Compound Structure, Bonds, Allotropes

This is part of preliminary HSC Chemistry course under the topic of Bonding

HSC Chemistry Syllabus

  • Investigate the differences between ionic and covalent compounds through:
    • Using nomenclature, valency, and chemical formulae (including Lewis dot diagrams) (ACSCH029)
    • Examining the spectrum of bonds between atoms of varying degrees of polarity with respect to their constituent elements' positions on the periodic table. 
  • Investigate elements that possess the physical property of allotropy
  • Investigate the different chemical structures of atoms and elements, including but not limited to:
    • Ionic networks
    • Covalent networks (including diamond and silicon dioxide)
    • Covalent molecular
    • Metallic Structure

    Compound Structure, Bonds, Allotropes

    Ionic Compounds

    Ionic compounds are compounds that have structures which are made up on ions – particles that have either gained or lost electrons. Metals will typically lose electrons to form positively charged cations, while non-metals will gain electrons to form negatively charged anions. Because of the complementary nature of these ions, ionic compounds commonly consist of a combination of metal(s) and non-metal(s). Ionic compounds can also be classified as either binary (containing one cation and an anion) or polyatomic (containing multiple cations and anions). 

    These types of compounds form lattice structures, where ions organise into a regular, repeating arrangement to generate a crystal. The diagram below illustrates the lattice structure of sodium chloride (NaCl), an ionic compound. It shares a likeness to scaffolding, regularly seen in construction sites.

     

    Ionic bonds, the "glue" holding ions in the ionic lattices together, form when valence electrons completely transfer from one atom (usually a metal) to another (typically a non-metal). Unlike covalent bonds where electrons are shared, this complete exchange leads to a more stable structure with noble-gas-like electron configurations for both participating atoms. This transfer of electrons results from the electrostatic attraction between oppositely charged ions, which is why ionic bonds are generally stronger than covalent bonds.

    Ionic crystals, structured by a continuous 3D arrangement of cations and anions, are held together by these robust ionic bonds. Due to the strong lattice structures, ionic solids tend to have high melting points and hardness. However, their lattice structure also contributes to their brittleness; when subjected to shearing forces, they can shatter as similarly charged ions in close proximity repel each other

    Ionic lattices, unlike metallic ones (discussed below), lack free electrons or ions to carry a charge when a voltage is applied. Therefore, in their solid state, they do not conduct electricity. However, when melted or dissolved in an appropriate solvent such as water, the ions detach from the lattice and become mobile. Consequently, solutions of ionic substances or molten ionic compounds can conduct electricity. The solubility of ionic substances can vary greatly, depending on the substance's polarity.

     

      In summary, ionic compounds generally: 

        • Boast high melting and boiling points due to their strong ionic bonds.
        • Are solid at room temperature, again due to their strong bonding.
        • Are hard yet brittle, susceptible to shattering under shearing forces.
        • Do not conduct electricity in the solid state, as they lack free charge carriers.
        • Become good electrical conductors in the liquid state or when dissolved in a solvent, due to their dissociation into free ions.
        • Vary in their solubility.



      Ionic Structure & Empirical Formulae

      Ionic substances, during crystal formation, align themselves in distinct arrangements known as close packing arrays. These arrays typically assume either a hexagonal or a cubic shape. However, for the scope of high school education, we'll primarily concentrate on cubic structures. A prime example of a compound exhibiting a cubic packing structure is sodium chloride (NaCl).

       



      Ionic Structures and Empirical Formulas

      The empirical formula of a compound denotes its atomic or ionic composition, expressed in the simplest whole number ratio.


      Examples:

        1. Ethane, a covalent molecular compound, has a molecular formula of `C_2H_6`. Each molecule consists of two carbon atoms bonded to six hydrogen atoms. The C:H ratio is 2:6, which can be simplified to 1:3. Therefore, the empirical formula of ethane is `CH_3`.
        2. Benzene, another covalent molecular compound, has a chemical formula of `C_6H_6`. The C:H ratio is 6:6, which can be reduced to 1:1. Hence, the empirical formula for benzene is CH.


      Unlike covalent compounds, ionic compounds arrange themselves in lattice structures. Considering the illustration above, you can observe that in a single cubic unit there are:

      • One entire sodium ion located at the center of the cube.
      • Twelve sodium ions, each occupying a quarter of the 12 edges of the cube.
      • Eight chloride ions, each taking up an eighth of the cube's corners.
      • Six chloride ions, each occupying half of the centers of each face of the cell.


