ringing sound when they strike a hard surface. This property is called being sonorous, and it's why bells are made of metals.
• State: At room temperature, all metals are solids, with the exception of mercury, which is a liquid.

Non-metals

Non-metals are fewer in number compared to metals and have contrasting properties.

• Examples include carbon, sulphur, iodine, oxygen, and hydrogen.
• State: Non-metals exist as either solids or gases at room temperature. The only exception is bromine, which is a liquid.

Exceptions to General Physical Properties

While these general properties are useful for classification, there are several important exceptions that show we cannot group elements based on physical properties alone.

• State and Melting Point of Metals: While most metals are hard solids with high melting points, there are exceptions.
  • Mercury is a liquid at room temperature.
  • Alkali metals like lithium, sodium, and potassium are so soft they can be cut with a knife. They also have low densities and low melting points.
  • Gallium and caesium have such low melting points that they will melt if you keep them on your palm.
• Lustre of Non-metals: Although non-metals are typically dull, iodine is a non-metal that is lustrous (shiny).
• Hardness and Conductivity of Non-metals:
  • Carbon is a non-metal that can exist in different forms called allotropes.
  • 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 good conductor of electricity, which is unusual for a non-metal.

CHEMICAL PROPERTIES OF METALS

The chemical properties of elements provide a clearer basis for classifying them as metals or non-metals. Generally, metals form basic oxides, while non-metals form acidic oxides.

What happens when Metals are burnt in Air?

Almost all metals combine with oxygen to form metal oxides. Metal + Oxygen → Metal oxide

• For example, when copper is heated, it combines with oxygen to form a black substance, copper(II) oxide. $2 \mathrm{Cu}+\mathrm{O}_{2} \rightarrow 2 \mathrm{CuO}$
• Similarly, aluminium forms aluminium oxide. $4 \mathrm{Al}+3 \mathrm{O}_{2} \rightarrow 2 \mathrm{Al}_{2} \mathrm{O}_{3}$

Nature of Metal Oxides Metal oxides are generally basic in nature. Some, like sodium oxide ($Na_2O$) and potassium oxide ($K_2O$), are soluble in water and form alkalis (bases that dissolve in water). $\mathrm{Na}_{2} \mathrm{O}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow 2 \mathrm{NaOH}(\mathrm{aq})$ $\mathrm{K}_{2} \mathrm{O}(\mathrm{~s})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow 2 \mathrm{KOH}(\mathrm{aq})$

However, some metal oxides, like aluminium oxide and zinc oxide, exhibit both acidic and basic properties. These are known as amphoteric oxides. They can react with both acids and bases to produce salt and water.

• Reaction with acid: $\mathrm{Al}_{2} \mathrm{O}_{3}+6 \mathrm{HCl} \rightarrow 2 \mathrm{AlCl}_{3}+3 \mathrm{H}_{2} \mathrm{O}$
• Reaction with base: $\mathrm{Al}_{2} \mathrm{O}_{3}+2 \mathrm{NaOH} \rightarrow 2 \mathrm{NaAlO}_{2} \text{ (Sodium aluminate)} +\mathrm{H}_{2} \mathrm{O}$

Reactivity with Oxygen Varies Metals show different levels of reactivity towards oxygen.

• Potassium (K) and Sodium (Na) react so vigorously that they catch fire in open air. To prevent this, they are stored immersed in kerosene oil.
• Magnesium (Mg), Aluminium (Al), Zinc (Zn), and Lead (Pb) form a thin, protective layer of oxide on their surface at room temperature. This layer prevents further oxidation (corrosion).
• Iron (Fe) does not burn on heating, but iron filings burn vigorously when sprinkled in a flame.
• Copper (Cu) does not burn but gets coated with a black layer of copper(II) oxide when heated.
• Silver (Ag) and Gold (Au) do not react with oxygen, even at high temperatures.

What happens when Metals react with Water?

Metals react with water to produce a metal oxide and hydrogen gas. If the metal oxide is soluble in water, it further dissolves to form a metal hydroxide. Metal + Water → Metal oxide + Hydrogen Metal oxide + Water → Metal hydroxide

The reactivity with water also varies greatly among metals.

