Comprehensive Study Guide on Metals, Non-Metals, and Metallurgy

Physical Properties of Metals

Metals exhibit distinct physical characteristics that define their behavior and appearance. Most metals exist in a solid state at room temperature, with the notable exceptions of mercury and gallium, which are liquids under these conditions. A defining visual trait is lustre, which refers to the property allowed by metals to reflect light from their polished surfaces, as seen in examples like gold and silver. Mechanically, metals are recognized for their malleability, allowing them to be hammered or formed into thin sheets called foils. Zinc, however, is an exception to this as it is fragile. Furthermore, metals possess ductility, meaning they can be drawn into wires, though zinc remains an exception here as well due to its fragility.

In terms of hardness, almost all metals are hard, except for sodium and potassium, which are soft enough to be sliced with a knife. Their ability to conduct heat and electricity is attributed to the presence of free electrons. Silver and copper are identified as the best conductors, whereas lead is the least efficient. Iron, bismuth, and mercury are also categorized as poor conductors. Metals generally have high densities and substantial weight, with iridium and osmium having the highest densities and lithium having the lowest. Their melting and boiling points are typically high; tungsten has the highest melting point, while silver has the lowest boiling point. Sodium and potassium are exceptions, possessing low melting values. Metals also form alloys, which are homogeneous combinations such as brass, a mixture of copper and zinc. Additionally, metals are sonorous, meaning they produce a specific sound when struck by a solid object.

Physical Properties of Non-Metals

Non-metals exhibit properties that are largely the opposite of metals. At ambient temperature, most non-metals exist as either solids or gases. Examples of solid non-metals include iodine, carbon, and sulphur, while oxygen serves as a common gaseous example. They generally lack a metallic sheen and do not reflect light; the primary exceptions are iodine and carbon in its diamond form. Unlike metals, non-metals are extremely fragile and possess neither ductility nor malleability, meaning they cannot be drawn into wires or hammered into sheets. Diamond, an allotrope of carbon, is the exception to this fragility as it is the hardest substance in the world.

Non-metals are generally poor conductors of heat and electricity. Notable exceptions include graphite, which conducts heat, and both graphite and gas carbon, which conduct electricity. Chemically, non-metals have an electronegative character, meaning they have a proclivity for gaining or sharing electrons with neighboring atoms. They are reactive, producing acidic or neutral oxides upon contact with oxygen. Compared to metals, non-metals have low melting and boiling points.

Comparison of Physical Properties between Metals and Non-Metals

The fundamental differences between metals and non-metals can be summarized across several physical dimensions. Metals are efficient conductors of heat and electricity, whereas non-metals are poor conductors. Metals are malleable and ductile, allowing for the creation of sheets and wires, while non-metals lack these properties. Metals are sonorous, but non-metals are not. From a structural standpoint, metals possess high tensile strength due to the high attraction between their molecules, while non-metals have low tensile strength resulting from lower molecular attraction. Finally, metals are characterized by high density, whereas non-metals generally have low density.

Chemical Properties of Metals and Reaction with Oxygen

Almost all metals react with oxygen to produce metal oxides, which typically exhibit basic characteristics. The general reaction is represented as:

$\text{Metal} + \text{Oxygen} \rightarrow \text{Metal oxide (basic)}$

Reactivity varies greatly across different metals. Sodium ($Na$) and potassium ($K$) are the most reactive; they react vigorously with the oxygen in the air at room temperature and can catch fire. To prevent this, they are stored immersed in kerosene oil. The reaction for potassium is:

$4K_{(s)} + O_{2(g)} \rightarrow 2K_2O_{(s)}$

Magnesium ($Mg$) does not react with oxygen at room temperature but will burn in air upon heating, producing an intense light and heat to form magnesium oxide. The reaction is:

$2Mg_{(s)} + O_{2(g)} \rightarrow 2MgO_{(s)}$

Magnesium burns specifically with a white dazzling light. In contrast, metals like silver, platinum, and gold do not burn or react with air at all.

