Carbon and its Compounds

Occurrence and Abundance of Carbon

• Carbon exists in both the free state and combined state in nature.

  • Free State Forms: Graphite, diamond, and fullerenes.

  • Combined State Forms: Carbon dioxide ($CO_2$), glucose, carbonates, hydrogen carbonates, coal, petroleum, and organic substances.

• Abundance in Nature:

  • Earth's Crust: Contains $0.02\%$ carbon present in the form of minerals such as carbonates, hydrogen carbonates, coal, and petroleum.

  • Atmosphere: Contains $0.03\%$ carbon dioxide ($CO_2$).

• Despite its small natural abundance, carbon is of immense significance as all living structures and countless daily items are carbon-based.

• Identification of Carbon Compounds (Activity 4.1):

  • Sorting everyday items used or consumed yields three major categories: things made of metal, things made of glass or clay, and others.

  • Most items listed under the "others" category (food, clothes, medicines, books) are compounds of carbon.

  • Burning a carbon-containing compound produces carbon dioxide gas ($CO_2$), which can be confirmed using the lime water test (turning lime water milky).

Properties and Bonding: Ionic vs. Covalent Compounds

• Properties of Ionic Compounds:

  • High melting and boiling points.

  • Conduct electricity in solution or in the molten state.

  • Formation involves the complete transfer of electrons between a metal and a non-metal, producing charged ions held together by strong electrostatic forces.

• Properties of Carbon Compounds:

  • Low melting and boiling points compared to ionic compounds.

  • Largely non-conductors or poor conductors of electricity.

  • Non-conducting nature proves that bonding in carbon compounds does not give rise to charged ions.

  • Low physical transition temperatures indicate that intermolecular forces of attraction (such as Van der Waals forces) between molecules are relatively weak.

• Melting and Boiling Points of Selected Carbon Compounds:

  • Acetic Acid ($CH_3COOH$): Melting point = $290\,K$, Boiling point = $391\,K$

  • Chloroform ($CHCl_3$): Melting point = $209\,K$, Boiling point = $334\,K$

  • Ethanol ($CH_3CH_2OH$): Melting point = $156\,K$, Boiling point = $351\,K$

  • Methane ($CH_4$): Melting point = $90\,K$, Boiling point = $111\,K$

Nature of the Covalent Bond in Carbon

• Electronic Configuration of Carbon:

  • Atomic number of Carbon = $6$.

  • Electron distribution across shells: K shell = $2$, L shell = $4$.

  • Carbon has $4$ valence electrons in its outermost shell (tetravalent).

• Inability to Form Ionic Bonds:

  • Elements react to attain a completely filled outer shell (noble gas configuration).

  • To achieve octet configuration ionically, carbon would need to gain or lose $4$ electrons:

  • Gain of 4 electrons ($C^{4-}$ anion): It would be extremely difficult for the nucleus containing $6$ protons to hold on to $10$ electrons ($4$ extra electrons), leading to nuclear instability.

  • Loss of 4 electrons ($C^{4+}$ cation): It would require a massive amount of energy to remove $4$ electrons, leaving behind a carbon cation with $6$ protons holding on to just $2$ electrons in its K shell.

• Covalent Bond Formation:

  • Carbon overcomes this obstacle by sharing its valence electrons with other carbon atoms or with atoms of other elements.

  • Definition: A covalent bond is a chemical bond formed by the sharing of electron pairs between two atoms.

  • Shared electrons belong to the outermost shells of both bonding atoms, allowing each atom to achieve the stable noble gas configuration.

  • Covalently bonded molecules possess strong intramolecular forces (bonds within the molecule) but weak intermolecular forces.

Covalent Bonding in Homonuclear and Simple Molecules

• Hydrogen Molecule ($H_2$):

  • Atomic number of Hydrogen = $1$ (has $1$ electron in its K shell).

  • Requires $1$ additional electron to fill the K shell and attain the electronic configuration of Helium ($2$ electrons).

  • Two hydrogen atoms share their single valence electrons to form an $H_2$ molecule.

  • The single shared pair of electrons constitutes a single covalent bond, represented by a single line ($H-H$).