      Thus, for each unit of sodium chloride, there are four sodium ions and four chloride ions. This leads to an empirical formula of NaCl for sodium chloride.

      Ion Formation (Cations & Anions) and Bond Creation

      Ionic bonds are formed during the ionisation of a compound's constituents, resulting in cations (positively charged ions) and anions (negatively charged ions). Commonly, a metal atom loses an electron to become a cation, while a non-metal atom gains an electron to become an anion.

      Consider the formation of sodium fluoride (NaF) as an example. Sodium, a metal, and fluorine, a non-metal, react together. In this reaction, the sodium atom relinquishes its single valence electron to the fluorine atom. This addition fulfills the 'octet rule' in fluorine, making its outer shell complete. The newly formed ions, having opposite charges, are attracted to each other due to the electrostatic forces between them.

      The process of valence electrons being transferred from a metal to a non-metal generates a cation and an anion. The resulting ionic bond creates a crystalline lattice structure, typically resulting in a solid-state compound under standard conditions.

      The charge on the ions is influenced by the number of transferred electrons needed to achieve a stable noble gas configuration. For an ionic compound to exist, the overall charge must be neutral. A good example is the formation of calcium chloride (`CaCl_2`) from calcium and chlorine.

      By inspecting the electronic configurations of these atoms, we can understand how the formation of calcium chloride results in an overall neutral charge:

      Ca: [Ar]`4s^2`

      Cl: [Ne]`3s^2``3p^5`


      To reach a noble gas configuration, the calcium atom has to lose its two valence electrons in the 4s orbital. Conversely, chlorine atoms, having seven valence electrons, need an additional electron to complete their outer shell. With two chlorine atoms, the calcium atom can distribute its two electrons to fulfill the octet rule for all involved atoms. This principle is applicable to any other halogens that may replace chlorine in the compound.

      The diagram above illustrates how a calcium atom reacts with chlorine by donating its two electrons to help complete the octet of the chlorine atoms. This type of diagram is called an electron transfer diagram.

      Covalent Bonding & Compounds

      Covalent compounds encompass a diverse range of substances composed solely of non-metal elements. These compounds, such as nitrogen dioxide (NO2), arise from the formation of covalent bonds - shared pairs of electrons between atoms, leading to the formation of molecules.

      Covalent Bonds

      Covalent bonds are essentially shared electron pairs between atoms or ions, typically formed between non-metals with similar electronegativities. Some exceptions to this rule may involve covalent bonds between metals and non-metals, as seen in Aluminum Chloride (`AlCl_3`). The electron pairs participating in covalent bonds are often referred to as "bonding electrons." Covalent bonds manifest in three distinct forms: single, double, and triple bonds.

      • Single Bond

      A single bond forms when a pair of electrons is shared between two atoms. It is symbolised by a single line linking the two atoms. Single bonds are the most unreactive of all bonds due to the stability derived from the shared electron pair.


      • Double Bond

      A double bond arises when two electron pairs are shared between two atoms, represented by two parallel lines between the atoms in a molecule. Though stronger than a single bond, double bonds are considerably more reactive due to the electron-rich centre, a result of loosely-held mobile electrons in the weak pi bond—one of the two component bonds in a double bond.


      • Triple Bond
      A triple bond is formed when three pairs of electrons are shared between two atoms in a molecule. It represents the strongest and most reactive bond type owing to its electron-rich center. Analogous to double bonds, two of the three constituent bonds in a triple bond are weak pi bonds.

       

      Covalent Molecular Substances

      Covalent molecular compounds encompass a range of substances, including ice and dry ice (solid carbon dioxide), which are made up of discrete molecules arranged in diverse geometric configurations. These molecules are held together within the lattice by intermolecular forces, which are inherently weaker than standard chemical bonds.

      Due to the absence of mobile charge carriers like electrons or ions, covalent molecular substances do not conduct electricity, even when melted or dissolved in a solvent. Moreover, covalent molecular crystals, seen in elements such as sulfur, iodine, and phosphorus, are susceptible to shearing forces.