• Reaction with Cold Water:

  • Potassium and sodium react violently with cold water. The reaction is so exothermic (releases so much heat) that the hydrogen gas produced immediately catches fire. $2 \mathrm{~K}(\mathrm{~s})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow 2 \mathrm{KOH}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{~g})+\text { heat energy }$ $2 \mathrm{Na}(\mathrm{~s})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow 2 \mathrm{NaOH}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{~g})+\text { heat energy }$
  • Calcium's reaction is less violent, and the heat produced is not enough to ignite the hydrogen. Calcium starts floating because the bubbles of hydrogen gas stick to its surface. $\mathrm{Ca}(\mathrm{~s})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{Ca}(\mathrm{OH})_{2}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{~g})$

• Reaction with Hot Water:

  • Magnesium does not react with cold water but reacts with hot water to form magnesium hydroxide and hydrogen. It also floats due to hydrogen bubbles sticking to its surface.

• Reaction with Steam:

  • Metals like aluminium, iron, and zinc do not react with cold or hot water. They react with steam to form the metal oxide and hydrogen. $2 \mathrm{Al}(\mathrm{~s})+3 \mathrm{H}_{2} \mathrm{O}(\mathrm{g}) \rightarrow \mathrm{Al}_{2} \mathrm{O}_{3}(\mathrm{~s})+3 \mathrm{H}_{2}(\mathrm{~g})$ $3 \mathrm{Fe}(\mathrm{~s})+4 \mathrm{H}_{2} \mathrm{O}(\mathrm{g}) \rightarrow \mathrm{Fe}_{3} \mathrm{O}_{4}(\mathrm{~s})+4 \mathrm{H}_{2}(\mathrm{~g})$

• No Reaction with Water:

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

What happens when Metals react with Acids?

Metals generally react with dilute acids to produce a salt and hydrogen gas. Metal + Dilute acid → Salt + Hydrogen

The vigor of the reaction depends on the metal's reactivity.

• The rate of bubble formation (hydrogen gas) indicates the speed of the reaction. The reactivity order with dilute hydrochloric acid is: Mg > Al > Zn > Fe.
• In the case of copper, no bubbles are seen, indicating it does not react with dilute HCl.

How do Metals react with Solutions of other Metal Salts?

A more reactive metal can displace a less reactive metal from its compound in a solution or molten form. This is known as a displacement reaction.

Metal A + Salt solution of B → Salt solution of A + Metal B

This principle provides excellent evidence for the relative reactivities of metals. If metal A displaces metal B, then metal A is more reactive than metal B.

The Reactivity Series

The reactivity series is a list of metals arranged in the order of their decreasing chemical activity. It helps predict the outcomes of displacement reactions.

How do Metals and Non-metals React?

The reactions between metals and non-metals are governed by their tendency to achieve a stable electron configuration, like that of a noble gas (a full outer electron shell).

• Metals have 1, 2, or 3 electrons in their outermost shell. They tend to lose these valence electrons to form positively charged ions, called cations.
• Non-metals typically have 4, 5, 6, or 7 electrons in their outermost shell. They tend to gain electrons to complete their outer shell, forming negatively charged ions, called anions.

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

Metals and Non-metals

Elements, the basic building blocks of matter, can be broadly classified into two main groups: metals and non-metals. This classification is based on their distinct physical and chemical properties, which also determine their uses in our daily lives.

PHYSICAL PROPERTIES

Metals

Let's explore the common physical characteristics that help us identify metals.

• Metallic Lustre: Metals, in their pure form, have a shiny surface. This property is called metallic lustre. When metals like iron or copper look dull, it's often because of a surface reaction. Rubbing them with sandpaper reveals their shiny nature.

• Hardness: Metals are generally hard. However, the degree of hardness can vary significantly from one metal to another. For example, iron is very hard, but alkali metals like sodium, lithium, and potassium are so soft they can be easily cut with a knife.

• Malleability: This is the property that allows metals to be beaten into thin sheets. If you strike a piece of iron with a hammer, it flattens instead of shattering. Gold and silver are the most malleable metals.

• Ductility: This is the ability of metals to be drawn into thin wires. Most wires you see, like those used for electricity, are made of metals like copper or aluminium. Gold is the most ductile metal; a single gram of gold can be drawn into a wire about 2 km long!

