11th Standard - Chemistry Tutorial - NCERT - CBSE Pattern

We will cover below chapters.
  • Unit 1 - Some Basic Concepts of Chemistry
  • Unit 2 - Structure of Atom
  • Unit 3 - Classification of Elements and Periodicity in Properties
  • Unit 4 - Chemical Bonding and Molecular Structure
  • Unit 5 - Thermodynamics
  • Unit 6 - Equilibrium
  • Unit 8 - Organic Chemistry - Some Basic Principles and Techniques
  • Unit 9 - Hydrocarbons

Unit 1 - Some Basic Concepts of Chemistry

Chemistry is the science of atoms and molecules. Drugs like cisplatin and taxol are used in cancer therapy. The drug AZT (Azidothymidine) is used for helping AIDS patients. CFCs (chlorofluorocarbons) is responsible for ozone depletion in the stratosphere. But Chemistry can help us create the alternatives that will be more environment friendly. We are still looking for ways to reduce the impact of green house gases like methane, carbon dioxide!


Anything which has mass and occupies space is called matter. Here are some characteristics of a matter!
  • States of Matter - Matter can exist in 3 states - solid, liquid and gas. Solids have definite volume and definite shape. Liquids have definite volume but do not have definite shape. They take the shape of the container in which they are placed. Gases have neither definite volume nor definite shape. They completely occupy the space in the container in which they are placed. On heating, a solid usually changes to a liquid, and the liquid on further heating changes to gas (or vapour).
  • Classification of Matter - Mixture and Pure Substance. In a homogeneous mixture , the components completely mix with each other. This means particles of components of the mixture are uniformly distributed throughout the bulk of the mixture and its composition is uniform throughout. Sugar solution and air are the examples of homogeneous mixtures. In contrast to this, in a heterogeneous mixture , the composition is not uniform throughout and sometimes different components are visible. For example, mixtures of salt and sugar, grains and pulses along with some dirt (often stone pieces), are heterogeneous mixtures. A pure substance is defined as a substance that is composed of a single type of particle with a uniform and definite composition. Copper, silver, gold, water and glucose are some examples of pure substances. Glucose contains carbon, hydrogen and oxygen in a fixed ratio and its particles are of same composition. Hence, like all other pure substances, glucose has a fixed composition. A pure substances can be divided into 2 types - Elements and Compounds! Particles of an element consist of only one type of atoms. These particles may exist as atoms or molecules.o2o_2 , h2h_2, O, N, H fall under Element type! When two or more atoms of different elements combine together in a definite ratio, the molecule of a compound is obtained. Water ( h2oh_2o ), ammonia, carbon dioxide (co2co_2), sugar fall under compound! The properties of a compound are different from those of its constituent elements. For example, hydrogen and oxygen are gases, whereas, the compound formed by their combination i.e., water is a liquid. It is interesting to note that hydrogen burns with a pop sound and oxygen is a supporter of combustion, but water is used as a fire extinguisher.
Here are some properties of a matter!
  • Physical properties such as colour, odour, melting point, boiling point, density. Measurement can be done using English, Metric or SI system .
  • Chemical properties like composition, combustibility, ractivity with acids and bases
The combination of different atoms is governed by basic laws of chemical combination
  • Law of Conservation of Mass
  • Law of Definite Proportions
  • Law of Multiple Proportions
  • Gay Lussac's Law of Gaseous Volumes
  • Avogadro Law

Dalton's atomic theory

This states that atoms are building blocks of matter. The atomic mass of an element is expressed relative to 12C^{12}C   isotope of carbon, which has an exact value of 12u. Usually, the atomic mass used for an element is the average atomic mass obtained by taking into account the natural abundance of different isotopes of that element. The molecular mass of a molecule is obtained by taking sum of the atomic masses of different atoms present in a molecule. The molecular formula can be calculated by determining the mass per cent of different elements present in a compound and its molecular mass. The number of atoms, molecules or any other particles present in a given system are expressed in the terms of Avogadro constant (6.022x1023)(6.022 x 10^23) . This is known as 1 mol of the respective particles or entities.