      Covalent Network Lattice

      Covalent network compounds, exemplified by substances like diamond, comprise a three-dimensional grid of atoms interconnected by robust covalent bonds. This interconnected lattice forms a remarkably strong and rigid structure, often reflected in the substance's hardness. Nonetheless, akin to ionic crystals, these structures can shatter under intense shearing forces.

      Covalent network compounds and elements, including diamond, boast exceptionally high melting points. For instance, silicon dioxide melts at a scorching 1713°C. Owing to the lack of free electrons, these network lattices do not conduct electricity in their solid state. This absence of free electrons also contributes to their high insolubility in most common solvents, including water.

      The hardness of diamond can be attributed to the four strong covalent bonds each carbon atom forms with four other carbon atoms. In contrast, a buckminsterfullerene, another form of carbon, forms bonds with only three other carbon atoms at most, resulting in a substance that is softer than diamond

       

      The compound shown above is silicon dioxide, a covalent network substance featuring a lattice of silicon and oxygen atoms. Naturally occurring as the primary component of quartz and present in some living organisms, its intrinsic hardness can be attributed to its robust lattice structure

       

      Property

      Reasoning

      Non-conductors

      There are no mobile charge carriers in the form of either ions or electrons, which are able to carry heat or charge through the material

      Very Brittle and Hard

      The strong covalent bonds of the lattice help to make it hard, but become susceptible to strong shearing forces

      Very High M.P and B.P

      Since high energy is required to break the strong covalent bonds holding the atoms together in the structure, they are not susceptible to heat until extreme temperatures.

       

      Properties of Covalent Compounds

      1. Covalent compounds, more often than not, possess low melting and boiling points. This characteristic enables them to exist in gaseous or liquid states at room temperature under standard pressure, a consequence of their constituent particles being bound together by comparatively weak intermolecular forces that readily yield to thermal energy.

      2. Covalent network structures are typically hard yet brittle, subject to shearing forces due to the inherent stress exerted on their lattice structure.

      3. Predominantly, covalent substances are non-conductive. This is attributed to the absence of free electrons or charge carriers - essential elements for electricity conduction. Consequently, they remain non-conductive even when dissolved in water.

      4. Generally, covalent compounds exhibit poor solubility in water. This is because most are non-polar and thus are insoluble in polar substances like water. As a principle, polar compounds tend to dissolve in polar solvents, and non-polar compounds dissolve in non-polar solvents. This is exemplified by the immiscibility of oil and water - oil being non-polar does not mix with water, leading it to float on water's surface.

      5. Noble gases from Group 8 primarily exist in monoatomic form, as they already possess a complete outer valence electron shell. This trait negates the necessity for them to engage in chemical reactions.



      Metallic Bonding Model

      Metals, barring mercury, are predominantly solid at room temperature, exhibiting relatively high melting points and excellent electrical conductivity. Many are hard substances, although certain members of Groups 1 and 2, such as alkali and alkali earth metals including sodium and calcium, show softer properties, making them easily cut with a knife.

      The properties metals exhibit are a result of their structure: a meticulously arranged three-dimensional lattice of positive ions interlaced with a mobile 'sea' of delocalised electrons.

      To comprehend the origins of these structures, it's essential to understand the formation of positive ions in metals. These ions arise when valence electrons dissociate from their parent atoms, leaving behind positive ions. The detached electrons, now "free," are termed "delocalised" as they no longer belong to any specific atom's valence shell. Instead, they meander randomly through the lattice and are shared among a multitude of positive ions.

      Electricity conduction necessitates a medium capable of facilitating a continuous flow of electrons. In metals, due to the ability of delocalised electrons to roam freely throughout the lattice, they are predominantly excellent conductors of electricity. When discussing electricity flow through a metallic wire, we're essentially referring to the movement of electrons through a conductive material.

      Metals demonstrate malleability, meaning they can be bent, rolled into sheets, and drawn into wires or rods. This property arises because, when metal's positive ions undergo shearing, the mobile electrons can adapt to the new arrangement of positive ions and 'cement' the entire structure.

      The diagram below presents a 2D  representation of metallic structures at the microscopic level. This kind of bonding, which involves positive metal ions and delocalised electrons, is known as metallic bonding.

       

       

       

      Property of metal

      Explanation

      Hardness

      The strong electrostatic forces of attraction between the positive ions and the sea of delocalized electrons in the metal helps to hold the metallic lattice together. Generally we can assume that if a substance is harder, it must have stronger bonds.