• Good Conductors of Heat: Metals transfer heat efficiently. This is why cooking vessels are made of metals like aluminium or copper. The best conductors of heat are silver and copper. In contrast, lead and mercury are relatively poor conductors of heat.

• Good Conductors of Electricity: Metals allow electricity to pass through them easily. This is why electrical wires are made from metals. However, these wires are coated with materials like polyvinylchloride (PVC) or rubber, which are insulators and prevent electric shocks.

• Sonorous: Metals produce a ringing sound when they are struck against a hard surface. This property is called being sonorous. It's the reason school bells are made of metal.

• State: Most metals are solid at room temperature. The only exception is mercury, which is a liquid.

Non-metals

Non-metals have properties that are generally the opposite of metals.

• Examples include carbon, sulphur, iodine, oxygen, and hydrogen.
• They exist as either solids or gases at room temperature, with the exception of bromine, which is a liquid.
• They are generally not lustrous, malleable, ductile, or sonorous.
• They are typically poor conductors of heat and electricity.

Exceptions to General Physical Properties

While these properties are useful for classification, there are many exceptions that show us why physical properties alone are not enough to group elements.

• State: All metals except mercury are solids at room temperature. However, gallium and caesium have such low melting points that they will melt on your palm.
• Lustre: Although non-metals are typically dull, iodine is a non-metal that is naturally lustrous (shiny).
• Hardness: Carbon, a non-metal, has different forms called allotropes. One allotrope, diamond, is the hardest natural substance known and has an extremely high melting and boiling point. Alkali metals like sodium and potassium are metals, but they are very soft.
• Conductivity: While non-metals are poor conductors of electricity, graphite, another allotrope of carbon, is an excellent conductor.

CHEMICAL PROPERTIES OF METALS

Chemical properties provide a clearer way to classify elements.

What happens when Metals are burnt in Air?

Almost all metals react with oxygen to form metal oxides. $\text{Metal} + \text{Oxygen} \rightarrow \text{Metal oxide}$ However, different metals react with oxygen at different rates.

• Nature of Metal Oxides: Most metal oxides are basic in nature. When dissolved in water, they form bases that turn red litmus paper blue.

  Note

  Some metal oxides, like sodium oxide ($Na_2O$) and potassium oxide ($K_2O$), are soluble in water and form alkalis (soluble bases). $\mathrm{Na}_{2} \mathrm{O}(\mathrm{~s})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow 2 \mathrm{NaOH}(\mathrm{aq})$ $\mathrm{K}_{2} \mathrm{O}(\mathrm{~s})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow 2 \mathrm{KOH}(\mathrm{aq})$

• Amphoteric Oxides: Some metal oxides, such as aluminium oxide ($Al_2O_3$) and zinc oxide (ZnO), exhibit both acidic and basic properties. These are called amphoteric oxides. They can react with both acids and bases to form salt and water.

  • Reaction with acid: $\mathrm{Al}_{2} \mathrm{O}_{3}+6 \mathrm{HCl} \rightarrow 2 \mathrm{AlCl}_{3}+3 \mathrm{H}_{2} \mathrm{O}$
  • Reaction with base: $\mathrm{Al}_{2} \mathrm{O}_{3}+2 \mathrm{NaOH} \rightarrow 2 \mathrm{NaAlO}_{2} \text{ (Sodium aluminate)} +\mathrm{H}_{2} \mathrm{O}$

• Reactivity with Oxygen:

  • Potassium (K) and Sodium (Na) react so vigorously with oxygen that they catch fire if left in the open. To prevent this, they are stored under kerosene oil.
  • Magnesium (Mg), Aluminium (Al), Zinc (Zn), and Lead (Pb) are covered with a thin, protective layer of oxide at ordinary temperatures. This layer prevents the metal from further oxidation.
  • Iron (Fe) does not burn on heating, but iron filings burn vigorously when sprinkled into a flame.
  • Copper (Cu) does not burn but gets coated with a black layer of copper(II) oxide ($CuO$) when heated.
  • Silver (Ag) and Gold (Au) do not react with oxygen, even at high temperatures.

Anodising

Anodising is an industrial process used to create a thick, protective oxide layer on aluminium. An aluminium article is made the anode in an electrolytic cell with dilute sulphuric acid. The oxygen gas produced at the anode reacts with the aluminium, forming a thicker, more durable oxide layer. This layer protects the aluminium from corrosion and can be dyed to give it an attractive finish.

