Metals and Non-metals-QnA

Malleable: A metal is malleable if it can be hammered or pressed into thin sheets without cracking. This property allows metals like aluminium and gold to be shaped into foils and sheets.

Ductile: A metal is ductile if it can be drawn into thin wires without breaking. Copper is a common example — its high ductility makes it ideal for electrical wiring.

Reason: Sodium is kept immersed in kerosene oil because it reacts extremely fast with the moisture and oxygen present in air.

If left exposed, it can catch fire or even explode due to its vigorous reaction. Kerosene acts as a protective layer and prevents sodium from coming in contact with air or water.

(i) Reaction of iron with steam:

When iron is heated with steam, it forms magnetite (iron(II,III) oxide) and hydrogen gas.

Equation: \(\mathrm{3Fe + 4H_2O \xrightarrow{\text{Steam}} Fe_3O_4 + 4H_2\uparrow}\)

(ii) Reaction of calcium with water:

Calcium reacts slowly with cold water to produce calcium hydroxide and hydrogen gas.

Equation: \(\mathrm{Ca + 2H_2O \rightarrow Ca(OH)_2 + H_2\uparrow}\)

(iii) Reaction of potassium with water:

Potassium reacts very vigorously with water, producing potassium hydroxide and releasing hydrogen gas with heat.

Equation: \(\mathrm{2K + 2H_2O \rightarrow 2KOH + H_2\uparrow}\)

Gas Produced: Hydrogen gas (H2) is released when a reactive metal reacts with dilute hydrochloric acid.

Chemical Reaction (Iron with dilute \(\mathrm{H_2SO_4}\)):

\[ \mathrm{Fe} + \mathrm{H_2SO_4} \;\longrightarrow\; \mathrm{FeSO_4} + \mathrm{H_2}\uparrow \]

This happens because iron displaces hydrogen from the acid, forming iron(II) sulfate and releasing hydrogen gas as bubbles.

Observation:

When zinc is added to an iron(II) sulphate solution, the green colour of the solution gradually fades. A greyish deposit of iron forms on the surface of the zinc metal. This happens because zinc is more reactive than iron and displaces it from the solution.

Chemical Reaction:

\[ \mathrm{Zn} + \mathrm{FeSO_4} \;\longrightarrow\; \mathrm{ZnSO_4} + \mathrm{Fe} \]

Zinc replaces iron from its salt due to its higher reactivity, forming zinc sulphate while iron separates as solid metal.

Electron-dot structures:

1. Sodium (Na):

Sodium has one electron in its outermost shell. So, its electron-dot symbol shows a single dot around Na.

\[ \mathrm{Na}\,\cdot \]

2. Oxygen (O):

Oxygen has six valence electrons. These are shown as six dots arranged in pairs around the symbol O.

\[ \begin{array}{c} \cdot\quad \mathrm{O} \quad \cdot \\ \cdot\qquad \quad \cdot \end{array} \]

3. Magnesium (Mg):

Magnesium has two electrons in its valence shell. Therefore, its electron-dot structure shows two dots around Mg.

\[ \mathrm{Mg}\, \cdot\;\cdot \]

These structures represent only the valence electrons responsible for chemical bonding.

Formation of Sodium Oxide (Na₂O):

Sodium has one valence electron. Each sodium atom loses one electron to form a sodium ion.

\[ \mathrm{Na} \;\longrightarrow\; \mathrm{Na^+} + e^- \]

Oxygen requires two electrons to complete its octet. Therefore, two sodium atoms each donate one electron to the oxygen atom.

\[ 2\mathrm{Na} + \mathrm{O} \;\longrightarrow\; \mathrm{Na_2O} \]

Formation of Magnesium Oxide (MgO):

Magnesium has two valence electrons, which it loses to form a magnesium ion with a +2 charge.

\[ \mathrm{Mg} \;\longrightarrow\; \mathrm{Mg^{2+}} + 2e^- \]

Oxygen gains these two electrons to form an oxide ion with a –2 charge.

\[ \mathrm{O} + 2e^- \;\longrightarrow\; \mathrm{O^{2-}} \]

These oppositely charged ions combine to form magnesium oxide.

\[ \mathrm{Mg} + \mathrm{O} \;\longrightarrow\; \mathrm{MgO} \]

Thus, both Na₂O and MgO are formed by complete transfer of electrons from metal atoms to oxygen.

The ions present in the given compounds are as follows:

(a) Sodium chloride (NaCl):

\[ \text{Ions: } \mathrm{Na^+} \text{ and } \mathrm{Cl^-} \]

(b) Magnesium chloride (MgCl₂):

\[ \text{Ions: } \mathrm{Mg^{2+}} \text{ and } 2\mathrm{Cl^-} \]

(c) Sodium oxide (Na₂O):

\[ \text{Ions: } 2\mathrm{Na^+} \text{ and } \mathrm{O^{2-}} \]

(d) Magnesium oxide (MgO):

\[ \text{Ions: } \mathrm{Mg^{2+}} \text{ and } \mathrm{O^{2-}} \]

Each compound contains positively charged metal ions and negatively charged non-metal ions formed by electron transfer.

Ionic compounds have high melting points because the ions in these solids are held together by very strong electrostatic forces, known as ionic bonds.

Each positive ion (cation) is strongly attracted to the surrounding negative ions (anions), creating a rigid and stable lattice structure.

To melt such a compound, a large amount of heat energy is needed to overcome these strong forces.