      Conductivity

      Because of the presence of a sea of free-moving, delocalized electrons are able to move towards positive electrodes and away from a negative ones

      Malleability and ductility

      When a force causes metal ions to move past each other, layers of ions are still held together by their electrostatic attraction to the delocalized electrons between them

      Metals are highly lustrous

      The surface structure of metals allows light to reflect readily. Light photons are rapidly absorbed and released by the mobile electron cloud.

       

      Characteristics of Transition Metals

      When juxtaposed with main group metals, transition metals exhibit certain distinct properties:

      1. Hardness: Transition metals typically possess a high degree of hardness.

      2. High Tensile Strength: They exhibit robust tensile strength, making them resistant to deformation and breakage under tension.

      3. Elevated Densities: These metals are generally denser compared to main group metals.

      4. High Melting Points: Transition metals have notably higher melting points. This can be attributed to the additional d orbital electrons that contribute to a stronger core charge, enhancing the electrostatic attraction between the nucleus and the outer electrons. Consequently, these metals have smaller atomic radii, enabling them to pack more densely within the lattice structure.

      5. Magnetic Properties: Certain transition metals display magnetic properties.

      6. Lustrous Appearance: Generally, transition metals bear a shiny or lustrous appearance. This characteristic stems from the presence of free electrons, which facilitate the rapid release of wavelengths within the visible light spectrum.


      Nomenclature and Polyatomic Ions

      Polyatomic ions are unique entities consisting of two or more atoms of differing elements bound together.

      Characteristics of polyatomic ions include:

      1. Fixed Ratio of Elements: The elements within polyatomic ions always maintain a constant ratio.

      2. Unified Behavior: Despite being composed of multiple elements, these ions behave as a singular group or unit in chemical reactions.

      3. Denoted by Subscripts and Superscripts: The quantity of each element present within the ion is signified by subscripts, whereas the overall charge of the ion is represented by superscripts. For instance, the carbonate polyatomic ion, `CO_3^{2-}`, comprises one carbon atom, three oxygen atoms, and carries a charge of -2.

      To successfully navigate ionic compound nomenclature, it is necessary to acquaint yourself with several polyatomic ions. The table below lists some commonly encountered polyatomic ions.

      NH­4+ – Ammonium

      CH3COO- - Acetate

      CN- - cyanide

      HCO3- - hydrogen carbonate

      HS- - hydrogen sulfide

      HSO4- - hydrogen sulfate

      OH- - hydroxide

      NO3- - nitrate

      NO2- - nitrite

      MnO4- - permanganate

      CO32- - carbonate

      CrO42- - chromate

      Cr2O72- - dichromate

      HPO42- - hydrogen phosphate

      C2O42- - oxalate

      SO32- - sulfite

      SO42- - sulfate

      PO43- - phosphate

       

      1. Ionic Compounds: Cation + Anion

      • Naming Protocol:

        • In ionic compounds, the cation is typically named first, followed by the anion. For instance, in the ionic compound NaCl, the cation is the metal Sodium, while the anion is the non-metal Chlorine.

      • Suffix Change: 

        • When the anion is a single non-metal, its name generally concludes with the suffix "-ide". Hence, for the ionic compound NaCl, it would be named Sodium Chloride, where 'Chlorine' is replaced with 'Chloride'.

      • Polyatomic Anion:

        • If the anion is polyatomic, its suffix remains unchanged. For example, CaCO3 is named Calcium Carbonate, as CO3^2- is recognized as the 'carbonate' anion.

      • Polyatomic Cation

        • When the cation is polyatomic, the same naming rules apply. For instance, NH4Cl is named Ammonium Chloride, where NH4+ is identified as the 'ammonium' cation.

      • Transition Metals:

        • Transition metals, capable of forming ions with varying charges, necessitate the charge indication by inserting a Roman numeral within parentheses after the metal name. For example, iron oxide, comprising of Fe and O, can be ambiguous as iron can exist as either `Fe^{2+}` or `Fe{3+}`, altering the compound formula to either FeO or `Fe_2O_3`. Therefore, if the iron ion is `Fe^{3+}`, the formula would be `Fe_2O_3`, named as Iron(III) Oxide. Conversely, if the iron ion is `Fe^{2+}`, the formula would be FeO, and it would be named Iron(II) Oxide

      1. Covalent Compounds: Non-Metal + Non-metal

      • Naming Order:

        • For covalent compounds, the less electronegative element is usually named first. This rule will typically be provided.

