The Chemistry Of Selected Metals And Non-Metals

Introduction to the Classification of Elements

The periodic table is primarily composed of metals, whereas non-metals constitute a smaller portion of approximately twenty ($20$) elements. Geographically, within the table, metals are positioned on the left side, while non-metals reside on the right. Between these two extremes lies a transitional set of elements known as metalloids or semimetals, which exhibit a blend of characteristics from both categories. These elements are predominantly found in groups $3A$ and $5A$. A defining feature of metalloids is their role as semiconductors, meaning they can conduct electricity but with less efficiency than pure metals. This specific physical property has made them indispensable in the electronics industry, particularly in the creation of diodes and transistors. Understanding the distinct physical and chemical properties of metals, non-metals, and metalloids is essential for a comprehensive analysis of chemical behavior.

Defining Metals and Non-Metals through Atomic Ionization

Metals are fundamentally defined by the behavior of their atoms during ionization, specifically their tendency to lose electrons. Atoms of metallic elements possess a small number of electrons in their outermost s and p shells, which they can easily shed to achieve an electron configuration identical to the nearest noble gas. For example, sodium ($Na$) has the configuration $[Ne]\,3s^1$. When it loses its valence electron, it becomes a sodium ion ($Na^+$) with the configuration $[Ne]$. Similarly, magnesium ($Mg$) with the configuration $[Ne]\,3s^2$ loses two electrons to become $Mg^{2+}$ with the configuration $[Ne]$. This characteristic electron loss is the hallmark of metallic nature.

Conversely, non-metals are elements that undergo ionization by gaining electrons. Elements such as those in the oxygen family or the halogens acquire electrons to fill their valence shells and reach the stable configuration of the nearest noble gas atom. For instance, chlorine ($Cl$), which has the electron configuration $[Ne]\,3s^2\,3p^5$, gains one electron to form a chloride ion ($Cl^-$) with the argon configuration $[Ar]$. Sulfur ($S$), with the configuration $[Ne]\,3s^2\,3p^4$, gains two electrons to form the sulfide ion ($S^{2-}$) with the configuration $[Ar]$. Metalloids remain the intermediate group, displaying properties that bridge these two distinct domestic chemical behaviors.

Comparative Physical and Chemical Properties of Metals and Non-Metals

The physical state of metals is generally solid at room temperature, with the notable exception of mercury ($Hg$). They are characterized by high melting and boiling points, and they serve as excellent conductors of both heat and electricity, with aluminum ($Al$) serving as a prime example. Metals are physically distinguished by their malleability and ductility, and they possess a characteristic metallic lustre. Structurally, metals exist as crystal lattices held together by strong metallic bonds. Chemically, they form positive ions (cations), act as effective reducing agents, and possess low electronegativity values.

Non-metals exhibit significantly different properties. They can exist in gaseous form (such as oxygen, $O_2$), liquid form (such as bromine, $Br_2$), or solid form (such as carbon, $C$). With the exception of certain forms of carbon, non-metals generally have low melting and boiling points. They are generally poor conductors of heat and electricity, though graphite is a significant exception. Unlike metals, solid non-metals are brittle and cannot be hammered into shapes. Most lack a characteristic lustre, though carbon is again an exception. At the molecular level, non-metals (excluding diamond) exist as covalent molecules held by weak van der Waals forces. Chemically, they form negative ions (anions), act as oxidizing agents, and maintain high electronegativities.

Distribution and Periodic Trends of Elements

The boundary between metals and non-metals in the periodic table is defined by a diagonal demarcation line. This line passes through boron ($B$), silicon ($Si$), and astatine ($At$), and also encompasses germanium ($Ge$), antimony ($Sb$), and polonium ($Po$). According to standard chemical definitions, elements to the left of this divide are predominantly metallic. Along the line itself, boron and silicon are classified as non-metals, while the others are identified as metalloids. In terms of chemical trends, the metallic nature of elements increases as one moves down any specific group. Across a period, there is a clear transition point where properties shift from primarily metallic to primarily non-metallic.

Metals typically form basic oxides and exhibit positive oxidation states. While they rarely form chemical combinations among themselves, they can physically merge to form alloys. Non-metals are more versatile in their bonding, combining with themselves, other non-metals, and metals. In binary compounds with metals, non-metals show negative oxidation states. However, when bonding with other non-metals, their oxidation state can be either positive or negative, depending on the relative electronegativities of the involved elements.

Characteristics of Groups I, II, and VII

Metals in Groups I and II are designated as s-block metals. Group I metals, known as Alkali Metals, are the most powerful reducing agents, particularly those located toward the bottom of the group. This is because their single outermost s electron is extremely well-shielded from the nuclear charge, making it weakly held and easily lost. When reacting with water, these metals produce strong alkalis. Group II metals, known as Alkaline Earth Metals, produce weaker alkalis upon reaction with water. On the non-metal side, Group VII elements (the Halogens) showcase the most complete non-metallic properties. Fluorine ($F$), positioned at the top of this group, is the most powerful oxidizing agent because its nuclear charge is poorly shielded by its seven electrons, allowing it to readily capture an electron to form a fluoride ion ($F^-$).