Chemical Reactions

Chemical reactions represent the chemical changes undergone by different elements and compounds. A balanced chemical equation provides a lot of information. The coefficients indicate the molar ratios and the respective number of particles taking part in a particular reaction. The quantitative study of the reactants required or the products formed is called stoichiometry. Using stoichiometric calculations, the amount of one or more reactant(s) required to produce a particular amount of product can be determined and vice-versa. The amount of substance present in a given volume of a solution is expressed in number of ways, e.g., mass per cent, mole fraction, molarity and molality.

Unit 2 - Structure of Atom

The first atomic theory, proposed by John Dalton in 1808, regarded atom as the ultimate indivisible particle of matter. Towards the end of the nineteenth century, it was proved experimentally that atoms are divisible and consist of three fundamental particles: electrons, protons and neutrons. The discovery of sub-atomic particles led to the proposal of various atomic models to explain the structure of atom.

Black Body Radiation

From 1850 to 1900, scientists were trying to find the solution to the problem called as "Ultra Violet Catastrophe". It is also called as "Rayleigh-Jeans catastrophe". As per the classical physics, an ideal black body at thermal equilibrium would emit an unbounded quantity of energy as wavelength decreased into the ultraviolet range! Spectral radiance is directly proportional to the frequency. So if frequency is infinte, spectral radiance will also be infinte. and that's the error in "Ultra Violet Catastrophe". Below you can play with the black body radiation spectrum! Max Planck In 1900, Max Planck derived the correct form for the intensity spectral distribution function by assuming that radiation can be emitted or absorbed only in discrete packets, called quanta, of energy. Albert Einstein (in 1905) confirmed Planc's theory stating that Planck's quanta were physical particles - photons! Photoelectric effect animation can be found at

Rutherford Model

Thomson in 1898 proposed that an atom consists of uniform sphere of positive electricity with electrons embedded into it. This model in which mass of the atom is considered to be evenly spread over the atom was proved wrong by Rutherford's famous alpha-particle scattering experiment in 1909. Rutherford concluded that atom is made of a tiny positively charged nucleus, at its centre with electrons revolving around it in circular orbits. Rutherford model, which resembles the solar system, was no doubt an improvement over Thomson model but it could not account for the stability of the atom i.e., why the electron does not fall into the nucleus. Further, it was also silent about the electronic structure of atoms i.e., about the distribution and relative energies of electrons around the nucleus. The difficulties of the Rutherford model were overcome by Niels Bohr in 1913 in his model of the hydrogen atom. Bohr postulated that electron moves around the nucleus in circular orbits. Only certain orbits can exist and each orbit corresponds to a specific energy. Bohr calculated the energy of electron in various orbits and for each orbit predicted the distance between the electron and nucleus.

Bohr Model

Bohr model, though offering a satisfactory model for explaining the spectra of the hydrogen atom, could not explain the spectra of multi-electron atoms. The reason for this was soon discovered. In Bohr model, an electron is regarded as a charged particle moving in a well defined circular orbit about the nucleus. The wave character of the electron is ignored in Bohr's theory. An orbit is a clearly defined path and this path can completely be defined only if both the exact position and the exact velocity of the electron at the same time are known. This is not possible according to the Heisenberg uncertainty principle. Bohr model of the hydrogen atom, therefore, not only ignores the dual behaviour of electron but also contradicts Heisenberg uncertainty principle.

Schrödinger equation

Erwin Schrödinger, in 1926, proposed an equation called Schrödinger equation to describe the electron distributions in space and the allowed energy levels in atoms. This equation incorporates de Broglie's concept of wave-particle duality and is consistent with Heisenberg uncertainty principle. When Schrödinger equation is solved for the electron in a hydrogen atom, the solution gives the possible energy states the electron can occupy [and the corresponding wave function(s) (ψ) (which in fact are the mathematical functions) of the electron associated with each energy state]. These quantized energy states and corresponding wave functions which are characterized by a set of three quantum numbers (principal quantum number n, azimuthal quantum number l and magnetic quantum number ml) arise as a natural consequence in the solution of the Schrödinger equation. The restrictions on the values of these three quantum numbers also come naturally from this solution. The quantum mechanical model of the hydrogen atom successfully predicts all aspects of the hydrogen atom spectrum including some phenomena that could not be explained by the Bohr model.