• Chlorine Molecule ($Cl_2$):

  • Atomic number of Chlorine = $17$ (electronic configuration: $2, 8, 7$).

  • Valence electrons = $7$; needs $1$ electron to achieve octet.

  • Two chlorine atoms share $1$ pair of electrons to form a diatomic $Cl_2$ molecule via a single covalent bond ($Cl-Cl$).

• Oxygen Molecule ($O_2$):

  • Atomic number of Oxygen = $8$ (electronic configuration: $2, 6$ in L shell).

  • Requires $2$ electrons to complete its octet.

  • Each oxygen atom contributes $2$ electrons, giving rise to $2$ shared pairs of electrons.

  • This constitutes a double covalent bond between the two oxygen atoms ($O=O$).

• Nitrogen Molecule ($N_2$):

  • Atomic number of Nitrogen = $7$ (electronic configuration: $2, 5$).

  • Combining capacity (valency) = $3$; requires $3$ electrons to complete octet.

  • Each nitrogen atom contributes $3$ electrons, giving rise to $3$ shared pairs of electrons.

  • This constitutes a triple covalent bond between the two nitrogen atoms ($N \equiv N$).

• Water Molecule ($H_2O$):

  • Consists of $1$ oxygen atom sharing electron pairs with $2$ hydrogen atoms.

  • Contains two single covalent bonds ($H-O-H$).

• Ammonia Molecule ($NH_3$):

  • Consists of $1$ central nitrogen atom sharing single electron pairs with $3$ hydrogen atoms.

  • Formed by single covalent bonds, allowing all four atoms to achieve noble gas configuration.

• Methane Molecule ($CH_4$):

  • One of the simplest carbon compounds.

  • Widely used as a fuel and serves as the major component of bio-gas and Compressed Natural Gas (CNG).

  • Carbon shares its $4$ valence electrons with $4$ separate hydrogen atoms to form $4$ single covalent bonds.

Allotropes of Carbon

• Definition: Allotropes are different physical forms of the same element existing in nature, exhibiting widely varying physical properties but identical chemical properties.

• Diamond:

  • Structure: Each carbon atom is bonded to $4$ other carbon atoms forming a rigid, three-dimensional tetrahedral structure.

  • Properties: Hardest natural substance known; non-conductor of electricity.

  • Synthetic Diamonds: Created by subjecting pure carbon to very high pressure and temperature. Synthetic diamonds are small but otherwise physically and chemically indistinguishable from natural diamonds.

• Graphite:

  • Structure: Each carbon atom is bonded to $3$ other carbon atoms in the same plane, producing a hexagonal array. One of these bonds is a double bond to satisfy carbon's tetravalency. The hexagonal arrays are stacked in layers one above another.

  • Properties: Smooth and slippery (due to sliding layers held by weak forces); excellent conductor of electricity due to free valence electrons.

• Fullerenes:

  • Class of carbon allotropes forming closed cages.

  • $C_{60}$ (Buckminsterfullerene): The first identified fullerene, consisting of $60$ carbon atoms arranged in the shape of a football.

  • Named after the US architect Buckminster Fuller because the molecule's structural shape resembles his geodesic dome design.

• Sulfur Molecule ($S_8$):

  • Consists of $8$ sulfur atoms joined together in a closed ring arrangement.

Versatile Nature of Carbon

• Millions of carbon compounds are known to chemists, far outnumbering the compounds formed by all other elements combined.

• Carbon's extraordinary compound-forming ability is driven by two key structural features:

1. Catenation

• Definition: Catenation is the unique ability of carbon to form strong covalent bonds with other carbon atoms, giving rise to large molecules.

• Structural forms created by catenation include:

  • Long straight carbon chains.

  • Branched carbon chains.

  • Carbon rings (cyclic structures).

• Carbon atoms in these structures can be linked by single, double, or triple covalent bonds.

• Comparison with Silicon: Silicon forms compounds with hydrogen having chains of up to $7$ or $8$ atoms, but these compounds are highly reactive and unstable. In contrast, the carbon-carbon ($C-C$) bond is exceptionally strong and stable, allowing long carbon chains to exist.