Physical Properties of Group I Elements

For metals in Groups I and II, properties such as melting point, boiling point, and hardness decrease as the atomic number increases. These metals are good conductors but possess lower densities compared to transition metals. There is a significant difference in the hydration energy between Group IA and Group IIA ions; an alkaline earth ion typically has a hydration energy roughly five times that of an alkali metal ion in the same period due to a higher ionic charge to ionic radius ratio.

Numerical data for Group I elements reveals the following trends:

Lithium ($Li$): Configuration $[He]\,2s^1$, Electronegativity $1.0$, I.E. $520\,kJ/mol$, Melting Point $181^\circ C$, Boiling Point $1331^\circ C$, Atomic Radius $123\,pm$, Ionic Radius $68\,pm$, $E^\circ$ of $-3.03\,V$.

Sodium ($Na$): Configuration $[Ne]\,3s^1$, Electronegativity $0.9$, I.E. $513\,kJ/mol$, Melting Point $98^\circ C$, Boiling Point $890^\circ C$, Atomic Radius $157\,pm$, Ionic Radius $98\,pm$, $E^\circ$ of $-2.71\,V$.

Potassium ($K$): Configuration $[Ar]\,4s^1$, Electronegativity $0.8$, I.E. $419\,kJ/mol$, Melting Point $63^\circ C$, Boiling Point $766^\circ C$, Atomic Radius $203\,pm$, Ionic Radius $133\,pm$, $E^\circ$ of $-2.92\,V$.

Rubidium ($Rb$): Configuration $[Kr]\,5s^1$, Electronegativity $0.8$, I.E. $400\,kJ/mol$, Melting Point $39^\circ C$, Boiling Point $688^\circ C$, Atomic Radius $216\,pm$, Ionic Radius $148\,pm$, $E^\circ$ of $-2.93\,V$.

Caesium ($Cs$): Configuration $[Xe]\,6s^1$, Electronegativity $0.7$, I.E. $380\,kJ/mol$, Melting Point $29^\circ C$, Boiling Point $685^\circ C$, Atomic Radius $235\,pm$, Ionic Radius $167\,pm$, $E^\circ$ of $-3.04\,V$.

Extraction and Properties of Sodium

Due to high reactivity, sodium occurs in nature as sodium chloride in sea water. It cannot be extracted by conventional reduction or aqueous electrolysis due to its high negative reduction potential. Instead, it is produced via the Downs Process in a specialized Downs cell. The process involves the electrolysis of a molten mixture of sodium chloride and calcium chloride in a $2:3$ ratio. This mixture lowers the melting point to approximately $600^\circ C$, which is $200^\circ C$ lower than pure $NaCl$. During electrolysis, sodium is discharged at the steel cathode ring, while chlorine is released at the graphite anode. The reaction at the anode is $2Cl^- - 2e^- \rightarrow Cl_2$, and at the cathode is $Na^+ + e^- \rightarrow Na$. A large current at low voltage ensures the mixture remains molten and the sodium is effectively discharged and collected in inverted troughs.

Sodium is a soft, silvery metal with a density of $0.97\,g/cm^3$. It tarnishes rapidly in air, forming sodium oxide: $4Na(s) + O_2(g) \rightarrow 2Na_2O(s)$. This oxide reacts with atmospheric moisture to form sodium hydroxide: $Na_2O(s) + H_2O(g) \rightarrow 2NaOH(aq)$, which eventually absorbs $CO_2$ to become hydrated sodium carbonate: $2NaOH(aq) + CO_2(g) \rightarrow Na_2CO_3(s) + H_2O(l)$. If heated in excess air, sodium burns with a golden yellow flame to produce sodium peroxide: $2Na(s) + O_2(g) \rightarrow Na_2O_2(s)$. Because of this reactivity, it must be stored under paraffin oil or toluene. It reacts vigorously with cold water ($2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g)$), explosively with dilute acids ($2Na(s) + 2HCl(aq) \rightarrow 2NaCl(aq) + H_2(g)$), and forms sodamide with ammonia ($2Na(s) + 2NH_3(g) \rightarrow 2NaNH_2(s) + H_2(g)$).