Quantum mechanical model

According to the quantum mechanical model of the atom, the electron distribution of an atom containing a number of electrons is divided into shells. The shells, in turn, are thought to consist of one or more subshells and subshells are assumed to be composed of one or more orbitals, which the electrons occupy. While for hydrogen and hydrogen like systems (such as He+, Li2+ etc.) all the orbitals within a given shell have same energy, the energy of the orbitals in a multi-electron atom depends upon the values of n and l: The lower the value of (n + l ) for an orbital, the lower is its energy. If two orbitals have the same (n + l ) value, the orbital with lower value of n has the lower energy. In an atom many such orbitals are possible and electrons are filled in those orbitals in order of increasing energy in accordance with Pauli exclusion principle (no two electrons in an atom can have the same set of four quantum numbers) and Hund's rule of maximum multiplicity (pairing of electrons in the orbitals belonging to the same subshell does not take place until each orbital belonging to that subshell has got one electron each, i.e., is singly occupied). This forms the basis of the electronic structure of atoms.

Unit 3 - Classification of Elements and Periodicity in Properties

Periodic Law

Mendeleev's Periodic Table

Mendeleev's Periodic Table was based on atomic masses.

Modern Periodic Table

Modern Periodic Table arranges the elements in the order of their atomic numbers in seven horizontal rows (periods) and eighteen vertical columns (groups or families). Atomic numbers in a period are consecutive, whereas in a group they increase in a pattern. Elements of the same group have similar valence shell electronic configuration and, therefore, exhibit similar chemical properties. However, the elements of the same period have incrementally increasing number of electrons from left to right, and, therefore, have different valencies. Four types of elements can be recognized in the periodic table on the basis of their electronic configurations. These are s-block, p-block, d-block and f-block elements. Hydrogen with one electron in the 1s orbital occupies a unique position in the periodic table. Metals comprise more than seventy eight per cent of the known elements. Non-metals, which are located at the top of the periodic table, are less than twenty in number. Elements which lie at the border line between metals and non-metals (e.g., Si, Ge, As) are called metalloids or semi-metals. Metallic character increases with increasing atomic number in a group whereas decreases from left to right in a period. The physical and chemical properties of elements vary periodically with their atomic numbers.

Periodic trends

Periodic trends are observed in atomic sizes, ionization enthalpies, electron gain enthalpies, electronegativity and valence. The atomic radii decrease while going from left to right in a period and increase with atomic number in a group. Ionization enthalpies generally increase across a period and decrease down a group. Electronegativity also shows a similar trend. Electron gain enthalpies, in general, become more negative across a period and less negative down a group. There is some periodicity in valence, for example, among representative elements, the valence is either equal to the number of electrons in the outermost orbitals or eight minus this number.

Chemical reactivity

Chemical reactivity is highest at the two extremes of a period and is lowest in the centre. The reactivity on the left extreme of a period is because of the ease of electron loss (or low ionization enthalpy). Highly reactive elements do not occur in nature in free state; they usually occur in the combined form. Oxides formed of the elements on the left are basic and of the elements on the right are acidic in nature. Oxides of elements in the centre are amphoteric or neutral.