For example, in an ionic compound like:

\[ \mathrm{Na^+ \;\; Cl^-} \]

The attraction between \(\mathrm{Na^+}\) and \(\mathrm{Cl^-}\) is very strong, which results in a high melting point.

Therefore, the strong ionic bonds in their lattice structure are the main reason for the high melting points of ionic compounds.

(i) Mineral: A mineral is a naturally occurring, inorganic solid with a definite (or at least limited) chemical composition and an ordered crystalline structure.
Example: Bauxite is a mineral mainly containing hydrated aluminium oxide, written as \(\mathrm{Al_2O_3\cdot xH_2O}\).

(ii) Ore: An ore is a mineral (or rock) that contains a metal or a useful substance in concentrations high enough to make extraction economically viable. In short, an ore is a mineral that is worth mining.
Example: Haematite, an iron ore, is represented as \(\mathrm{Fe_2O_3}\).

(iii) Gangue: Gangue refers to the unwanted, non-metallic impurities or the worthless rock surrounding or mixed with an ore. These are removed during the concentration process (e.g., by hand-picking, flotation, or washing) before extraction of the metal.

Two common metals found in nature in the free (native) state are gold and silver.

In chemical notation these are written as \(\mathrm{Au}\) (gold) and \(\mathrm{Ag}\) (silver). These metals are chemically unreactive, which is why they often occur in their elemental form in the earth's crust.

Answer: The chemical process used to obtain a metal from its oxide is reduction. In reduction, oxygen is removed from the metal oxide to yield the free metal.

Common methods:

• Reduction by carbon (smelting): A carbon-based reducing agent (coke) removes oxygen from many metal oxides. Example for iron extraction:

  \[ \mathrm{Fe_2O_3} + 3\mathrm{C} \;\longrightarrow\; 2\mathrm{Fe} + 3\mathrm{CO} \]

• Reduction by carbon monoxide (blast furnace):

  \[ \mathrm{Fe_2O_3} + 3\mathrm{CO} \;\longrightarrow\; 2\mathrm{Fe} + 3\mathrm{CO_2} \]

• Electrolytic reduction (for very reactive metals): Highly reactive metal oxides (like \(\mathrm{Al_2O_3}\)) are reduced by electrolysis. Example (conceptual):

  \[ \mathrm{Al_2O_3} \xrightarrow{\text{electrolysis}} 2\mathrm{Al} + \tfrac{3}{2}\mathrm{O_2} \]

Summary: The general principle is removal of oxygen (reduction); the specific method—smelting with carbon or electrolysis—depends on the metal's reactivity and economic feasibility.

Some metals that do not corrode easily are listed below, with brief reasons:

• Gold (\(\mathrm{Au}\)) — Extremely unreactive; it does not combine with oxygen or most chemicals, so it remains metallic in appearance for a very long time.
• Platinum (\(\mathrm{Pt}\)) — Another noble metal; highly resistant to oxidation and chemical attack, used where corrosion resistance is critical.
• Silver (\(\mathrm{Ag}\)) — Relatively low reactivity; it tarnishes (forms \(\mathrm{Ag_2S}\)) on exposure to sulphur compounds but does not undergo widespread corrosive degradation like iron.
• Aluminium (\(\mathrm{Al}\)) — Forms a thin, tightly adherent oxide layer (\(\mathrm{Al_2O_3}\)) on exposure to air that protects the underlying metal from further corrosion (passivation).
• Stainless steel (alloy containing Fe–Cr–Ni) — Chromium in the alloy forms a protective chromium oxide film that prevents rusting, so stainless steels are much less prone to corrosion than plain iron.\

Summary: Noble metals like \(\mathrm{Au}\) and \(\mathrm{Pt}\) are naturally non-corroding. Other metals (e.g., \(\mathrm{Al}\), stainless steel) resist corrosion because they form stable, protective oxide films or are alloyed to enhance protection.

Alloys are homogeneous mixtures of two or more elements, where at least one component is a metal. They are created to combine desirable properties of the constituent elements — for example, greater strength, corrosion resistance, lower melting point, or improved workability — than those of the pure metals.

Key points:

• Alloys are not simple chemical compounds but solid solutions or intermetallic mixtures with atoms of different elements occupying positions in a metal lattice.
• The properties of an alloy depend on its composition and the proportions of its constituents.

Common examples:

• Brass: primarily copper and zinc — often written conceptually as Cu–Zn. Used for decorative items and musical instruments.
• Bronze: copper and tin — Cu–Sn. Known for hardness and wear resistance (e.g., statues, bearings).
• Steel: iron with carbon and other elements (Cr, Ni, etc.) — often represented as Fe–C or more specifically stainless steel as Fe–Cr–Ni alloy. Used in construction, tools, and appliances.
• Solder: tin–lead or tin–silver mixtures — used for joining electrical components.

Example (conceptual notation with MathJax):

\[\scriptsize \text{Brass: } \\\mathrm{Cu} + \mathrm{Zn} \;\longrightarrow\; \text{Brass (Cu–Zn alloy)} \]
\[\scriptsize \text{Steel: } \\\mathrm{Fe} + \mathrm{C} \;(+\;\text{Cr, Ni, etc.}) \\\longrightarrow\; \text{Steel (Fe–C alloy)} \]

Why alloys are useful: By varying the type and amount of each component, manufacturers tailor mechanical, thermal, and chemical properties to suit specific applications — for example, adding chromium to iron produces stainless steel that resists corrosion.