Reaction of Metals with Water

Metals react with water to produce a metal oxide and hydrogen gas. If the resulting metal oxide is soluble in water, it dissolves further to form a metal hydroxide. The reaction sequence is:

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

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

Reactivity with water varies. Sodium and potassium react vigorously even with cold water, evolving hydrogen that immediately catches fire due to the heat produced. For example:

$2K_{(s)} + 2H_2O_{(l)} \rightarrow 2KOH_{(aq)} + H_{2(g)} + \text{heat energy}$

$2Na_{(s)} + 2H_2O_{(l)} \rightarrow 2NaOH_{(aq)} + H_{2(g)} + \text{heat energy}$

Other metals like aluminium ($Al$), zinc ($Zn$), and iron ($Fe$) do not react with cold or hot water but do react with steam to form oxides and hydrogen:

$2Al_{(s)} + 3H_2O_{(g)} \rightarrow Al_2O_{3(s)} + 3H_{2(g)}$

$3Fe_{(s)} + 4H_2O_{(g)} \rightarrow Fe_3O_{4(s)} + 4H_{2(g)}$

Metals such as lead, copper, silver, and gold do not react with water or steam in any form.

Reaction of Metals with Acids

Metals generally react with dilute acids to produce a salt and hydrogen gas:

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

When reacting with dilute hydrochloric acid ($HCl$), metals form metal chlorides. For example:

$Mg + 2HCl \rightarrow MgCl_2 + H_2$

When reacting with sulphuric acid ($H_2SO_4$), they form metal sulphates:

$Fe + H_2SO_4 \rightarrow FeSO_4 + H_2$

Nitric acid ($HNO_3$) usually does not result in the evolution of hydrogen because it is a strong oxidising agent that oxidises the produced hydrogen into water while itself reducing into nitrogen oxides. However, magnesium ($Mg$) and manganese ($Mn$) are exceptions and will react with dilute nitric acid to evolve hydrogen gas:

$Mg + 2HNO_3 \rightarrow Mg(NO_3)_2 + H_2$

$Mn + 6HNO_3 \rightarrow Mn(NO_3)_2 + H_2$

Reactivity Series and Displacement Reactions

The reactivity series is the arrangement of metals in order of their decreasing chemical reactivity. This hierarchy determines how metals interact with the salt solutions of other metals. A more reactive metal will displace a less reactive metal from its salt solution. A common example is the placement of an iron nail in a copper sulphate solution. Iron is more reactive than copper, so the blue color of the $CuSO_4$ solution fades as iron replaces the copper, forming light green iron sulphate ($FeSO_4$) and reddish-brown copper metal:

$CuSO_{4(aq)} + Fe_{(s)} \rightarrow FeSO_{4(aq)} + Cu_{(s)}$

Reaction of Metals with Chlorine

Metals react with chlorine to form ionic metal chlorides. Sodium reacts readily to form sodium chloride:

$2Na_{(s)} + Cl_{2(g)} \rightarrow 2NaCl_{(s)}$

Calcium reacts vigorously with chlorine to form calcium chloride:

$Ca_{(s)} + Cl_{2(g)} \rightarrow CaCl_{2(s)}$

Properties of Ionic Compounds

Ionic compounds are formed through the attraction of oppositely charged ions. They exist as hard solids due to the strong electrostatic forces between positive and negative ions, though they are brittle and tend to shatter under pressure. They possess high melting and boiling points because significant energy is required to overcome these strong intermolecular attractions. While they are soluble in water, they remain insoluble in organic solvents like kerosene or petrol.

In terms of electrical conductivity, ionic compounds do not conduct electricity in their solid state because the ions are fixed in place by strong electrostatic forces. However, they are conductive when fused (molten) or in an aqueous state because the ions become mobile and can carry a current.