Unit 4 - Chemical Bonding and Molecular Structure

The attractive force which holds various constituents (atoms, ions, etc.) together in different chemical species is called a chemical bond

KÖssel-Lewis Approach to Chemical Bonding

Kössel's first insight into the mechanism of formation of electropositive and electronegative ions related the process to the attainment of noble gas configurations by the respective ions. Electrostatic attraction between ions is the cause for their stability. This gives the concept of electrovalency. Lewis postulated that atoms achieve the stable octet when they are linked by chemical bonds. In the formation of a molecule, only the outer shell electrons take part in chemical combination and they are known as valence electrons. Sodium (Na) has 1 valence electron and Chlorine(Cl) has 7 valence electrons. So Sodium gives away 1 electron to chlorine and thus NaCl compound is formed. The bond formed, as a result of the electrostatic attraction between the positive (Na) and negative (Cl) ions is called as Electrovalent Bond. Kössel and Lewis in 1916 developed an important theory of chemical combination between atoms known as electronic theory of chemical bonding. According to this, atoms can combine either by transfer of valence electrons from one atom to another (gaining or losing) or by sharing of valence electrons in order to have an octet in their valence shells. This is known as octet rule.

covalent bonding

The first description of covalent bonding was provided by Lewis in terms of the sharing of electron pairs between atoms and he related the process to the attainment of noble gas configurations by reacting atoms as a result of sharing of electrons. The Lewis dot symbols show the number of valence electrons of the atoms of a given element and Lewis dot structures show pictorial representations of bonding in molecules. An ionic compound is pictured as a three-dimensional aggregation of positive and negative ions in an ordered arrangement called the crystal lattice. In a crystalline solid there is a charge balance between the positive and negative ions. The crystal lattice is stabilized by the enthalpy of lattice formation.

Lewis dot structure

While a single covalent bond is formed by sharing of an electron pair between two atoms, multiple bonds result from the sharing of two or three electron pairs. Some bonded atoms have additional pairs of electrons not involved in bonding. These are called lone-pairs of electrons. A Lewis dot structure shows the arrangement of bonded pairs and lone pairs around each atom in a molecule. Important parameters, associated with chemical bonds, like: bond length, bond angle, bond enthalpy, bond order and bond polarity have significant effect on the properties of compounds. A number of molecules and polyatomic ions cannot be described accurately by a single Lewis structure and a number of descriptions (representations) based on the same skeletal structure are written and these taken together represent the molecule or ion. This is a very important and extremely useful concept called resonance. The contributing structures or canonical forms taken together constitute the resonance hybrid which represents the molecule or ion.

VSEPR model

The VSEPR model used for predicting the geometrical shapes of molecules is based on the assumption that electron pairs repel each other and, therefore, tend to remain as far apart as possible. According to this model, molecular geometry is determined by repulsions between lone pairs and lone pairs; lone pairs and bonding pairs and bonding pairs and bonding pairs. The order of these repulsions being : lplp>lpbp>bpbplp-lp > lp-bp > bp-bp


The valence bond (VB) approach to covalent bonding is basically concerned with the energetics of covalent bond formation about which the Lewis and VSEPR models are silent. Basically the VB theory discusses bond formation in terms of overlap of orbitals. For example the formation of the H2 molecule from two hydrogen atoms involves the overlap of the 1s orbitals of the two H atoms which are singly occupied. It is seen that the potential energy of the system gets lowered as the two H atoms come near to each other. At the equilibrium inter-nuclear distance (bond distance) the energy touches a minimum. Any attempt to bring the nuclei still closer results in a sudden increase in energy and consequent destabilization of the molecule. Because of orbital overlap the electron density between the nuclei increases which helps in bringing them closer. It is however seen that the actual bond enthalpy and bond length values are not obtained by overlap alone and other variables have to be taken into account. For explaining the characteristic shapes of polyatomic molecules Pauling introduced the concept of hybridisation of atomic orbitals. sp, sp2, sp3 hybridizations of atomic orbitals of Be, B, C, N and O are used to explain the formation and geometrical shapes of molecules like BeCl2BeCl_2 , BCl3BCl_3 ,CH4CH_4 ,NH3NH_3 and H2OH_2O. They also explain the formation of multiple bonds in molecules like C2H2C_2H_2 and C2H4C_2H_4.

