Comprehensive Study Guide on Metals and Non-metals

Physical Properties of Metals

• Metallic Lustre: Metals in their pure state possess a shining surface. This characteristic can be enhanced by cleaning the surface with sand paper to remove oxidation layers.

• Hardness: Metals are generally hard, though the degree of hardness varies between different metals.

  • Exceptions: Alkali metals like lithium, sodium, and potassium are exceptionally soft and can be cut easily with a knife. They also possess low densities and low melting points.

• Malleability: This is the property that allows metals to be beaten into thin sheets using a hammer. Gold and silver are considered the most malleable of all metals.

• Ductility: The ability of a metal to be drawn into thin wires is known as ductility.

  • Gold is the most ductile metal.

  • A wire of approximately $2\,km$ in length can be drawn from just $1\,g$ of gold.

• Thermal Conductivity: Metals are efficient conductors of heat and typically have high melting points.

  • Silver and copper are the best conductors of heat.

  • Lead and mercury are comparatively poor conductors of heat.

  • Melting Point Exceptions: Gallium and caesium have very low melting points and will melt if held in the palm of a hand.

• Electrical Conductivity: Metals conduct electricity. For safety, electrical wires used in homes are coated with insulating materials like polyvinylchloride (PVC) or rubber-like substances.

• Sonority: Metals that produce a distinct sound upon striking a hard surface are described as sonorous. This property explains why school bells are manufactured from metals.

• State at Room Temperature: Almost all metals are solids at room temperature.

  • Exception: Mercury is the only metal that exists as a liquid at room temperature.

Physical Properties of Non-metals

• Variety of States: Non-metals exist as either solids or gases.

  • Exception: Bromine is the only non-metal that is a liquid at room temperature.

• Examples: Common non-metals include carbon, sulphur, iodine, oxygen, and hydrogen.

• General Characteristics: Non-metals generally lack the malleability, ductility, and lustre associated with metals.

• Exceptions and Allotropes:

  • Iodine: Unlike most non-metals, iodine is lustrous.

  • Carbon Allotropes: Carbon exists in various forms called allotropes.

    • Diamond: An allotrope of carbon that is the hardest known natural substance. It possesses an extremely high melting and boiling point.

    • Graphite: An allotrope of carbon that serves as a good conductor of electricity.

Chemical Properties of Metals

Reaction with Oxygen (Burning in Air)

• Almost all metals combine with oxygen to form metal oxides, which are generally basic in nature.

  • $\text{Metal} + \text{Oxygen} \rightarrow \text{Metal oxide}$

• Specific Reactions:

  • Copper: When heated, it forms copper(II) oxide, a black substance.

    • $2Cu + O_2 \rightarrow 2CuO$

  • Aluminium: Forms aluminium oxide.

    • $4Al + 3O_2 \rightarrow 2Al_2O_3$

• Amphoteric Oxides: Some metal oxides, such as aluminium oxide and zinc oxide, exhibit both acidic and basic behavior. They react with both acids and bases to produce salt and water.

  • Reaction with Acid: $Al_2O_3 + 6HCl \rightarrow 2AlCl_3 + 3H_2O$

  • Reaction with Base: $Al_2O_3 + 2NaOH \rightarrow 2NaAlO_2 \text{ (Sodium aluminate)} + H_2O$

• Solubility and Alkalis: Most metal oxides are insoluble, but some dissolve in water to form alkalis.

  • $Na_2O(s) + H_2O(l) \rightarrow 2NaOH(aq)$

  • $K_2O(s) + H_2O(l) \rightarrow 2KOH(aq)$

• Reactivity Variations:

  • Sodium ($Na$) and Potassium ($K$) react so vigorously that they catch fire if exposed to air. They are stored immersed in kerosene oil for safety.

  • Magnesium ($Mg$), Aluminium ($Al$), Zinc ($Zn$), and Lead ($Pb$) form a thin protective oxide layer at ordinary temperatures, preventing further oxidation.

  • Iron ($Fe$) does not burn on heating, but iron filings burn vigorously in a flame.

  • Silver ($Ag$) and Gold ($Au$) do not react with oxygen even at high temperatures.

Reaction with Water

• Metals react with water to produce metal oxides and hydrogen gas. If the oxide is soluble, it further reacts to form a metal hydroxide.

