Comprehensive Notes on Alkali Metals, Reactions, and Metal Extraction

General Properties and Distinctions of Metals and Non-Metals

Metals and non-metals exhibit vastly different physical and chemical characteristics. Most metals are found in the solid state, with the notable exception of mercury, which is liquid at room temperature. Physically, metals are recognized for being hard, strong, and shiny. Their surfaces are typically difficult to scratch and they cannot be broken easily due to their inherent strength. Two critical mechanical properties of metals include malleability, which is the ability to be shaped or flattened by a hammer without breaking, and ductility, which allows them to be drawn into thin wires. Furthermore, metals are sonorous, meaning they produce a ringing sound when struck. Thermally and electrically, metals are excellent conductors and generally possess high melting points, high boiling points, and high density.

In contrast, non-metals are primarily gases or liquids. They are generally soft, weak, and dull in appearance. Unlike metals, they are brittle and tend to break easily when subjected to impact. When struck, they produce a dull sound rather than a ringing one. Non-metals are characterized by low melting points, low boiling points, and low density. With few exceptions, they are non-conductors of both heat and electricity.

Comparative Characteristics of Alkali and Transition Metals

Alkali metals, located in Group I of the periodic table, possess distinct properties compared to transition metals. Alkali metals are remarkably soft and can be cut easily with a knife. While they are shiny, this luster is only visible when they are freshly cut, as they tarnish quickly. Compared to transition metals, alkali metals have lower melting points, lower boiling points, and lower densities. Chemically, they are highly reactive, and their reactivity increases as you move down the group from Lithium ($Li$) to Sodium ($Na$) to Potassium ($K$). They typically exhibit only one oxidation state, or valency, and form white compounds that result in colorless solutions when dissolved in water.

Transition metals are markedly different, being hard, strong, and shiny with high melting points, boiling points, and densities. They are characterized by having more than one oxidation state; for example, Iron can exist as $Fe^{2+}$ or $Fe^{3+}$. These metals often form colored compounds and produce colored solutions. A significant application of transition metals and their compounds is their ability to act as catalysts in chemical reactions.

The Reactivity Series and Metal Extraction Methods

The reactivity series ranks metals in order of their chemical reactivity, decreasing from the most reactive at the top to the least reactive at the bottom. The series includes Potassium ($K$), Sodium ($Na$), Calcium ($Ca$), Magnesium ($Mg$), Aluminium ($Al$), Carbon ($C$), Zinc ($Zn$), Iron ($Fe$), Tin ($Sn$), Lead ($Pb$), Hydrogen ($H$), Copper ($Cu$), Silver ($Ag$), Gold ($Au$), and Platinum ($Pt$). Metals at the top of the series, from Potassium to Aluminium, are highly reactive and react vigorously or fast with water and acids. Specifically, Potassium and Sodium are never added to acids because the reaction is expected to be too vigorous and dangerous. These highly reactive metals are extracted from their ores using electrolysis.

Metals positioned in the middle of the series, such as Zinc, Iron, Tin, and Lead, are less reactive with water and acids. For instance, Lead exhibits a very slow reaction with steam. These metals are typically extracted by heating the metal oxide with Carbon. Metals at the very bottom, including Silver, Gold, and Platinum, are unreactive and do not react with water or acids. They are often found in their native state.

Displacement Reactions and Observational Chemistry

A displacement reaction is a competitive reaction where a more reactive metal displaces a less reactive metal from its compound. An example of this is the reaction between Zinc and Copper(II) sulfate. Because Zinc is more reactive than Copper, it displaces the Copper ions. The chemical equation for this reaction is $Zn + CuSO_4 \rightarrow ZnSO_4 + Cu$. In this process, the shiny grey Zinc atoms react with the blue Copper(II) sulfate solution. As the reaction proceeds, the blue color of the solution fades or becomes paler, eventually becoming a colorless solution of Zinc sulfate ($ZnSO_4$). Simultaneously, the Zinc disappears and is replaced by a pink-brown solid, which is the Copper metal being formed. The fading of the blue color occurs because the concentration of $Cu^{2+}$ ions, which are responsible for the blue color, decreases as they are converted into Copper atoms.

While most metals are shiny grey, there are exceptions such as Copper, which is pink-brown, and Gold, which is yellow. Another notable reaction is the burning of Magnesium in air or oxygen: $2Mg + O_2 \rightarrow 2MgO$. This reaction produces a bright white flame and results in the formation of a white solid called Magnesium oxide ($MgO$).

Electrochemical Principles: Redox and Electron Transfer

A redox reaction is a process where reduction and oxidation take place at the same time. These processes can be defined in terms of oxygen transfer or electron transfer. Regarding oxygen, oxidation is the gain of oxygen, while reduction is the loss of oxygen. For example, in the reaction $C + 2ZnO \rightarrow CO_2 + 2Zn$, Carbon is oxidized because it gains oxygen to become $CO_2$, and Zinc oxide ($ZnO$) is reduced because it loses oxygen to become Zinc metal.