2. Tetravalency and Atomic Size

• Carbon has a valency of $4$, enabling it to bond with $4$ other carbon atoms or monovalent/polyvalent atoms of other elements (e.g., oxygen, hydrogen, nitrogen, sulfur, chlorine, bromine).

• These non-carbon atoms impart specific chemical properties to the compound.

• Effect of Small Atomic Size: The carbon atom is extremely small in size. This allows its nucleus to hold on strongly to shared pairs of electrons, producing strong, stable covalent bonds. Elements with larger atoms form significantly weaker bonds.

History of Organic Chemistry and Vital Force Theory

• Originally, carbon compounds extracted from natural substances were believed to form only within living systems under the influence of a non-physical "vital force".

• Disproof: Friedrich Wöhler disproved the vital force theory in $1828$ by synthesizing urea (an organic compound) directly from ammonium cyanate (an inorganic compound).

• Modern Definition: Carbon compounds continue to be studied under organic chemistry, with the exception of simple forms such as carbides, oxides of carbon, carbonate salts, and hydrogen carbonate salts.

Saturated and Unsaturated Carbon Compounds

• Saturated Carbon Compounds:

  • Carbon compounds in which all carbon atoms are linked exclusively by single bonds.

  • Also known as Alkanes.

  • Generally exhibit low chemical reactivity.

• Unsaturated Carbon Compounds:

  • Carbon compounds containing one or more double bonds or triple bonds between carbon atoms.

  • Hydrocarbons with double bonds are called Alkenes; hydrocarbons with triple bonds are called Alkynes.

  • Considerably more reactive than saturated carbon compounds.

• Step-by-Step Method for Deriving Carbon Structures:

  1. Step 1: Link the carbon atoms together using single covalent bonds ($C-C$).

  2. Step 2: Attach hydrogen atoms to satisfy the remaining valencies of carbon.

  3. Step 3: If valencies remain unsatisfied, form double or triple bonds between carbon atoms.

• Examples:

  • Ethane ($C_2H_6$):

  • Step 1: Link carbons ($C-C$).

  • Step 2: Attach $3$ hydrogen atoms to each carbon atom ($H_3C-CH_3$). All valencies satisfied by single bonds (saturated).

  • Propane ($C_3H_8$):

  • Chain of $3$ single-bonded carbon atoms satisfied with $8$ hydrogen atoms (saturated).

  • Ethene ($C_2H_4$):

  • Step 1: Link carbons ($C-C$).

  • Step 2: Attach $2$ hydrogen atoms to each carbon ($H_2C-CH_2$), leaving $1$ valency unsatisfied per carbon.

  • Step 3: Insert a double bond between the carbons ($H_2C=CH_2$) (unsaturated).

  • Ethyne ($C_2H_2$):

  • Linked carbons require a triple bond ($H-C \equiv C-H$) to satisfy carbon's tetravalency (unsaturated).

Chains, Branches, Rings, and Structural Isomerism

• Straight Chain Saturated Hydrocarbons (Alkanes):

  • Methane ($1$ Carbon): Formula = $CH_4$

  • Ethane ($2$ Carbons): Formula = $C_2H_6$

  • Propane ($3$ Carbons): Formula = $C_3H_8$

  • Butane ($4$ Carbons): Formula = $C_4H_{10}$

  • Pentane ($5$ Carbons): Formula = $C_5H_{12}$

  • Hexane ($6$ Carbons): Formula = $C_6H_{14}$

• Structural Isomers:

  • Definition: Compounds having identical molecular formulas but different structural arrangements of atoms.

  • Example - Butane ($C_4H_{10}$):

  • Straight carbon skeleton: $C-C-C-C$ ($n$-butane).

  • Branched carbon skeleton: $C-C(C)-C$ (isobutane / $2$-methylpropane).

  • Both structures share the identical molecular formula $C_4H_{10}$.

• Cyclic Carbon Compounds (Rings):

  • Carbon atoms arranged in a ring structure; can be saturated or unsaturated.