Metallurgy: Definitions and Extraction Steps

Metallurgy involves several key terms and processes used to obtain pure metals from the earth. Minerals are naturally occurring compounds of metals found with impurities. Ores are specific minerals from which metals can be extracted profitably and conveniently. Gangue refers to the earthly impurities, such as mud and silica, that are associated with the ore. Metallurgy itself is the comprehensive process of extracting pure metals from these ores.

The extraction process depends on the reactivity of the metal. For less reactive metals (those at the bottom of the reactivity series), extraction is simpler. For example, Mercury is extracted from its ore, Cinnabar ($HgS$), by heating it in air to form mercuric oxide ($HgO$), which is then further reduced by heat to produce liquid mercury:

$2HgS_{(s)} + 3O_{2(g)} \rightarrow 2HgO_{(s)} + 2SO_{2(g)}$

$2HgO_{(s)} \rightarrow 2Hg_{(s)} + O_{2(g)}$

Extraction of Moderately Reactive Metals

Moderately reactive metals in the middle of the reactivity series are typically extracted by reducing their oxides using agents like carbon, aluminium, sodium, or calcium. It is easier to reduce metal oxides than carbonates or sulphides, so ores are first converted into oxides using two primary methods:

1. Calcination: Carbonate ores are heated strongly in the absence of air (or very limited air) to remove carbonates and moisture. Example: $ZnCO_{3(s)} \rightarrow ZnO_{(s)} + CO_{2(g)}$.
2. Roasting: Sulphide ores are heated strongly in the presence of air to convert them into oxides. Example: $2ZnS_{(s)} + 3O_{2(g)} \rightarrow 2ZnO_{(s)} + 2SO_{2(g)}$.

Following these processes, the resulting zinc oxide ($ZnO$) is reduced by heating it with carbon to produce zinc metal. In the case of iron, aluminium can be used to reduce iron oxide in a process called the Thermite reaction. This reaction generates significant heat, producing molten iron used for joining railway tracks or repairing cracked machinery:

$Fe_2O_{3(s)} + 2Al_{(s)} \rightarrow 2Fe_{(l)} + Al_2O_{3(s)} + \text{Heat}$

Extraction of Highly Reactive Metals through Electrolytic Reduction

Highly reactive metals at the top of the reactivity series have a high affinity for oxygen and cannot be reduced by carbon. These metals are obtained through electrolytic reduction of their molten salts. For instance, sodium is extracted by the electrolysis of molten sodium chloride ($NaCl$):

$2NaCl_{(l)} \rightarrow 2Na_{(s)} + Cl_{2(g)}$

During this process, sodium ions are reduced at the cathode, and chloride ions are oxidized at the anode:

At Cathode: $Na^+ + e^- \rightarrow Na$

At Anode: $2Cl^- \rightarrow Cl_2 + 2e^-$

Refining of Metals and the Electrolysis Process

The most common method for purifying impure metals is electrolytic refining. Metals such as copper, zinc, tin, lead, chromium, nickel, silver, and gold are refined this way. In electrorefining, a block of impure metal acts as the anode, and a thin sheet of pure metal acts as the cathode. The electrolytic cell contains an aqueous solution of the metal's salt. When current flows, the metal dissolves from the anode ($M \rightarrow M^{+n} + ne^-$) and deposits onto the cathode ($M^{+n} + ne^- \rightarrow M$).

In the refining of copper, an acidified copper sulphate solution is used as the electrolyte. Pure copper is deposited on the cathode from the impure copper anode. This method is also suitable for refining volatile metals with lower boiling points than their impurities, such as mercury and zinc. An electrolyte is a substance that carries current in solution or molten form and is decomposed by that current.

Corrosion and Methods of Prevention

Corrosion occurs when metals react with atmospheric components, forming oxides, sulphides, or carbonates that "eat away" the metal over time. Rusting of iron is the most prevalent form, occurring when iron reacts with moisture and oxygen to form hydrated ferric oxide ($Fe_2O_3 \cdot xH_2O$), a brown-red substance known as rust. Necessary conditions for roasting are the presence of air (oxygen) and water (moisture).