      • Utilizing Prefixes:

        • Numerical prefixes like mono, di, tri, tetra, and so on are used to indicate the number of atoms of each element in the compound.

      • Addition of Suffix:
        • The name of the compound generally ends with the suffix "-ide" or "-ine".
      • Example 1: Consider a compound consisting of two atoms of oxygen and one atom of nitrogen. As Nitrogen is less electronegative, it is named first, followed by the prefix "di-" for two oxygen atoms, leading to 'dioxide'. Hence, the compound `NO_2` is named Nitrogen Dioxide.
      • Example 2: In a compound comprising two chlorine and seven oxygen atoms, Chlorine being the less electronegative element is named first, using the prefix "di-" for two atoms, leading to 'dichlorine', and then the prefix "hepta-" for seven oxygen atoms, forming 'heptoxide'. Therefore, the compound `Cl_2O_7` is named Dichlorine Heptoxide.

      Allotropy

      Allotropy refers to the phenomenon in which an element can exist in more than one distinct crystalline form, known as allotropes. Allotropes are manifestations of a single element in the same physical state (solid, liquid, or gas) that showcase significantly different physical properties such as density, hardness, electrical conductivity, and color. Elements such as carbon, sulfur, phosphorous, and oxygen exhibit allotropy.

      Allotropes typically emerge due to the differences in stability at specific temperature and pressure conditions. In other words, certain forms of an element may be more stable under a particular set of environmental conditions, leading to the formation of different allotropes.


      Allotropes of Carbon
       


      Diamond:

      One of the allotropes of carbon, diamond, boasts a robust covalent network structure. This structure comprises a limitless matrix of carbon atoms, forming a continuous chain of tetrahedral configurations. The specific alignment of the carbon atoms' layers plays a pivotal role in defining the unique properties of diamond.

      In a diamond, each carbon atom is covalently bonded to four other carbon atoms at the vertices of a tetrahedron. This strategic formation imparts extraordinary strength to the structure, making it nearly impervious to breakage. Consequently, diamonds possess a remarkable hardness, elevated melting points, high boiling points, and display low reactivity.

      Interestingly, the tetrahedral configuration of a diamond also harbors gaps within its structure. These gaps or voids contribute to the diamond's optical transparency, allowing light to pass through it unhindered.

      Property

      Reasoning

      Hard

      The three-dimensional lattice structure formed from a repetition of tetrahedral structures connected by strong covalent bonds between the carbon atoms make it incredibly difficult to break

      High M.P and B.P

      Because of the strength of the intermolecular bonds (bonds within molecules) between the carbon atoms, there must be lots of energy required to break them

      Transparency

      The carbon atoms are arranged in a very orderly fashion formed by tetrahedral structures with large spaces between them. Because of this, light is able to pass between the atoms to give diamond its colourless appearance and high light-refractive index.

      Not electrically conductive

      Since diamond is a covalent network substance, this means that the bonds between the carbon atoms are covalent bonds and the valence electrons of the carbon atoms are not free to move.

      High Density

      The structure of the diamond shows that the atoms are tightly bound in a strong three dimensional lattice structure.

      Insoluble

      Diamond is extremely hard and unreactive due to the covalent network structure and a lack of free valence electrons. Thus there is no substance that will be able to chemically react with diamond.

      Excellent Thermal Conductivity

      Since the carbon atoms in diamond are strongly bonded via covalent bonds which are precisely aligned, the diamond is known as an ideal crystal. The atoms in the crystal lattices of solids vibrate which allow for thermal conduction but since the carbon atoms in diamonds are so ideally constructed, they do not interact with one another.

       


      Graphite:

      Graphite's structure significantly diverges from that of diamond, which in turn substantially influences its distinct physical and chemical properties. Although graphite, like diamond, is a covalent network solid, it possesses a unique architecture. Each carbon atom in graphite is covalently bonded to merely three other carbon atoms, unlike the four in diamond.

      The structural design of graphite comprises layers of six-membered carbon rings. Intriguingly, one electron from each carbon atom is not engaged in a covalent bond. These surplus valence electrons amalgamate to form a 'sea' of delocalized electrons. This sea of electrons grants graphite its electrical conductivity, a distinctive property not shared by diamond.

