Chemistry Chapter Metal

Part 1: General Physical Properties of Metals and Non-Metals

Metals and non-metals can be distinguished through several distinct physical properties:

Thermal Conductivity: Metals serve as highly efficient thermal conductors. This conductivity is attributed to the presence of delocalised electrons, which are capable of moving throughout the structure to transfer thermal energy between locations. In contrast, non-metals are generally poor conductors of heat.
Electrical Conductivity: Metals are excellent electrical conductors because their delocalised electrons are free to move and transport an electrical charge. Non-metals typically act as insulators, exhibiting poor conductivity for both electricity and heat.
Malleability and Ductility: Metals are characterised as malleable (able to be reshaped or bent without breaking) and ductile (able to be drawn out into wires). These traits arise because the positive metal ions are organized in uniform, regular rows, allowing them to slide over one another easily when force is applied.
Melting and Boiling Points: Unlike non-metals, which generally have low melting and boiling points, metals possess high transition temperatures. Metals are defined by giant structures held together by strong metallic bonding. This involves a potent electrostatic attraction between positive metal ions and negative delocalised electrons, requiring substantial energy to overcome.

General Chemical Properties and Reactivity of Metals

Metals exhibit specific patterns of chemical behavior when reacting with acids, water, and oxygen:

Reactions with Dilute Acids

When a metal reacts with a dilute acid, it produces a salt and hydrogen gas. This is classified as a displacement reaction. If a metal resides above hydrogen in the reactivity series, its ions will displace the hydrogen ions within the acid.

General Equation: $\text{Metal} + \text{Acid} \rightarrow \text{Salt} + \text{Hydrogen}$
Example: $\text{Sodium} + \text{Hydrochloric acid} \rightarrow \text{Sodium chloride} + \text{Hydrogen}$
Chemical Equation: $2Na(s) + 2HCl(aq) \rightarrow 2NaCl(aq) + H_2(g)$
Metals positioned below hydrogen in the reactivity series cannot displace it and therefore do not react with dilute acids to form salts.

Reactions with Cold Water and Steam

Metals reacting with cold water ( $H_2O(l)$ ) produce a metal hydroxide and hydrogen gas. The vigor of this reaction depends on the metal's placement in the reactivity series:

Potassium ( $K$ ): Reacts very vigorously and exothermically, burning with a characteristic lilac flame.
General Equation: $\text{Metal} + \text{Water} \rightarrow \text{Metal hydroxide} + \text{Hydrogen gas}$
Example: $2K(s) + 2H_2O(l) \rightarrow 2KOH(aq) + H_2(g)$

When metals react with steam ( $H_2O(g)$ ), they produce a metal oxide and hydrogen gas. Less reactive metals, such as Magnesium, react more vigorously with steam than with liquid water.

Example: $\text{Magnesium} + \text{Water (Steam)} \rightarrow \text{Magnesium oxide} + \text{Hydrogen gas}$
Chemical Equation: $Mg(s) + H_2O(g) \rightarrow MgO(s) + H_2(g)$

Reactions with Oxygen

Metals react with oxygen to form metal oxides. The rate of the reaction is determined by the metal's reactivity.

General Equation: $\text{Metal} + \text{Oxygen} \rightarrow \text{Metal oxide}$
Example: $2Ca(s) + O_2(g) \rightarrow 2CaO(s)$

Part 2: Practical Applications of Metals

The industrial use of metals is determined by their physical properties:

Aluminium in Aircraft: Utilized due to its low density. The lightweight nature of aluminium atoms makes structures easy to maneuver.
Aluminium in Overhead Cables: Selected for its low density and high electrical conductivity. The delocalised electrons allow for efficient charge transport while the lightness aids in installation.
Aluminium in Food Containers: Chosen for its high resistance to corrosion. Aluminium naturally forms a protective aluminium oxide layer that halts further oxidation.
Copper in Electrical Wiring: Ideal due to its high electrical conductivity and ductility, allowing it to be stretched into thin wires without fracturing.

Part 3: Characteristics and Structures of Alloys

An alloy is a mixture consisting of a metal combined with other elements (metallic or non-metallic). This process is intentional, aimed at improving strength or corrosion resistance.

Common Examples

Brass: A mixture of copper ( $Cu$ ) and zinc ( $Zn$ ).
Stainless Steel: A mixture of iron ( $Fe$ ) combined with carbon ( $C$ ), chromium ( $Cr$ ), and nickel ( $Ni$ ).

Strength and Structure

Aloys are typically harder and stronger than pure metals. In a pure metal, identical atoms are arranged in neat, uniform layers that slide over each other easily, making the metal soft. In an alloy, the introduction of different-sized atoms distorts these regular layers. This disruption prevents the layers from sliding over one another, resulting in a harder and more durable material.

Corrosion Resistance: Alloys like stainless steel are used for cutlery because they resist rusting (the chemical destruction of materials by environmental oxygen and water) better than pure iron.