The molecular orbital (MO) theory

The molecular orbital (MO) theory describes bonding in terms of the combination and arrangment of atomic orbitals to form molecular orbitals that are associated with the molecule as a whole. The number of molecular orbitals are always equal to the number of atomic orbitals from which they are formed. Bonding molecular orbitals increase electron density between the nuclei and are lower in energy than the individual atomic orbitals. Antibonding molecular orbitals have a region of zero electron density between the nuclei and have more energy than the individual atomic orbitals. The electronic configuration of the molecules is written by filling electrons in the molecular orbitals in the order of increasing energy levels. As in the case of atoms, the Pauli exclusion principle and Hund's rule are applicable for the filling of molecular orbitals. Molecules are said to be stable if the number of elctrons in bonding molecular orbitals is greater than that in antibonding molecular orbitals. Hydrogen bond is formed when a hydrogen atom finds itself between two highly electronegative atoms such as F, O and N. It may be intermolecular (existing between two or more molecules of the same or different substances) or intramolecular (present within the same molecule). Hydrogen bonds have a powerful effect on the structure and properties of many compounds.

Unit 5 - Thermodynamics

Thermodynamics deals with energy changes in chemical or physical processes and enables us to study these changes quantitatively and to make useful predictions. For these purposes, we divide the universe into the system and the surroundings. Chemical or physical processes lead to evolution or absorption of heat (q), part of which may be converted into work (w). These quantities are related through the first law of thermodynamics via
ΔU=q+wΔU\Delta U = q + w \cdot \Delta U
change in internal energy, depends on initial and final states only and is a state function, whereas q and w depend on the path and are not the state functions. We follow sign conventions of q and w by giving the positive sign to these quantities when these are added to the system. We can measure the transfer of heat from one system to another which causes the change in temperature. The magnitude of rise in temperature depends on the heat capacity (C) of a substance. Therefore, heat absorbed or evolved is
q=CΔTq = C \Delta T
Work can be measured by
w=pexΔVw = - p_{ex} \Delta V
in case of expansion of gases. Under reversible process, we can put
pex=pp_{ex} = p
for infinitesimal changes in the volume making
wrev=pdVw_{rev} = - p dV
In this condition, we can use gas equation,
pV=nRTpV = nRT


At constant volume, w = 0, then ΔU=qv ΔU = q_v , heat transfer at constant volume. But in study of chemical reactions, we usually have constant pressure. We define another state function enthalpy. Enthalpy change, ΔH=ΔU+ΔngRTΔH = ΔU + Δn_g RT , can be found directly from the heat changes at constant pressure, ΔH=qpΔH = q_p . There are varieties of enthalpy changes. Changes of phase such as melting, vaporization and sublimation usually occur at constant temperature and can be characterized by enthalpy changes which are always positive. Enthalpy of formation, combustion and other enthalpy changes can be calculated using Hess's law. Enthalpy change for chemical reactions can be determined by
ΔrH=f(aiΔfHproducts)i(biΔfHreactions)\Delta_rH = \sum_f(a_i \Delta_fH_{products}) - \sum_i(b_i \Delta_f H_{reactions})
and in gaseous state by
ΔrHo=Σbond enthalpies of the reactantsΣbond enthalpies of the productsΔrH^o = Σ \text{bond enthalpies of the reactants} - Σ \text{bond enthalpies of the products}
First law of thermodynamics does not guide us about the direction of chemical reactions i.e., what is the driving force of a chemical reaction. For isolated systems,
ΔU=0ΔU = 0
. We define another state function, S, entropy for this purpose. Entropy is a measure of disorder or randomness. For a spontaneous change, total entropy change is positive. Therefore, for an isolated system, ΔU=0ΔU = 0, ΔS>0ΔS > 0 , so entropy change distinguishes a spontaneous change, while energy change does not. Entropy changes can be measured by the equation
ΔS=qrevT\Delta S = \frac {q_{rev}} {T}
for a reversible process.
qrevT\frac {q_{rev}} {T}
is independent of path. Chemical reactions are generally carried at constant pressure, so we define another state function Gibbs energy, G, which is related to entropy and enthalpy changes of the system by the equation
ΔrG=ΔrHTΔrS\Delta_r G = \Delta_r H - T \Delta_r S
For a spontaneous change,ΔGsys<0\Delta G_{sys} < 0and at equilibrium, ΔGsys=0\Delta G_{sys} = 0Standard Gibbs energy change is related to equilibrium constant by
ΔrGo=RTlnKΔ_rG^o = -RT ln K
K can be calculated from this equation, if we know ΔrGoΔ_rG^o which can be found from
ΔrGo=ΔrHoTΔrSoΔ_rG^o = Δ_rH^o - TΔ_rS^o
Temperature is an important factor in the equation. Many reactions which are non-spontaneous at low temperature, are made spontaneous at high temperature for systems having positive entropy of reaction.