Extraction and Properties of Aluminium

Aluminium is the most abundant metal in the Earth's crust, primarily sourced from Bauxite, which contains up to $60\%$ Alumina ($Al_2O_3$). Other sources include cryolite, kaolin, corundum, and mica. Extraction occurs in two stages. Stage I is the purification of Bauxite, where it is heated with concentrated $NaOH$ under pressure to form soluble sodium aluminate ($III$): $Al_2O_3(s) + 2NaOH(aq) + 3H_2O \rightarrow 2NaAl(OH)_4(aq)$. After filtering out impurities like iron ($III$) oxide, the filtrate is seeded with pure aluminium hydroxide crystals to precipitate $Al(OH)_3$, which is then heated to yield pure Alumina: $2Al(OH)_3(s) \rightarrow Al_2O_3(s) + 3H_2O(l)$.

Stage II is the electrolysis of Alumina in a graphite-lined cell. The alumina is dissolved in molten cryolite ($Na_3AlF_6$) to lower the operational temperature to $1000^\circ C$. Graphite anodes are consumed as they react with evolved oxygen to form $CO_2$. At the cathode, reduction occurs: $2Al^{3+} + 6e^- \rightarrow 2Al(s)$. At the anode, oxygen is produced: $2O^{2-} \rightarrow O_2(g) + 4e^-$. Aluminium is a silvery-white metal that is corrosion-proof due to a thin oxide layer that forms in air. It reacts with acids like $HCl$ to produce hydrogen and forms soluble aluminates when reacting with alkalis: $2Al(s) + 2NaOH(aq) + 6H_2O(l) \rightarrow 2NaAl(OH)_4(aq) + 3H_2(g)$. It is also used in the thermite process to reduce iron oxide: $2Al(s) + Fe_2O_3(s) \rightarrow Al_2O_3(s) + 2Fe(s)$.

The Chemistry of Nitrogen and Chlorine

Nitrogen ($N_2$) makes up $78\%$ of the atmosphere and is chemically unreactive due to the high bond energy of its triple covalent bond. It is industrially prepared via the fractional distillation of liquid air, boiling at $-196^\circ C$. While inert at room temperature, it reacts at high temperatures with metals to form nitrides, such as magnesium nitride ($Mg_3N_2$), which can be hydrolyzed to produce ammonia ($NH_3$). Nitrogen is used as a coolant (liquid nitrogen), as a carrier gas in chromatography, and in fertilizer manufacturing.

Chlorine ($Cl_2$) was isolated by Scheele in $1774$. It is a greenish-yellow, poisonous gas with a choking smell. Industrially, it is prepared by the electrolysis of brine or molten chlorides using the Castner-Kellner or Solvay cells. Chemically, chlorine is a powerful oxidizing agent that displaces other halogens (except fluorine) from their salts. It has a high affinity for hydrogen, removing it from compounds to form hydrogen chloride ($HCl$). It is widely used as a bleaching agent for textiles and wood pulp, and as a germicide for water sterilization.

Hydrogen: Preparation, Properties, and Hydrides

Hydrogen, though nearly absent from the atmosphere in its free state, is abundant in water, living matter, and stars. It has three isotopes: Protium ($^1H$), Deuterium ($^2H$), and Tritium ($^3H$, which is radioactive). Industrially, hydrogen is produced through the steam reforming of methane at $35\,atm$ and $800^\circ C$ with a nickel catalyst ($CH_4(g) + H_2O(g) \rightarrow CO(g) + 3H_2(g)$), or via the Bosch Reaction, where water gas ($CO + H_2$) is reacted with steam over an iron catalyst. In the lab, it is prepared by reacting zinc with dilute $HCl$ or $H_2SO_4$. Hydrogen is the lightest known substance and burns with a characteristic "pop" sound in the presence of a flame.

Hydrogen forms different types of hydrides. Active metals form ionic hydrides (e.g., $NaH$, $CaH_2$), which are crystalline solids that react with water to produce hydroxides and hydrogen gas. Aluminum and boron form complex covalent hydrides like lithium tetrahydridoaluminate ($III$) ($LiAlH_4$), which are vital reducing agents in organic chemistry. Non-metals create simple covalent hydrides, most of which are volatile gases at room temperature, though water ($H_2O$) and hydrogen fluoride ($HF$) are liquids due to hydrogen bonding.

Questions & Discussion

What are ALLOYS? Mention eight ($8$) examples of alloys, stating their constituents and their uses.

1. Define metals and non-metals.
2. Outline, in a tabular form, the differences in physical and chemical properties between metals and non-metals.
3. Show the distribution of metals, non-metals and metalloids in the periodic table.
4. Describe how the following metals are extracted from their ores: (a) Sodium (b) Aluminum.
5. State five chemical properties each of (a) sodium (b) Aluminum.
6. Describe the chemical test for (a) sodium (b) aluminum.
7. Explain how nitrogen can be obtained industrially.
8. With the help of diagram, describe how hydrogen gas can be prepared in the laboratory.
9. State five chemical properties of hydrogen gas.
10. Outline five uses of hydrogen gas.
11. Show how chlorine can be prepared in the laboratory.
12. State the chemical properties of chlorine.
13. Outline the uses of chlorine.