       

      Property

      Reasoning

      Good Lubricant

      Because the carbon atoms in graphite are connected in hexagonal rings which stack in sheets which are held together by weak dispersion forces, the layers can easily slip past each other to give graphite its slippery nature. This is the same reason why graphite is used in pencils as this property allows the substance to mark paper when pressure is applied.

      Extremely Soft

      For the same reason graphite is a good lubricant, the weak dispersion forces between the layers means that the layers of carbon easily slide over each other to make it soft and greasy.

      Medium Density

      Graphite’s density is less than that of diamond. This is due to the structural layout. The layers of the graphite are separated by large distances since the forces which hold the layers together are weak and unable to bind them together tightly.

      High M.P and B.P

      Graphite can still be considered as a covalent network substance despite not having any bonding in the vertical direction between layers. The carbon atoms are connected via strong covalent bonds which extend throughout the horizontal lattice. These intramolecular bonds (bonds between multiple molecules) are hard to break and thus more emergy is required to break them.

      Good Electrical Conductor

      As previously explained, because each carbon atom in graphite is bound to only three other atoms, there is one single electron that is not covalently bound for each of them. These extra valence electrons form a sea of delocalized electrons which helps to make it an excellent electrical conductor. However electricity is only conducted along the plane of the layers so at 90 degrees it is unable to conduct.  

       

      Uses of Diamond

      Use

      Relevant property

      Jewellery

      Lustrous, High light refractive index, hard, scratch resistant, transparent

      Industrial cutting tools and drills

      Because diamond is the hardest substance on earth and has a high melting point, this means that it can cut through anything without melting

      Heat sinks

      Because Of the high thermal conductivity of diamond, it is able to quickly extract heat from any sensitive areas such as computer chips.

       

      Uses of Graphite 

      Use

      Relevant property

      Lead Pencils

      Soft and slippery with layers easily separated

      Electrodes

      Good electrical conductivity

      Polishes and paints

      Soft, slippery nature, metallic luster

      Lubricant in machines

      Slippery nature and high melting point

      Dry cell battery

      Good electrical conductivity, high melting point

       

      Allotropes of Oxygen

      Property

      O2

      O3

       

      Colour

      Colourless

      Blue

       

      Boiling Point

      -183°C

      -111°C

      The boiling point of the diatomic oxygen is lower than that of the ozone since is has a lower molecular mass

      Solubility in water

      Sparingly Soluble

      More soluble than oxygen

      Non-Polar O2 is unable to form strong intermolecular forces in the polar water while O3has a bent structure allowing for some polarity.

      Bonding

      One double covalent bond

      Single bond and a partial bond

      Ozone is much more reactive that oxygen

      Chemical stability

      Stable

      Highly reactive

      Ozone is easily decomposed into much stabler oxygen molecules

      2O­3(g) 3O2(g)

      Oxidation ability

      Weak oxidant

      Strong Oxidant

      Oxidation is the reaction with oxygen to form an oxide. Since only Ozone is highly reactive, it more readily reacts with metal to form metal oxides.

       

      Allotropes of Phosphorus

      Phosphorus possesses three distinct allotropes - white, red, and black - each exhibiting unique physical and chemical properties.

      White phosphorus, characterised by its garlic-like odor, is a white, waxy solid, the most reactive, least stable, and most toxic of the phosphorus allotropes. Its high reactivity with sunlight and air is utilised in the manufacture of smoke, illumination, and incendiary munitions, where it serves as a primary combustible element. Structurally, white phosphorus takes the form of a tetrahedron, denoted as P4, where four phosphorus atoms interlink to form a regular tetrahedral shape.

      Red phosphorus is a derivative of white phosphorus, obtained by heating the latter between 270-300°C in an air-free environment. This conversion enhances phosphorus's density and melting point while significantly reducing its reactivity and toxicity. Such improved stability and ease of handling have made red phosphorus the preferred allotrope in the manufacture of commonplace matchsticks. Like its white counterpart, red phosphorus also assumes a P4 tetrahedral structure and sublimes upon heating. Interestingly, red phosphorus exhibits insolubility in solvents that readily dissolve white phosphorus.