What happens when Metals react with Water?

Metals react with water to produce a metal oxide and hydrogen gas. If the metal oxide is soluble in water, it further dissolves to form a metal hydroxide. $\text{Metal} + \text{Water} \rightarrow \text{Metal oxide} + \text{Hydrogen}$ $\text{Metal oxide} + \text{Water} \rightarrow \text{Metal hydroxide}$ The reactivity of metals with water varies greatly:

• Metals that react with cold water:

  • Potassium (K) and Sodium (Na) react violently with cold water. The reaction is so exothermic (releases so much heat) that the hydrogen gas produced immediately catches fire. $2 \mathrm{~K}(\mathrm{~s})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow 2 \mathrm{KOH}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{~g})+\text { heat energy }$ $2 \mathrm{Na}(\mathrm{~s})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow 2 \mathrm{NaOH}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{~g})+\text { heat energy }$
  • Calcium (Ca) reacts less violently. The heat produced is not enough to ignite the hydrogen. Calcium starts to float because bubbles of hydrogen gas stick to its surface. $\mathrm{Ca}(\mathrm{~s})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{Ca}(\mathrm{OH})_{2}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{~g})$

• Metals that react with hot water:

  • Magnesium (Mg) does not react with cold water but reacts with hot water to form magnesium hydroxide and hydrogen. It also floats due to hydrogen bubbles sticking to its surface.

• Metals that react with steam:

  • Metals like Aluminium (Al), Iron (Fe), and Zinc (Zn) do not react with cold or hot water. They react with steam to form the metal oxide and hydrogen. $2 \mathrm{Al}(\mathrm{~s})+3 \mathrm{H}_{2} \mathrm{O}(\mathrm{~g}) \rightarrow \mathrm{Al}_{2} \mathrm{O}_{3}(\mathrm{~s})+3 \mathrm{H}_{2}(\mathrm{~g})$ $3 \mathrm{Fe}(\mathrm{~s})+4 \mathrm{H}_{2} \mathrm{O}(\mathrm{~g}) \rightarrow \mathrm{Fe}_{3} \mathrm{O}_{4}(\mathrm{~s})+4 \mathrm{H}_{2}(\mathrm{~g})$

• Metals that do not react with water:

  • Lead (Pb), Copper (Cu), Silver (Ag), and Gold (Au) do not react with water at all.

What happens when Metals react with Acids?

Metals generally react with dilute acids to produce a salt and hydrogen gas. $\text{Metal} + \text{Dilute acid} \rightarrow \text{Salt} + \text{Hydrogen}$ The vigour of the reaction depends on the metal's reactivity. For dilute hydrochloric acid, the order of reactivity is: Mg > Al > Zn > Fe. Copper does not react with dilute HCl.

Aqua Regia

Aqua regia (Latin for "royal water") is a freshly prepared, highly corrosive, fuming liquid. It is a mixture of concentrated hydrochloric acid (HCl) and concentrated nitric acid ($HNO_3$) in a ratio of 3:1. It is one of the few substances that can dissolve noble metals like gold and platinum, which are unreactive to either acid alone.

How do Metals react with Solutions of other Metal Salts?

A more reactive metal can displace a less reactive metal from its salt solution. This is a type of chemical reaction known as a displacement reaction. $\text{Metal A} + \text{Salt solution of B} \rightarrow \text{Salt solution of A} + \text{Metal B}$ This principle gives us a clear way to compare the reactivities of metals. For example, if you place an iron nail in a copper sulphate solution, the iron will displace the copper because iron is more reactive than copper. $\text{Fe(s) + CuSO}_4\text{(aq)} \rightarrow \text{FeSO}_4\text{(aq) + Cu(s)}$

The Reactivity Series

By studying displacement reactions, scientists have arranged metals in a list according to their decreasing reactivity. This list is called the reactivity series or activity series.

Metals above hydrogen in the series can displace hydrogen from dilute acids, while those below cannot.

How do Metals and Non-metals React?

The reaction between metals and non-metals can be understood by looking at their electron configurations. Atoms react to achieve a stable electron configuration, usually a full outer shell (an octet), like the noble gases.