Unit 6 - Equilibrium

When the number of molecules leaving the liquid to vapour equals the number of molecules returning to the liquid from vapour, equilibrium is said to be attained and is dynamic in nature. Equilibrium can be established for both physical and chemical processes and at this stage rate of forward and reverse reactions are equal. Equilibrium constant, Kc is expressed as the concentration of products divided by reactants, each term raised to the stoichiometric coefficient.
Kc=[C]c[D]d[A]a[B]bK_c = \frac{[C]^c [D]^d}{[A]^a [B]^b}
Equilibrium constant has constant value at a fixed temperature and at this stage all the macroscopic properties such as concentration, pressure, etc. become constant. For a gaseous reaction equilibrium constant is expressed as Kp and is written by replacing concentration terms by partial pressures in Kc expression. The direction of reaction can be predicted by reaction quotient Qc which is equal to Kc at equilibrium. Le Chatelier's principle states that the change in any factor such as temperature, pressure, concentration, etc. will cause the equilibrium to shift in such a direction so as to reduce or counteract the effect of the change. It can be used to study the effect of various factors such as temperature, concentration, pressure, catalyst and inert gases on the direction of equilibrium and to control the yield of products by controlling these factors. Catalyst does not effect the equilibrium composition of a reaction mixture but increases the rate of chemical reaction by making available a new lower energy pathway for conversion of reactants to products and vice-versa. All substances that conduct electricity in aqueous solutions are called electrolytes. Acids, bases and salts are electrolytes and the conduction of electricity by their aqueous solutions is due to anions and cations produced by the dissociation or ionization of electrolytes in aqueous solution. The strong electrolytes are completely dissociated. In weak electrolytes there is equilibrium between the ions and the unionized electrolyte molecules. According to Arrhenius, acids give hydrogen ions while bases produce hydroxyl ions in their aqueous solutions. Brönsted-Lowry on the other hand, defined an acid as a proton donor and a base as a proton acceptor. When a Brönsted-Lowry acid reacts with a base, it produces its conjugate base and a conjugate acid corresponding to the base with which it reacts. Thus a conjugate pair of acid-base differs only by one proton. Lewis further generalised the definition of an acid as an electron pair acceptor and a base as an electron pair donor. The expressions for ionization (equilibrium) constants of weak acids (Ka) and weak bases (Kb) are developed using Arrhenius definition. The degree of ionization and its dependence on concentration and common ion are discussed. The pH scale
(pH=log[H+])(pH = –log[H+])
for the hydrogen ion concentration (activity) has been introduced and extended to other quantities(pOH=log[OH]);(pOH = – log[OH–]);pKa=log[Ka];pKa = –log[Ka];pKb=log[Kb];pKb = –log[Kb];and pKw=log[Kw];pKw = –log[Kw];The ionization of water has been considered and we note that the equation:pH+pOH=pKwpH + pOH = pKw is always satisfied. The salts of strong acid and weak base, weak acid and strong base, and weak acid and weak base undergo hydrolysis in aqueous solution. The definition of buffer solutions, and their importance are discussed briefly. The solubility equilibrium of sparingly soluble salts is discussed and the equilibrium constant is introduced as solubility product constant (Ksp). Its relationship with solubility of the salt is established. The conditions of precipitation of the salt from their solutions or their dissolution in water are worked out. The role of common ion and the solubility of sparingly soluble salts is also discussed.