  • Cyclohexane ($C_6H_{12}$): Saturated ring of $6$ carbon atoms linked by single bonds, with each carbon bonded to $2$ hydrogen atoms.

  • Benzene ($C_6H_6$): Unsaturated ring of $6$ carbon atoms containing alternating single and double bonds, with each carbon bonded to $1$ hydrogen atom.

  • Additional examples of saturated cyclic hydrocarbons: Cyclobutane ($C_4H_8$) and Cyclopentane ($C_5H_{10}$).

• Classification Summary of Hydrocarbons:

  • Hydrocarbons: Compounds consisting exclusively of carbon and hydrogen atoms.

  • Alkanes: Saturated hydrocarbons with single bonds. General formula = $C_n H_{2n+2}$.

  • Alkenes: Unsaturated hydrocarbons containing one or more double bonds. General formula = $C_n H_{2n}$.

  • Alkynes: Unsaturated hydrocarbons containing one or more triple bonds. General formula = $C_n H_{2n-2}$.

Heteroatoms and Functional Groups

• Heteroatom: An atom other than carbon and hydrogen (such as halogens, oxygen, nitrogen, sulfur) that replaces one or more hydrogen atoms in a hydrocarbon chain while keeping carbon's valency satisfied.

• Functional Group: A heteroatom or group of atoms containing heteroatoms that confers specific chemical properties to an organic compound, regardless of the length and nature of the carbon chain.

• Important Functional Groups:

  • Halo- (Chloro / Bromo) alkanes:

  • Heteroatoms: Chlorine ($Cl$), Bromine ($Br$).

  • Formula of group: $-Cl$, $-Br$ (substitutes for H).

  • Alcohol:

  • Heteroatom: Oxygen ($O$).

  • Formula of group: $-OH$

  • Aldehyde:

  • Heteroatom: Oxygen ($O$).

  • Formula of group: $-CHO$ (represented structurally as $-C(=O)H$).

  • Ketone:

  • Heteroatom: Oxygen ($O$).

  • Formula of group: $-C(=O)-$ or $-CO-$ (non-terminal carbonyl group).

  • Carboxylic Acid:

  • Heteroatom: Oxygen ($O$).

  • Formula of group: $-COOH$ (represented structurally as $-C(=O)OH$).

Homologous Series

• Definition: A series of organic compounds in which the same functional group substitutes for hydrogen in a carbon chain of varying lengths.

• Key Characteristics of a Homologous Series:

  • Successive members differ structurally by a $-CH_2-$ (methylene) unit.

  • Successive members differ in molecular mass by $14\,u$ (Carbon = $12\,u$, Hydrogen = $1\,u \times 2 = 2\,u$).

  • All members share similar chemical properties because reactivity is determined by the functional group.

  • Physical properties (melting point, boiling point, density, solubility in specific solvents) show a gradual, consistent change (gradation) as molecular mass increases.

• Examples of Homologous Series:

  • Alcohol Series: Methanol ($CH_3OH$), Ethanol ($C_2H_5OH$), Propanol ($C_3H_7OH$), Butanol ($C_4H_9OH$).

  • Alkane Series: Methane ($CH_4$), Ethane ($C_2H_6$), Propane ($C_3H_8$), Butane ($C_4H_{10}$) (General formula: $C_n H_{2n+2}$).

  • Alkene Series: Ethene ($C_2H_4$), Propene ($C_3H_6$), Butene ($C_4H_8$) (General formula: $C_n H_{2n}$, where $n = 2, 3, 4$

  • Alkyne Series: Ethyne ($C_2H_2$), Propyne ($C_3H_4$), Butyne ($C_4H_6$) (General formula: $C_n H_{2n-2}$).

Nomenclature of Carbon Compounds

• Systematic rules for naming organic compounds based on carbon chain length and functional groups:

  1. Identify Carbon Chain Length: Determine the parent name based on the number of carbon atoms in the longest chain (e.g., $1$ carbon = meth-, $2$ carbons = eth-, $3$ carbons = prop-, $4$ carbons = but-).