Prevention methods include:

1. Galvanising: Coating iron or steel with a thin layer of zinc (e.g., pins, nails).
2. Tinning: Coating other metals with tin.
3. Electroplating: Using electrolysis to coat one metal with another (e.g., silver-plated spoons).
4. Alloying: Creating a homogeneous mixture of metals or metals and non-metals. Common alloys include brass ($Cu$ and $Zn$), bronze ($Cu$ and $Sn$), and stainless steel (iron, nickel, chromium, and carbon). Alloys generally do not corrode easily.

Questions & Discussion

Multiple Choice Questions:

1. Aluminium is used for making cooking utensils due to its good thermal conductivity and high melting point. (Correct Choice: d)
2. The most abundant metal in the earth's crust is Aluminium. (Correct Choice: b)
3. The poorest conductor of heat among metals is Lead. (Correct Choice: a)
4. The property used for making bells and stringed instruments is Sonorousness. (Correct Choice: a)
5. The reaction $Al_2O_3 + 2NaOH \rightarrow \text{?} + H_2O$ produces $NaAlO_2$. (Correct Choice: b/c)
6. Order of reactivity (ascending): $Sodium > Magnesium > Zinc > Iron$. (Correct Choice: c)
7. Displacement reactions occur between $AgNO_3$ solution and Copper metal. (Correct Choice: b)
8. Non-metals form covalent chlorides because they can share electrons with chlorine. (Correct Choice: c/b)
9. Prolonged reaction of iron with steam yields $Fe_3O_4$. (Correct Choice: c/d)
10. compounds that are not ionic include $HCl$ and $CCl_4$. (Correct Choice: b)

Very Short Questions:

• The most abundant metal in the earth's crust is Aluminium ($Al$), comprising about $8.1\%$ by mass.
• Calcination occurs without air; Roasting occurs with excess air.
• Rust is hydrated ferric oxide: $Fe_2O_3 \cdot xH_2O$.
• Sulphide ores are enriched via the Froth Floatation process.
• Zinc evolves hydrogen more readily than iron when reacting with dilute $HCl$ because it is higher in the reactivity series.
• Brass consists of copper and zinc; Bronze consists of copper and tin.
• German silver contains copper, zinc, and nickel; it contains no actual silver.
• Main ore of iron: Haematite ($Fe_2CO_3$); Main ore of aluminium: Bauxite ($Al_2O_3 \cdot 2H_2O$).
• Graphite is a non-metal that conducts electricity.
• Neutral oxides include carbon monoxide ($CO$) and nitrous oxide ($N_2O$).

Short and Long Questions Summary:

Aluminium is highly demanded because it is a good conductor, is not attacked by water, and is a powerful reducing agent. Duralumin is an alloy of $Al$ (93%), $Cu$ (4%), $Mg$ (0.5%), and $Mn$ (0.5%). Titanium is a strategic metal because it is vital for war equipment, being lightweight, stronger than other metals, and corrosion-resistant.

Displacement reactions explain why an iron knife in copper sulphate turns the solution light green (forming ferrous sulphate) and why copper in silver nitrate deposits silver on the plate. If a bronze medal gains a green layer, it is due to the formation of basic copper carbonate ($Cu(OH)_2 \cdot CuCO_3$).

To demonstrate that ionic compounds conduct electricity only in solution, one can dip graphite rods into solid sodium chloride (bulb won't glow) and then add water (bulb glows). In solids, ions are immobile; in solution, $Na^+$ and $Cl^-$ ions move freely to carry current.

Case Study Discussion:

Metals like Titanium, Zirconium, and Manganese are essential for defense and are called strategic metals. Copper, silver, and gold are known as coinage metals. While metals are generally sonorous and conductive, some like potassium readily form cations. Conversely, non-metals like iodine form anions. Ionic compounds are generally soluble in water but not in organic solvents like petrol, and they conduct electricity only in molten or aqueous states.