Part 4: The Reactivity Series

The reactivity series ranks metals by their chemical reactivity, ranging from the most reactive to the least reactive:

Potassium ( $K$ ) - Most reactive
Sodium ( $Na$ )
Calcium ( $Ca$ )
Magnesium ( $Mg$ )
Aluminium ( $Al$ )
(Carbon) ( $C$ )
Zinc ( $Zn$ )
Iron ( $Fe$ )
(Hydrogen) ( $H$ )
Copper ( $Cu$ )
Silver ( $Ag$ )
Gold ( $Au$ ) - Least reactive

Experimental Observations with Water and Acid

Potassium with Cold Water: Melts, floats, and moves rapidly. It ignites with a lilac flame and may result in a small explosion. $2K(s) + 2H_2O(l) \rightarrow 2KOH(aq) + H_2(g)$ .
Sodium with Cold Water: Melts into a ball, fizzes rapidly, and moves across the surface. The reaction is less vigorous than potassium. $2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g)$ .
Calcium with Cold Water: Sinks; the reaction is slower and less violent, with hydrogen bubbles appearing after some time. $Ca(s) + 2H_2O(g) \rightarrow Ca(OH)_2(aq) + H_2(g)$ .
Magnesium with Steam: Magnesium reacts very slightly with cold water because magnesium hydroxide is insoluble. However, it reacts with steam to produce magnesium oxide. $Mg(s) + H_2O(g) \rightarrow MgO(aq) + H_2(g)$ .
Reactions with Hydrochloric Acid ( $HCl$ ):
- Magnesium: Fast reaction.
- Zinc: Slightly fast.
- Iron: Slightly fast.
- Copper, Silver, Gold: No reaction (positioned below hydrogen).

Ion Formation and Displacement

Reactivity is defined by a metal's tendency to lose outer shell electrons to form positive ions (oxidation). In displacement reactions, a more reactive metal replaces a less reactive metal within a compound.

Example: $ZnSO_4(aq) + Ca(s) \rightarrow Zn(s) + CaSO_4(aq)$ .
Half-reactions: $Ca(s) \rightarrow Ca^{2+}(aq) + 2e^-$ (Oxidation); $Zn^{2+}(aq) + 2e^- \rightarrow Zn(s)$ (Reduction).
Aluminium's Unreactivity: Despite its high position in the series, aluminium appears unreactive because it forms a tough, unreactive layer of aluminium oxide ( $Al_2O_3$ ) that protects the metal beneath from further oxidation.

Part 5: Corrosion and Protection

Rusting is specifically the oxidation of iron and steel, requiring both water and oxygen to produce hydrated iron(III) oxide.

Prevention Methods

Barrier Methods: Painting, greasing, oiling, or coating with plastic/galvanising protects the metal by physically excluding oxygen and water.
Sacrificial Protection: Involves using a more reactive metal (e.g., Zinc) to protect a less reactive one (e.g., Iron). The more reactive metal is oxidized more easily, losing electrons and preventing the iron from being oxidized.
Galvanising: Coating iron or steel with zinc. Zinc acts as both a barrier and as sacrificial protection.

Part 6: Extraction of Metals

Metals are extracted from ores found in the Earth's crust based on their reactivity:

Native Form: Unreactive metals like gold and silver are found as pure elements.
Reduction with Carbon: Metals like zinc and iron are extracted by reacting their oxides with carbon, which displaces the metal. This occurs for metals below carbon in the reactivity series.
Electrolysis: Metals above carbon (like aluminium) require electricity to be extracted from their molten compounds.

Extraction of Iron in the Blast Furnace

Iron is extracted from hematite (containing $Fe_2O_3$ ) using coke, limestone ( $CaCO_3$ ), and hot air.

Coke Combustion: $C(s) + O_2(g) \rightarrow CO_2(g)$ (releases heat).
Formation of Carbon Monoxide: $CO_2(g) + C(s) \rightarrow 2CO(g)$ .
Reduction of Hematite: $Fe_2O_3(s) + 3CO(g) \rightarrow 2Fe(l) + 3CO_2(g)$ .
Limestone Decomposition: $CaCO_3(s) \rightarrow CaO(s) + CO_2(g)$ .
Slag Formation: $CaO(s) + SiO_2(s) \rightarrow CaSiO_3(l)$ (removes sand/silica as calcium silicate slag).

Extraction of Aluminium

Extracted from purified bauxite ( $Al_2O_3$ ) via electrolysis.

Cryolite: Aluminium oxide is dissolved in molten cryolite to lower the melting point of the electrolyte, reducing energy costs.
Electrodes: Made of graphite (carbon). The anodes must be replaced regularly because they react with the oxygen produced to form $CO_2$ .
Cathode Reaction (Reduction): $Al^{3+} + 3e^- \rightarrow Al$ .
Anode Reaction (Oxidation): $2O^{2-} - 4e^- \rightarrow O_2$ .