  2. Identify Functional Group: Indicated using a prefix or a suffix attached to the parent name.

  3. Modify Suffix for Functional Groups:

  • If the suffix of the functional group begins with a vowel ($a, e, i, o, u$), delete the terminal letter 'e' from the carbon chain name and append the functional group suffix.

  • Example: A $3$-carbon chain with a ketone group: Propane - 'e' + 'one' = Propanone.

  • Alcohols: End with the suffix -ol (e.g., Methanol, Ethanol, Propanol, Butanol).

  • Aldehydes: End with the suffix -al.

  • Carboxylic Acids: End with the suffix -oic acid.

  1. Nomenclature for Unsaturated Chains:

  • Replace the ending '-ane' with '-ene' for double bonds (e.g., $3$-carbon chain with a double bond is Propene).

  • Replace the ending '-ane' with '-yne' for triple bonds (e.g., $3$-carbon chain with a triple bond is Propyne).

• Forms of Carbon: Carbon is found in two main ways:

  • Free State: This includes graphite (used in pencils), diamonds (super hard), and fullerenes (like a soccer ball shape).

  • Combined State: This includes substances like carbon dioxide (found in the air), glucose (sugar that gives us energy), and many other materials like coal and oil.

• Importance of Carbon: Even though carbon is in small amounts (0.02% in earth's crust, 0.03% in atmosphere), it's super important because all living things and many everyday items are made of carbon.

• Identifying Carbon Compounds:

  • When we burn things that contain carbon (like food), they produce carbon dioxide. We can check for carbon dioxide using lime water, which turns cloudy.

• Properties of Compounds:

  • Ionic Compounds:

  • They have high melting and boiling points.

  • Conduct electricity when dissolved in water.

  • Formed when metals give away electrons to non-metals.

  • Carbon Compounds:

  • Lower melting and boiling points compared to ionic compounds.

  • Poor conductors of electricity since they don’t form charged particles (ions).

  • Examples include acetic acid (vinegar), chloroform (used as solvent), ethanol (alcohol), and methane (natural gas).

• Bonding in Carbon:

  • How Carbon Bonds: Carbon has 4 outer electrons and prefers to share them to form bonds instead of giving them away.

  • This sharing creates covalent bonds, which are strong and help form stable compounds.

  • Examples of Covalent Bonds:

  • Hydrogen molecule (H₂): Two hydrogen atoms share their one electron each.

  • Chlorine molecule (Cl₂): Two chlorine atoms share one electron each.

  • Oxygen molecule (O₂): Two oxygen atoms share two electrons each to form a double bond.

  • Nitrogen molecule (N₂): Two nitrogen atoms share three electrons each to form a triple bond.

  • Water (H₂O): One oxygen atom shares electrons with two hydrogen atoms.

  • Ammonia (NH₃): One nitrogen shares electrons with three hydrogen atoms.

• Allotropes of Carbon:

  • Diamond: Has a strong structure, making it the hardest natural substance, doesn’t conduct electricity.

  • Graphite: Can conduct electricity, slippery because of its layered structure.

  • Fullerenes: Carbon in shapes like cages or spheres.

• Versatile Nature of Carbon:

  • Carbon can bond with itself to make long chains or rings, known as catenation.

  • It can also bond with many different elements due to its ability to form four bonds (tetravalency).

• History of Organic Chemistry:

  • People used to think that carbon compounds only came from living things due to "vital force."

  • This was shown to be wrong when a scientist made urea (a carbon compound) from non-living ingredients in 1828.

• Types of Carbon Compounds:

  • Saturated Compounds: Have only single bonds (like alkanes). They are generally stable and less reactive.

  • Unsaturated Compounds: Contain double or triple bonds (like alkenes and alkynes). They are more reactive.

• Structural Isomerism:

  • What are Structural Isomers?: Compounds that have the same formula but different structures. Example: Both butane and isobutane have the same formula (C₄H₁₀) but different arrangements of carbon atoms.

• Functional Groups:

  • Atoms other than carbon and hydrogen that give compounds special properties.

  • Common functional groups include haloalkanes (like chlorinated compounds), alcohols (like ethanol), aldehydes (like formaldehyde), and acids (like acetic acid).