Redox reactions form an important class of reactions in which oxidation and reduction occur simultaneously. Three tier conceptualisation viz, classical, electronic and oxidation number, which is usually available in the texts, has been presented in detail. Oxidation, reduction, oxidising agent (oxidant) and reducing agent (reductant) have been viewed according to each conceptualisation. Oxidation numbers are assigned in accordance with a consistent set of rules. Oxidation number and ion-electron method both are useful means in writing equations for the redox reactions. Redox reactions are classified into four categories: combination, decomposition displacement and disproportionation reactions. The concept of redox couple and electrode processes is introduced here. The redox reactions find wide applications in the study of electrode processes and cells.

Unit 8 - Organic Chemistry - Some Basic Principles and Techniques

Covalent bonding and reactions

The nature of the covalent bonding in organic compounds can be described in terms of orbitals hybridisation concept, according to which carbon can have sp3, sp2 and sp hybridised orbitals. The sp3, sp2 and sp hybridised carbons are found in compounds like methane, ethene and ethyne respectively. The tetrahedral shape of methane, planar shape of ethene and linear shape of ethyne can be understood on the basis of this concept. A sp3 hybrid orbital can overlap with 1s orbital of hydrogen to give a carbon - hydrogen (C-H) single bond (sigma, σ bond). Overlap of a sp2 orbital of one carbon with sp2 orbital of another results in the formation of a carbon-carbon σ bond. The unhybridised p orbitals on two adjacent carbons can undergo lateral (sideby- side) overlap to give a pi (π) bond. Organic compounds can be represented by various structural formulas. The three dimensional representation of organic compounds on paper can be drawn by wedge and dash formula.

IUPAC Naming

Organic compounds can be classified on the basis of their structure or the functional groups they contain. A functional group is an atom or group of atoms bonded together in a unique fashion and which determines the physical and chemical properties of the compounds. The naming of the organic compounds is carried out by following a set of rules laid down by the International Union of Pure and Applied Chemistry (IU PAC ). In IUPAC nomenclature, the names are correlated with the structure in such a way that the reader can deduce the structure from the name.

Organic Reactions

Organic reaction mechanism concepts are based on the structure of the substrate molecule, fission of a covalent bond, the attacking reagents, the electron displacement effects and the conditions of the reaction. These organic reactions involve breaking and making of covalent bonds. A covalent bond may be cleaved in heterolytic or homolytic fashion. A heterolytic cleavage yields carbocations or carbanions, while a homolytic cleavage gives free radicals as reactive intermediate. Reactions proceeding through heterolytic cleavage involve the complimentary pairs of reactive species. These are electron pair donor known as nucleophile and an electron pair acceptor known as electrophile. The inductive, resonance, electromeric and hyperconjugation effects may help in the polarisation of a bond making certain carbon atom or other atom positions as places of low or high electron densities. Organic reactions can be broadly classified into following types; substitution, addition, elimination and rearrangement reactions.

Organic Compound Analysis

Purification, qualitative and quantitative analysis of organic compounds are carried out for determining their structures. The methods of purification namely : sublimation, distillation and differential extraction are based on the difference in one or more physical properties. Chromatography is a useful technique of separation, identification and purification of compounds. It is classified into two categories : adsorption and partition chromatography. Adsorption chromatography is based on differential adsorption of various components of a mixture on an adsorbent. Partition chromatography involves continuous partitioning of the components of a mixture between stationary and mobile phases. After getting the compound in a pure form, its qualitative analysis is carried out for detection of elements present in it. Nitrogen, sulphur, halogens and phosphorus are detected by Lassaigne's test. Carbon and hydrogen are estimated by determining the amounts of carbon dioxide and water produced. Nitrogen is estimated by Dumas or Kjeldahl's method and halogens by Carius method. Sulphur and phosphorus are estimated by oxidising them to sulphuric and phosphoric acids respectively. The percentage of oxygen is usually determined by difference between the total percentage (100) and the sum of percentages of all other elements present.