      Black phosphorus is the most stable allotrope of the three, featuring a structure composed of puckered layers of phosphorus atoms, each bonded to three other phosphorus atoms. Strong covalent intermolecular bonds exist between phosphorus atoms within each layer, while weak dispersion forces hold the layers together. This layered structure affords black phosphorus greater stability and reduced reactivity compared to red phosphorus.

       

      Allotropes of Sulfur

      Sulfur exists in two primary forms: Crystalline and Amorphous. The Crystalline category includes Rhombic and Monoclinic sulfur, while the Amorphous category encompasses Plastic, Colloidal, and Milk of Sulphur.

      1. Rhombic Sulfur: This is a yellow, translucent crystalline form of sulphur, stable at temperatures below 96°C, making it the most prevalent form under ordinary conditions. It has a melting point of 114°C and a density of 2.08g/cm³. Above 96°C, it transitions to Monoclinic sulphur.

      2. Monoclinic Sulfur: Characterised by its transparent, amber crystals, this allotrope is less stable than its Rhombic counterpart below 96°C. However, it's predominant in conditions above this temperature. It melts at 119°C and has a density of 1.98g/cm³.

      3. Colloidal Sulfur: This form is obtained by passing hydrogen sulphide through a cooled, saturated solution of sulfur dioxide in water. It serves as a solvent in carbon disulphide and is commonly used in certain medications.

      4. Milk of Sulfur: This non-crystalline, white compound is produced by reacting weak hydrochloric acid with ammonium sulphide or by boiling sulfur with an aqueous solution of calcium hydroxide. Upon heating, it changes to the conventional yellow colour of sulphur. It is soluble in carbon disulphide.

      5. Plastic Sulfur: This form features long chains of sulfur atoms, singly bonded to one another. Above 120°C, it exists as a liquid, transitioning into a dark, viscous substance above 200°C due to the breaking of its sulfur rings into chains. When rapidly cooled in cold water, it forms a soft, elastic material, as the molecules don't have sufficient time to rearrange themselves into rhombic crystals. However, over time, it reverts to the rhombic form at room temperatures.

      Summary:

      Comparison of physical properties of different types of chemical structures

      Physical Property

      Metallic

      Ionic

      Covalent Molecular

      Covalent Network

      Hardness

      Variable

      Hard

      Soft

      Very Hard

      Malleability/ductility

      Malleable, Ductile

      Brittle

      Brittle

      Brittle

      MP/BP

      High

      High

      Low

      Very High

      Conductivity (solid)

      High

      Low

      Low

      Low

      Conductivity (liquid)

      High

      High

      Low

      Low

       

      Ionic

      Property

      Reasoning

      Non-conductors

      There are no mobile charge carriers in the form of either ions or electrons, which are able to carry heat or charge through the material

      Very Brittle and Hard

      The strong covalent bonds of the lattice help to make it hard, but become susceptible to strong shearing forces

      Very High M.P and B.P

      Since high energy is required to break the strong covalent bonds holding the atoms together in the structure, they are not susceptible to heat until extreme temperatures.

       

      Covalent Network

      Property

      Reasoning

      Non-conductors

      There are no mobile charge carriers in the form of either ions or electrons, which are able to carry heat or charge through the material

      Very Brittle and Hard

      The strong covalent bonds of the lattice help to make it hard, but become susceptible to strong shearing forces

      Very High M.P and B.P

      Since high energy is required to break the strong covalent bonds holding the atoms together in the structure, they are not susceptible to heat until extreme temperatures.

       

      Covalent molecular

      Property

      Reasoning

      Soft and Brittle

      Forces between molecules are weak intermolecular attractions

      Low conductivity

      Molecules are uncharged, and electrons are localised in covalent bonds

      Low melting and boiling points

      Forces between molecules are weak intermolecular attractions

        

      Metallic

      Property of metal

      Explanation

      Hardness

      The strong electrostatic forces of attraction between the positive ions and the sea of delocalized electrons in the metal helps to hold the metallic lattice together

      Conductivity

      Because of the presence of a sea of free-moving, delocalized electrons are able to move towards positive electrodes and away from a negative ones

      Malleability and ductility

      When a force causes metal ions to move past each other, layers of ions are still held together by their electrostatic attraction to the delocalized electrons between them

      Metals are highly lustrous

      The surface structure of metals allows light to reflect readily. Light photons are rapidly absorbed and released by the mobile elctron cloud.

       

       

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