• Metals tend to have 1, 2, or 3 electrons in their outermost shell. They can lose these electrons to form positively charged ions, called cations.
• Non-metals tend to have 5, 6, or 7 electrons in their outermost shell. They can gain electrons to form negatively charged ions, called anions.

The compounds formed by the transfer of electrons from a metal to a non-metal are known as ionic compounds or electrovalent compounds.

• A sodium atom (Na) has an electron configuration of 2, 8, 1. It loses one electron to become a sodium ion ($Na^+$) with a stable configuration of 2, 8.
• A chlorine atom (Cl) has an electron configuration of 2, 8, 7. It gains one electron to become a chloride ion ($Cl^−$) with a stable configuration of 2, 8, 8.
• The oppositely charged ions, $Na^+$ and $Cl^−$, are held together by strong electrostatic forces of attraction to form sodium chloride (NaCl).

• A magnesium atom (Mg) has a configuration of 2, 8, 2. It loses two electrons to form a magnesium ion ($Mg^{2+}$).
• Each of the two chlorine atoms (Cl) gains one electron to form two chloride ions ($Cl^−$).
• The cation is $Mg^{2+}$ and the anion is $Cl^−$. These are held together to form magnesium chloride ($MgCl_2$).

Properties of Ionic Compounds

Ionic compounds have distinct properties due to the strong forces between their ions.

1. Physical Nature: They are typically hard, crystalline solids because of the strong attraction between positive and negative ions. They are also brittle, meaning they break into pieces when pressure is applied.
2. Melting and Boiling Points: Ionic compounds have very high melting and boiling points. A large amount of energy is required to overcome the strong inter-ionic attraction and break apart the crystal structure.
3. Solubility: They are generally soluble in water but insoluble in organic solvents like kerosene and petrol.
4. Conduction of Electricity:
   • In the solid state, ionic compounds do not conduct electricity because the ions are held in fixed positions and cannot move.
   • In the molten state or when dissolved in water, they conduct electricity. The heat or the water breaks down the rigid structure, allowing the ions to move freely and carry charge.

OCCURRENCE OF METALS

The Earth's crust is the primary source of metals. Seawater also contains dissolved salts like sodium chloride and magnesium chloride.

• Minerals: The elements or compounds that occur naturally in the Earth's crust are called minerals.
• Ores: Minerals from which a metal can be extracted profitably are called ores. For example, aluminium is extracted from its ore, bauxite.
• Gangue: Ores mined from the earth are often contaminated with impurities like soil and sand. These impurities are called gangue.

Extraction of Metals

Metallurgy is the science and technology of extracting metals from their ores and refining them for use. The method used for extraction depends on the metal's position in the reactivity series.

Enrichment of Ores

The first step in metallurgy is the removal of the gangue from the ore. This process is called enrichment or concentration of the ore.

Extracting Metals Low in the Activity Series

Metals like mercury, copper, silver, and gold are low in the reactivity series and are very unreactive.

• Their oxides can be reduced to metal by heating alone.
• Example (Mercury): Cinnabar (HgS), an ore of mercury, is heated in air. It first turns into mercuric oxide (HgO), which then reduces to mercury (Hg) upon further heating. $2 \mathrm{HgS}(\mathrm{~s})+3 \mathrm{O}_{2}(\mathrm{~g}) \xrightarrow{\text { Heat }} 2 \mathrm{HgO}(\mathrm{~s})+2 \mathrm{SO}_{2}(\mathrm{~g})$ $2 \mathrm{HgO}(\mathrm{~s}) \xrightarrow{\text { Heat }} 2 \mathrm{Hg}(\mathrm{l})+\mathrm{O}_{2}(\mathrm{~g})$

Extracting Metals in the Middle of the Activity Series

Metals like iron, zinc, and lead are moderately reactive. They are usually found as sulphide or carbonate ores. It is easier to obtain a metal from its oxide than from its sulphide or carbonate.

