Comprehensive Notes on Metals and Non-metals

Metals and Non-metals

Introduction

• Elements are classified as metals or non-metals based on their properties.
• Metals and non-metals have various uses in daily life.

Physical Properties of Metals

Activity 3.1: Metallic Lustre

• Metals in their pure state have a shining surface called metallic lustre.
• Example metals used: iron, copper, aluminium, magnesium.

Activity 3.2: Malleability

• Some metals can be beaten into thin sheets; this property is called malleability.
• Gold and silver are the most malleable metals.

Activity 3.3: Hardness

• Metals are generally hard, but hardness varies.
• Metals used: iron, zinc, lead, copper.

Activity 3.4: Ductility

• The ability of metals to be drawn into thin wires is called ductility.
• Gold is the most ductile metal; 1 gram can be drawn into a 2 km long wire.
• Metals can be shaped according to needs due to malleability and ductility.

Activity 3.5: Conductivity of Heat

• Metals are good conductors of heat and have high melting points.
• Silver and copper are the best conductors; lead and mercury are poor conductors.

Conductivity of Electricity

• Electric wires are coated with polyvinylchloride (PVC) or a rubber-like material.
• Metals produce a sound when struck, called sonorous.
• School bells are made of metals due to their sonorous property.

Physical Properties of Non-metals

• Non-metals can be solids or gases, except bromine, which is a liquid.
• Non-metals do not typically share the same physical properties as metals.

Activity 3.6: Electric Circuit

• Testing conductivity by placing a material in a circuit between terminals A and B.
• If the bulb glows, the material conducts electricity.

Activity 3.7

• Non-metal samples: carbon (coal or graphite), sulphur, and iodine.
• Observations are compiled in a table to compare metals and non-metals focusing on hardness, malleability, ductility, conductivity, and sonority.

Exceptions in Physical Properties

• Elements cannot be grouped solely based on physical properties; there are exceptions.
• All metals are solid at room temperature except mercury.
• Gallium and caesium have very low melting points.
• Iodine is a lustrous non-metal.
• Carbon exists in different allotropes (diamond and graphite).
• Diamond is the hardest natural substance with a high melting and boiling point.
• Graphite is a conductor of electricity.
• Alkali metals (lithium, sodium, potassium) are soft and can be cut with a knife; they also have low densities and melting points.

Chemical Properties of Metals

Activity 3.8

• Magnesium ribbon and sulphur powder are burned to test acidic or basic properties.
• Burning magnesium produces a basic product.
• Burning sulphur produces an acidic product.

Question Examples

• Example of a metal liquid at room temperature: mercury.
• Metal easily cut with a knife: sodium.
• Best conductor of heat: silver or copper.
• Poor conductor of heat: lead.
• Malleable: Ability to be beaten into thin sheets.
• Ductile: Ability to be drawn into thin wires.

Reactions with Air

Activity 3.9

• Metals burn in air to form metal oxides: $Metal + Oxygen \rightarrow Metal \space oxide$
• Copper heated in air forms copper(II) oxide: $2Cu + O_2 \rightarrow 2CuO$
• Aluminium forms aluminium oxide: $4Al + 3O<em>2 \rightarrow 2Al</em>2O_3$
• Metal oxides are generally basic.
• Amphoteric oxides (e.g., aluminium oxide, zinc oxide) react with both acids and bases:
  • $Al<em>2O</em>3 + 6HCl \rightarrow 2AlCl<em>3 + 3H</em>2O$
  • $Al<em>2O</em>3 + 2NaOH \rightarrow 2NaAlO<em>2 + H</em>2O$
• Sodium oxide and potassium oxide dissolve in water to form alkalis:
  • $Na<em>2O(s) + H</em>2O(l) \rightarrow 2NaOH(aq)$
  • $K<em>2O(s) + H</em>2O(l) \rightarrow 2KOH(aq)$

Reactivity with Oxygen

• Different metals show different reactivities towards oxygen.
• Potassium and sodium react vigorously and are stored in kerosene oil.
• Magnesium, aluminium, zinc, and lead form a protective oxide layer preventing further oxidation.
• Iron filings burn vigorously when sprinkled in a flame.
• Copper is coated with a black layer of copper(II) oxide when heated.
• Silver and gold do not react with oxygen even at high temperatures.

Anodising

• Anodising is forming a thick oxide layer of aluminium for corrosion resistance using dilute sulphuric acid during electrolysis.

Reactions with Water

Activity 3.10

• Metals react with water to produce metal oxide and hydrogen gas: $Metal + Water \rightarrow Metal \space oxide + Hydrogen$
• Metal oxides soluble in water form metal hydroxides: $Metal \space oxide + Water \rightarrow Metal \space hydroxide$
• Potassium and sodium react violently with cold water:
  • $2K(s) + 2H<em>2O(l) \rightarrow 2KOH(aq) + H</em>2(g) + heat \space energy$
  • $2Na(s) + 2H<em>2O(l) \rightarrow 2NaOH(aq) + H</em>2(g) + heat \space energy$
• Calcium reacts less violently:
  • $Ca(s) + 2H<em>2O(l) \rightarrow Ca(OH)</em>2(aq) + H_2(g)$
• Magnesium reacts with hot water to form magnesium hydroxide and hydrogen.
• Aluminium, iron, and zinc react with steam:
  • $2Al(s) + 3H<em>2O(g) \rightarrow Al</em>2O<em>3(s) + 3H</em>2(g)$
  • $3Fe(s) + 4H<em>2O(g) \rightarrow Fe</em>3O<em>4(s) + 4H</em>2(g)$
• Lead, copper, silver, and gold do not react with water.