  • $\text{Metal} + \text{Water} \rightarrow \text{Metal oxide} + \text{Hydrogen}$

  • $\text{Metal oxide} + \text{Water} \rightarrow \text{Metal hydroxide}$

• Reactivity Levels:

  • Sodium and Potassium: React violently with cold water. The reaction is highly exothermic, causing the evolved hydrogen to ignite.

    • $2K(s) + 2H_2O(l) \rightarrow 2KOH(aq) + H_2(g) + \text{heat energy}$

  • Calcium ($Ca$): Reacts less violently with water. The hydrogen gas bubbles stick to the metal surface, causing it to float.

  • Magnesium ($Mg$): Does not react with cold water but reacts with hot water to form magnesium hydroxide and hydrogen, causing it to float.

  • Aluminium, Iron, and Zinc: Do not react with cold or hot water but react with steam.

    • $2Al(s) + 3H_2O(g) \rightarrow Al_2O_3(s) + 3H_2(g)$

    • $3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)$

  • Lead ($Pb$), Copper ($Cu$), Silver ($Ag$), and Gold ($Au$): Do not react with water or steam at all.

Reaction with Acids

• Metals typically react with dilute acids to produce a salt and hydrogen gas.

  • $\text{Metal} + \text{Dilute acid} \rightarrow \text{Salt} + \text{Hydrogen}$

• Reactivity Order with $HCl$: Mg > Al > Zn > Fe. Copper does not react with dilute $HCl$.

• Nitric Acid ($HNO_3$): Hydrogen gas is not usually evolved with $HNO_3$ because it is a strong oxidising agent that oxidises resulting $H_2$ into $H_2O$.

  • Exception: Magnesium ($Mg$) and Manganese ($Mn$) react with very dilute $HNO_3$ to evolve $H_2$ gas.

• Aqua Regia: A highly corrosive fuming liquid consisting of concentrated hydrochloric acid and concentrated nitric acid in a $3:1$ ratio. It is capable of dissolving gold and platinum.

Displacement Reactions

• A more reactive metal can displace a less reactive metal from its salt solution.

  • $\text{Metal A} + \text{Salt solution of B} \rightarrow \text{Salt solution of A} + \text{Metal B}$

The Reactivity Series

The reactivity series is a list of metals arranged in decreasing order of their chemical activity:

1. Potassium ($K$): Most reactive

2. Sodium ($Na$)

3. Calcium ($Ca$)

4. Magnesium ($Mg$)

5. Aluminium ($Al$)

6. Zinc ($Zn$)

7. Iron ($Fe$)

8. Lead ($Pb$)

9. Hydrogen ($[H]$): Reference non-metal

10. Copper ($Cu$)

11. Mercury ($Hg$)

12. Silver ($Ag$)

13. Gold ($Au$): Least reactive

Reactions Between Metals and Non-metals

• Chemical Reactivity: Elements react to attain a stable, completely filled valence shell (octet), similar to noble gases.

• Ionic Compounds (Electrovalent Compounds): Formed by the transfer of electrons from a metal to a non-metal.

  • Formation of $NaCl$: Sodium ($2, 8, 1$) loses one electron to become $Na^+$; Chlorine ($2, 8, 7$) gains one electron to become $Cl^-$.

  • Formation of $MgCl_2$: Magnesium ($2, 8, 2$) loses two electrons to become $Mg^{2+}$; two chlorine atoms each gain one electron.

• Properties of Ionic Compounds:

  • Physical Nature: Crystalline solids, generally hard and brittle due to strong electrostatic forces.

  • Melting and Boiling Points: Very high (e.g., $NaCl$ melts at $1074\,K$ and boils at $1686\,K$).

  • Solubility: Generally soluble in water but insoluble in organic solvents like kerosene or petrol.

  • Conduction of Electricity: Do not conduct in solid state (no free ions); conduct in molten or aqueous state as ions are free to move.

Occurrence and Enrichment of Metals

• Minerals: Elements or compounds occurring naturally in the earth's crust.

• Ores: Minerals from which metals can be extracted profitably.

• Gangue: Large amounts of impurities like soil and sand found in mined ores.