In terms of electrons, the mnemonic "OIL RIG" is used: Oxidation Is Losing electrons, and Reduction Is Gaining electrons. In the displacement reaction between Zinc and Copper(II) sulfate, Zinc atoms ($Zn^0$) lose electrons to become Zinc ions ($Zn^{2+}$), meaning Zinc undergoes oxidation. Consequently, the oxidation number of Zinc increases from $0$ to $2+$. Zinc acts as the reducing agent because it provides electrons to another species. Conversely, Copper ions ($Cu^{2+}$) gain electrons to become Copper atoms ($Cu^0$), meaning they undergo reduction. The oxidation number of Copper decreases from $2+$ to $0$. Copper ions act as the oxidizing agent because they accept electrons from the Zinc.

Atomic Structure and the Properties of Metallic Bonds

The physical properties of metals can be explained by their internal structure, which consists of a lattice of positive metal ions surrounded by a "sea" of delocalised electrons. Metals are malleable and ductile because they have regular layers of positive ions that can slide past each other when force is applied, without breaking the metallic bond. The high melting and boiling points of metals are due to their giant metallic structure. There are strong electrostatic forces of attraction between the positive ions and the delocalised electrons, which require a high amount of energy to overcome.

An electric current is defined as the flow of charged particles, which can be either electrons or ions. A substance is considered a conductor if it contains charged particles that are free to move. Metals are good conductors because their delocalised electrons are free to flow through the structure. Graphite is a unique non-metal conductor because each Carbon atom is bonded to three others, leaving one delocalised electron per atom that is free to flow. Ionic compounds do not conduct electricity in solid form because their ions are fixed in place; however, they become conductors when molten or in an aqueous solution because the ions are then free to move. Covalent compounds generally do not conduct electricity because they lack both ions and delocalised electrons.

Practical Applications and Chemical Resistance of Metals

Specific metals are chosen for various applications based on their properties. Copper is utilized for electrical wiring because it is a good conductor of electricity and is highly ductile. It is also used for cooking pans because it conducts heat well and has a high melting point. Aluminium is used in electrical wires due to its conductivity and ductility, but its high melting point and non-toxic nature also make it suitable for drink containers. Aluminium is particularly valued for aircraft bodies because it is strong yet has a low density.

Although Aluminium is high in the reactivity series, it unexpectedly resists corrosion. This is because it quickly forms a thin, tough outside layer of Aluminium oxide ($Al_2O_3$), which is highly unreactive and protects the underlying metal from further reaction with air or water.

The Nature of Alloys and the Process of Corrosion

An alloy is a mixture of a metal with another element, such as Brass, which is a mixture of Copper and Zinc. Alloys are typically stronger and harder than pure metals. In a pure metal, atoms are of the same size and arranged in regular layers that slide easily. In an alloy, the introduction of different-sized particles disrupts this regular layer structure, making it much more difficult for the layers to slide past each other. Alloys are also often more resistant to corrosion and may be shinier.

Corrosion occurs when a metal reacts with air and water, causing it to wear out and form a less useful metal oxide. "Rusting" is a term specifically reserved for the corrosion of iron. The chemical equation for rusting is: $Iron + Water + Oxygen \rightarrow \text{Hydrated iron(III) oxide}$. For rusting to occur, both air (oxygen) and water must be present. Experiments show that iron in boiled water (no air) or iron in dry air (using a drying agent like silica gel to remove water vapour) will not rust.

Methods for Rust Protection and Prevention

Rust protection can be achieved through several methods. The barrier method involves providing a physical coating to prevent air and water from reaching the metal surface. These coatings include paint, oil, or plastic. However, if the barrier is scratched, the protection fails. A more robust method is sacrificial protection, which involves using a more reactive metal that reacts with air and water instead of the iron. Galvanising is a specific type of sacrificial protection where iron is covered with Zinc. Unlike simple barriers, sacrificial protection continues to protect the iron even if the Zinc coating is scratched, because the more reactive Zinc will preferentially oxidize.

Industrial Extraction of Iron and Aluminium

Metal ores are rocks rich in a specific metal compound. Iron is extracted from Hematite (Iron(III) oxide, $Fe_2O_3$) in a blast furnace. The furnace is fed with Hematite, Coke (a source of Carbon to reduce the ore), and Limestone ($CaCO_3$) to remove impurities. Hot air is blasted in to facilitate combustion. The chemical reactions include:

1. Combustion: $C + O_2 \rightarrow CO_2$, which releases heat.
2. Production of reducing agent: $C + CO_2 \rightarrow 2CO$, where Carbon monoxide acts as a powerful reducing agent.
3. Reduction of Ore: $Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2$. The Limestone undergoes thermal decomposition ($CaCO_3 \rightarrow CaO + CO_2$), and the resulting Calcium oxide reacts with sand (silica) impurities to form slag ($CaO + SiO_2 \rightarrow CaSiO_3$). Molten slag floats on top of the iron and is used in road building. The extracted iron, known as pig iron or cast iron, is used for items like manhole covers or converted into steel.

Aluminium is extracted from Bauxite ($Al_2O_3$) via electrolysis because Aluminium is more reactive than Carbon. A major challenge is that Bauxite has a very high melting point (over $2000\,^{\circ}\text{C}$). To solve this, Cryolite is added, which lowers the melting point of the Bauxite and improves electrical conductivity, making the process more energy-efficient and cost-effective.