Reactions with Acids

• Metals react with dilute acids to give salt and hydrogen gas: $Metal + Dilute \space acid \rightarrow Salt + Hydrogen$

Activity 3.11

• Testing reaction rates with hydrochloric acid. Magnesium reacts fastest.
• Nitric acid ((HNO3)) is a strong oxidising agent and oxidises (H2) to water, reducing itself to nitrogen oxides.
• Magnesium ((Mg)) and manganese ((Mn)) react with very dilute (HNO3) to evolve (H2) gas.

Aqua Regia

• Aqua regia is a mixture of concentrated hydrochloric acid and concentrated nitric acid in a 3:1 ratio that can dissolve gold and platinum.

Reactions with Metal Salt Solutions

Activity 3.12

• Reactive metals displace less reactive metals from their compounds in solution.
• $Metal \space A + Salt \space solution \space of \space B \rightarrow Salt \space solution \space of \space A + Metal \space B$

Reactivity Series

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

How Metals and Non-metals React

• Noble gases have completely filled valence shells and show little chemical activity.
• Elements react to attain a completely filled valence shell.
• Sodium atom loses an electron to form (Na^+) cation.
• Chlorine gains an electron to form (Cl^-) anion.

Formation of Ionic Compounds

• Sodium chloride ((NaCl)) is formed by electrostatic attraction between (Na^+) and (Cl^-) ions.
• Magnesium chloride ((MgCl_2)) is another ionic compound formed by electron transfer.

Properties of Ionic Compounds

Activity 3.13

• Ionic compounds are solids with strong inter-ionic forces.
• They have high melting and boiling points.
• Generally soluble in water and insoluble in solvents like kerosene and petrol.
• Conduct electricity in solution or molten state but not in solid-state.

Occurrence of Metals

• The earth’s crust is the major source of metals.
• Minerals are elements or compounds in the earth’s crust.
• Ores are minerals with a high percentage of a particular metal.

Extraction of Metals

• Metals at the bottom of the activity series are found in a free state (e.g., gold, silver, platinum).
• Metals at the top of the series are highly reactive and found as compounds.
• Metals in the middle are moderately reactive and found as oxides, sulphides, or carbonates.

Enrichment of Ores

• Ores are contaminated with impurities like soil and sand, called gangue.
• Gangue is removed based on physical or chemical properties.

Extracting Metals Low in the Activity Series

• Oxides of these metals are reduced by heating alone.
• Cinnabar ((HgS)) is heated in air to form mercuric oxide ((HgO)), then reduced to mercury.
  • $2HgS(s) + 3O<em>2(g) \rightarrow 2HgO(s) + 2SO</em>2(g)$
  • $2HgO(s) \rightarrow 2Hg(l) + O_2(g)$
• (Cu_2S) is heated in air to obtain copper.
  • $2Cu<em>2S + 3O</em>2(g) \rightarrow 2Cu<em>2O(s) + 2SO</em>2(g)$
  • $2Cu<em>2O + Cu</em>2S \rightarrow 6Cu(s) + SO_2(g)$

Extracting Metals in the Middle of the Activity Series

• These metals are present as sulphides or carbonates.
• Sulphides are converted to oxides by roasting (heating in excess air).
• Carbonates are converted to oxides by calcination (heating in limited air).

Roasting and Calcination Examples

• Roasting: $2ZnS(s) + 3O<em>2(g) \rightarrow 2ZnO(s) + 2SO</em>2(g)$
• Calcination: $ZnCO<em>3(s) \rightarrow ZnO(s) + CO</em>2(g)$
• Metal oxides are reduced using reducing agents like carbon: $ZnO(s) + C(s) \rightarrow Zn(s) + CO(g)$

Displacement Reactions

• Highly reactive metals (Na, Ca, Al) displace less reactive metals.
• Thermit reaction: $Fe<em>2O</em>3(s) + 2Al(s) \rightarrow 2Fe(l) + Al<em>2O</em>3(s) + Heat$

Extracting Metals towards the Top of the Activity Series

• These metals are obtained by electrolytic reduction.
• Sodium, magnesium, and calcium are obtained by electrolysis of their molten chlorides.
  • At cathode: $Na^+ + e^- \rightarrow Na$
  • At anode: $2Cl^- \rightarrow Cl_2 + 2e^-$
• Aluminium is obtained by electrolytic reduction of aluminium oxide.

Refining of Metals

• Electrolytic refining is widely used for metals like copper, zinc, tin, nickel, silver, and gold.
  • Impure metal is the anode, pure metal is the cathode.
  • Metal salt solution is the electrolyte.
  • Pure metal from the anode dissolves and deposits on the cathode.
  • Insoluble impurities settle as anode mud.

Corrosion

• Silver articles become black due to reaction with sulphur: $Ag + S \rightarrow Ag_2S$
• Copper reacts with moist (CO2) to form basic copper carbonate: $Cu + CO</em>2 + H<em>2O \rightarrow CuCO</em>3 "." Cu(OH)_2$
• Iron rusts when exposed to moist air.

Activity 3.14

• Investigating conditions for iron rusting.
  • Tube A: Iron nails rust due to air and water.
  • Tube B: No rust due to oil preventing air dissolving in water.
  • Tube C: No rust due to anhydrous calcium chloride absorbing moisture.

Prevention of Corrosion

• Methods include painting, oiling, greasing, galvanising, chrome plating, anodising, or making alloys.
• Galvanisation: Coating with zinc.

Alloys

• Alloys are homogeneous mixtures of two or more metals or a metal and a non-metal.
• Examples: brass (copper and zinc), bronze (copper and tin), solder (lead and tin).