Unit 9 - Hydrocarbons

Hydrocarbons are those compounds formed from carbon and hydrogen!

Hydrocarbon Classification

Hydrocarbons can be classified into 3 categories.
  • Saturated - Alkanes (single bond)
  • Unsaturated - Alkenes (double bond) and Alkynes (triple bond)
  • Aromatic - Cyclic Compounds


Alkanes are saturated Hydrocarbons as single bonds exist between carbon atoms. Example: Methane (CH4), Ethane (C2H6), Propane (C3H8). The general formula for alkanes is CnH2n+2, where "n" represents the number of carbon atoms in the molecule. Alkanes are relatively stable and non-reactive compared to alkenes and alkynes. They are commonly found in natural gas and petroleum.


Alkenes (double bond) and Alkynes (triple bond) are unsaturated hydrocarbons as double or triple bond may exist between carbon atoms. The general formula for alkenes is CnH2n, where "n" represents the number of carbon atoms in the molecule. Alkenes are more reactive than alkanes due to the presence of double bonds, making them susceptible to addition reactions Example: Ethene (C2H4), Propene (C3H6), Butene (C4H8).


The general formula for alkynes is CnH2n-2, where "n" represents the number of carbon atoms in the molecule. Alkynes are even more reactive than alkenes due to the presence of triple bonds, making them highly susceptible to addition reactions. Example: Ethyne (C2H2), Propyne (C3H4), Butyne (C4H6).


Aromatic hydrocarbons are a class of organic compounds that contain a unique and stable ring of carbon atoms, often referred to as an "aromatic ring" or a "benzene ring." The most well-known and simplest aromatic hydrocarbon is benzene (C6H6), which consists of a hexagonal ring of six carbon atoms, each bonded to a hydrogen atom. Key Characteristics
  • Aromaticity: The term "aromatic" refers to the characteristic aroma or smell of some of the early-discovered compounds in this class. However, many aromatic compounds are odorless. Aromaticity is a specific property of these compounds, associated with the delocalization of electrons within the ring structure, making them highly stable.
  • Ring Structure: Aromatic hydrocarbons have one or more closed rings of carbon atoms. Benzene is the simplest example, but larger and more complex aromatic rings are also common.
  • Stability: Aromatic compounds are exceptionally stable due to the resonance structure and electron delocalization within the ring. This stability makes them less reactive than other hydrocarbons, like alkenes and alkynes.
  • Unsaturated: Aromatic hydrocarbons are unsaturated compounds, meaning they contain double bonds (pi bonds) within the ring structure. However, they are not as reactive as typical alkenes because of their stability.
  • Planar Structure: The carbon atoms in the aromatic ring are arranged in a planar, or flat, structure. This planarity is important for maintaining the aromaticity.
  • Substitution Reactions: Aromatic compounds are often subjected to substitution reactions, where one atom or group replaces another on the ring. Common substituents include alkyl groups (e.g., methyl, ethyl) and functional groups (e.g., halogens, nitro groups).
Common Examples
  • Benzene: The simplest aromatic hydrocarbon.
  • Toluene: A benzene ring with a methyl group attached.
  • Phenol: Benzene with a hydroxyl (OH) group attached.
  • Naphthalene: A fused aromatic ring system consisting of two benzene rings.
  • Aniline: Benzene with an amino (NH2) group attached.
  • TNT (Trinitrotoluene): An explosive compound containing both aromatic and nitro groups.

Aromatic hydrocarbons have a wide range of applications in industry and are found in many natural and synthetic compounds, including dyes, pharmaceuticals, and plastics. The stability and unique electronic properties of aromatic rings make them a focus of study in both organic chemistry and materials science.