• Metallurgy Groups:

  • Low Reactivity Metals: Found in free state (Gold, Silver, Platinum). Copper and silver also found as oxides/sulphides.

  • Medium Reactivity Metals: Found as oxides, sulphides, or carbonates ($Zn, Fe, Pb$).

  • High Reactivity Metals: Never found in free state ($K, Na, Ca, Mg, Al$).

Extraction of Metals

Metals Low in Activity Series

• Their oxides can be reduced by heating alone.

  • Cinnabar ($HgS$): $2HgS + 3O_2 \xrightarrow{\text{Heat}} 2HgO + 2SO_2 \rightarrow 2HgO \xrightarrow{\text{Heat}} 2Hg + O_2$

  • Copper Sulphide: $2Cu_2S + 3O_2 \xrightarrow{\text{Heat}} 2Cu_2O + 2SO_2 \rightarrow 2Cu_2O + Cu_2S \xrightarrow{\text{Heat}} 6Cu + SO_2$

Metals in the Middle of Activity Series

• Sulphide ores are converted to oxides via Roasting (heating in excess air).

  • $2ZnS + 3O_2 \xrightarrow{\text{Heat}} 2ZnO + 2SO_2$

• Carbonate ores are converted to oxides via Calcination (heating in limited air).

  • $ZnCO_3 \xrightarrow{\text{Heat}} ZnO + CO_2$

• Reduction: Metal oxides are reduced to metals using carbon (coke) or reactive metals (Aluminium).

  • $ZnO + C \rightarrow Zn + CO$

• Thermit Reaction: A highly exothermic displacement reaction used to join railway tracks.

  • $Fe_2O_3(s) + 2Al(s) \rightarrow 2Fe(l) + Al_2O_3(s) + \text{Heat}$

Metals High in Activity Series

• Obtained by Electrolytic Reduction of their molten chlorides.

  • Cathode (Negative): $Na^+ + e^- \rightarrow Na$

  • Anode (Positive): $2Cl^- \rightarrow Cl_2 + 2e^-$

Refining of Metals

• Electrolytic Refining: Used for copper, zinc, tin, nickel, silver, and gold.

  • Setup: Anode is impure metal; Cathode is a thin strip of pure metal; Electrolyte is a metal salt solution.

  • Result: Pure metal deposits on the cathode; soluble impurities enter the solution; insoluble impurities settle as anode mud.

Corrosion and Prevention

• Corrosion Examples:

  • Silver: Becomes black forming silver sulphide.

  • Copper: Gains a green coat of basic copper carbonate when exposed to moist $CO_2$.

  • Iron (Rusting): Requires both air and water. Forms a brown flaky substance called rust.

• Prevention Methods:

  • Painting, oiling, greasing, and chrome plating.

  • Anodising: Thickening the oxide layer on aluminium.

  • Galvanisation: Coating steel or iron with zinc. It protects even if the coating is damaged.

  • Alloying: Homogeneous mixture of metals or metals and non-metals.

    • Stainless Steel: Iron + Nickel + Chromium (hard, non-rusting).

    • Brass: Copper + Zinc.

    • Bronze: Copper + Tin.

    • Solder: Lead + Tin (low melting point for welding).

    • Amalgam: Any alloy involving mercury.

  • 22 Carat Gold: Used in India; 22 parts gold mixed with 2 parts copper or silver for hardness. (24 Carat is too soft for jewellery).

Questions & Discussion

• Why is school bells made of metals? Because metals are sonorous and produce sound when struck.

• Why are sodium and potassium stored under oil? They are highly reactive and catch fire when exposed to air/moisture.

• Can you write equations for burning magnesium and sulphur?

  • $2Mg + O_2 \rightarrow 2MgO$ (Basic oxide)

  • $S + O_2 \rightarrow SO_2$ (Acidic oxide)

• What is the solution used by the fake goldsmith to clean gold? Likely Aqua Regia, as it dissolves gold, explaining the drastic weight loss of the bangles.

• Why use copper for hot water tanks instead of steel? Copper is a better conductor of heat and does not react with water, whereas iron in steel can rust over time when exposed to steam/water.

• Distinguishing Metals and Non-metals: One can use a battery, bulb, wires, and switch to test for electrical conductivity. Metals will complete the circuit and light the bulb, while most non-metals (except graphite) will not.