1. Conversion to Oxide:

   • Roasting: Sulphide ores are converted into oxides by heating strongly in the presence of excess air. $2 \mathrm{ZnS}(\mathrm{~s})+3 \mathrm{O}_{2}(\mathrm{~g}) \xrightarrow{\text { Heat }} 2 \mathrm{ZnO}(\mathrm{~s})+2 \mathrm{SO}_{2}(\mathrm{~g})$
   • Calcination: Carbonate ores are converted into oxides by heating strongly in a limited supply of air. $\mathrm{ZnCO}_{3}(\mathrm{~s}) \xrightarrow{\text { Heat }} \mathrm{ZnO}(\mathrm{~s})+\mathrm{CO}_{2}(\mathrm{~g})$

2. Reduction to Metal: The metal oxides are then reduced to metal using a reducing agent, most commonly carbon (coke). $\mathrm{ZnO}(\mathrm{~s})+\mathrm{C}(\mathrm{~s}) \rightarrow \mathrm{Zn}(\mathrm{~s})+\mathrm{CO}(\mathrm{~g})$

   Sometimes, more reactive metals like sodium, calcium, or aluminium are used as reducing agents. These displacement reactions are highly exothermic.

   Example

   The Thermit Reaction The reaction between iron(III) oxide ($Fe_2O_3$) and aluminium (Al) is extremely exothermic, producing molten iron. This reaction, known as the thermit reaction, is used to weld railway tracks or join cracked machine parts. $\mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{~s})+2 \mathrm{Al}(\mathrm{~s}) \rightarrow 2 \mathrm{Fe}(\mathrm{l})+\mathrm{Al}_{2} \mathrm{O}_{3}(\mathrm{~s})+\text { Heat }$

Extracting Metals towards the Top of the Activity Series

Metals high in the reactivity series (K, Na, Ca, Mg, Al) are very reactive. They have a strong affinity for oxygen and cannot be reduced by carbon.

• These metals are extracted using electrolytic reduction (electrolysis) of their molten chlorides or oxides.
• During electrolysis, the metal ions are attracted to the negatively charged electrode (cathode), where they gain electrons and are deposited as pure metal.
• Example (Sodium): In the electrolysis of molten NaCl:
  • At the cathode: $\mathrm{Na}^{+}+\mathrm{e}^{-} \rightarrow \mathrm{Na}$ (metal deposited)
  • At the anode: $2 \mathrm{Cl}^{-} \rightarrow \mathrm{Cl}_{2}+2 \mathrm{e}^{-}$ (gas liberated)

Refining of Metals

The metals produced by reduction processes are often impure. The most common method for purifying them is electrolytic refining.

This process is used for metals like copper, zinc, tin, nickel, silver, and gold.

• Anode: A thick block of the impure metal.
• Cathode: A thin strip of the pure metal.
• Electrolyte: A solution of a salt of the metal being refined (e.g., acidified copper sulphate for refining copper).

When electric current is passed, the impure metal from the anode dissolves into the electrolyte. An equivalent amount of pure metal from the electrolyte then deposits onto the cathode. Soluble impurities dissolve in the solution, while insoluble impurities settle at the bottom of the anode as anode mud.

CORROSION

Corrosion is the gradual deterioration of a metal due to its reaction with substances in its environment, such as air, water, and chemicals.

• Silver: Tarnishes and turns black due to the formation of silver sulphide ($Ag_2S$) from reacting with sulphur compounds in the air.
• Copper: Slowly reacts with moist carbon dioxide in the air to form a green coat of basic copper carbonate.
• Iron: When exposed to moist air, it acquires a brown, flaky coating called rust. For iron to rust, both air (oxygen) and water must be present.

Prevention of Corrosion

Preventing the surface of a metal from coming into contact with air and moisture can stop corrosion. Common methods include:

• Painting, oiling, or greasing: Creates a barrier on the surface.
• Galvanisation: Coating iron or steel with a thin layer of zinc. Zinc is more reactive than iron, so it corrodes first, protecting the iron. This is known as sacrificial protection.
• Chrome plating and anodising: Coating with another metal or a stable oxide layer.
• Alloying: Mixing a metal with other metals or non-metals to improve its properties.

Alloys

An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal. Alloying can change a metal's properties, making it stronger, harder, or more resistant to corrosion.

• Steel: Pure iron is soft and stretches easily. Adding a small amount of carbon (about 0.05%) makes it hard and strong, creating steel.
• Stainless Steel: Mixing iron with nickel and chromium produces stainless steel, which is hard and does not rust.
• Jewellery: Pure gold (24 carat) is very soft. It is alloyed with silver or copper to make it harder and suitable for jewellery (e.g., 22 carat gold).