Organic functional groups

Environmental chemistry

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Syllabus Notes Assignments

Organic functional groups are specific groups of atoms within molecules that are responsible for the characteristic chemical reactions of those molecules. These functional groups play a crucial role in organic chemistry because they determine the properties and reactivity of the compounds in which they are found. Each functional group has a specific set of atoms arranged in a particular pattern, and this arrangement gives the group its unique chemical behaviour. Functional groups are used to classify organic compounds and predict their reactivity.

For example, compounds with hydroxyl groups (-OH) are typically alcohols and have certain properties such as solubility in water and the ability to form hydrogen bonds. Similarly, carboxyl groups (-COOH) are found in carboxylic acids, which are known for their acidity and ability to participate in esterification reactions.

Functional Group	Structure	Properties	Example
Hydroxyl Group	-OH	Polar, forms hydrogen bonds	Ethanol (CH₃CH₂OH)
Aldehyde	-CHO	Polar, reactive	Formaldehyde (HCHO)
Ketone	>C=O	Polar, less reactive than aldehydes	Acetone (CH₃COCH₃)
Carboxyl Group	-COOH	Highly polar, acidic	Acetic acid (CH₃COOH)
Amino Group	-NH₂	Basic, forms hydrogen bonds	Aniline (C₆H₅NH₂)
Thiol Group	-SH	More nucleophilic than hydroxyl group	Ethanethiol (CH₃CH₂SH)
Phosphate Group	-PO₄	Highly polar, acidic	ATP
Ether Group	R-O-R	Relatively non-polar	Diethyl ether (CH₃CH₂OCH₂CH₃)
Ester Group	-COOR	Polar, pleasant odours	Ethyl acetate (CH₃COOCH₂CH₃)
Amide Group	-CONH₂	Polar, forms hydrogen bonds	Acetamide (CH₃CONH₂)
Nitrile Group	-C≡N	Polar, high dipole moment	Acetonitrile (CH₃CN)

This table provides a concise overview of common organic functional groups, highlighting their structures, properties & examples.

Fundamental Chemistry: Elements and Chemical bonding

Elements

The chemical elements, the fundamental substances, are made from atoms. And atom is the smallest particle of an element which can take part in a chemical reaction. The main components of an atom are protons, which are positively charged, electrons, which are negatively charged and neutrons, which are uncharged. The all atoms of an element have the same atomic number (Z), which is equal to the number of protons in an atom. In an atom the number of protons and electrons is equal and therefore an atom is electrically neutral. The elements may exist in mono-atomic or polyatomic forms. For example in air, argon exists as a single atom, Ar, and nitrogen (N) exists as N₂.

A given element may exist in more than one polyatomic (molecular) forms. These forms are known as allotropes. For example, in atmosphere oxygen is found both as dioxygen, O₂, and ozone,O₃.

The elements and their compounds are building block of every material found in the universe. The relative abundance of the elements by mass in whole Earth is: 35% iron, 30% oxygen, 15% silicon. On the other hand Earth’s crust has 40% oxygen, 28% silicon, 8% aluminum 6% iron, etc.

Classification of Elements on the Basis of Their Properties

Metals

These elements have low ionization energies, low electron affinities and low electronegativities, and form cations. Out of 118 elements known at present, about 81 are metals. Metals are further classified as follows.

1. Alkali Metals

Alkali metals, viz., Li, Na, K, Rb, Cs are very reactive and react with O₂, S, and other nonmetals to form compounds. That is why these are never found in nature in native state. There outermost electronic configuration is [noble gas], ns¹. These elements are easily ionized to form cations. Alkali form strongly alkaline oxides, e.g., Na₂O and hydroxides, NaOH, all of which react with water to produce hydroxide ions, OH- , which make the aqueous solution alkaline.

2. Alkaline Earth Metals

The elements belonging to Group 2 are also reactive metals and commonly known as alkaline earth metals. The alkaline earth metals have the outermost electronic

configuration, ns². Thus these metals form di positive ions. The oxides and hydroxides and carbonates of Mg, Ca, Sr and Ba are alkaline and their aqueous solutions are alkaline. The carbonates of these metals are found in Earth crust. The last member of the group, Radium, is radioactive and has several applications.

3. Transition metals

There are four series of these elements. The first series includes Sc, Ti, V, Cr, Mn, Fe, Co, Ni and Cu. Chemistry of these elements is of great environmental significance. These elements exhibit the variable oxidation states. For example Mn shows all the oxidation states starting from -1 to +7. These elements form large number of complexes and also act as catalysts in all types of chemical systems.

4. Rare earth elements

The elements of lanthanide group are called the rare earths, which comprise the fourteen elements from Ce(z = 580 to Lu (Z = 71), but sometimes La, Sc and Y are also included. The prominent member of actinide series are U and Pu, which are used in production of nuclear energy.

5. Heavy metals

Because of their high relative atomic masses, As, Be, Cd, Pb, Mn, Hg, Ni and Se are called heavy metals. These elements concern us because of occupational or residual exposure. They persist in nature and can cause damage or death in animals, humans and plants, even at a very low concentration (1 or 2 microns in some cases). Industrial processes release these into air and water. Since heavy metals have a property to accumulate in the selected body organs, such as brain and liver, long term exposure may result in slowly progressing physical, muscular and neurological degenerative processes that mimic Alzheimer's diseases, Parkinson’s diseases, Muscular dystrophy and multiple sclerosis, Allergies are not uncommon. In all water analyses, measurement of heavy metals is necessary.

6. Base metals

The non-ferrous metals, excluding precious metals, are called base metals, e.g., iron, steel, aluminium, tin, tungsten, molybdenum, tantalum, cobalt, bismuth, cadmium, titanium, zirconium, antimony, manganese, beryllium, chromium, germanium, vanadium, gallium, hafnium, indium, niobium, rhenium, thallium, Copper

and Lead.

7. Ferrous and Nonferrous metals

Ferrous metals contain iron, for example carbon steel, stainless steel (both alloys; mixtures of metals) and wrought iron, Non-ferrous metals are metals that do not

contain iron, for example aluminium, brass, copper and titanium brass. Al, Be, Cu, Pb, Mg, Ni, Sn, Zn, and precious metals.

8. Metalloid

There is no rigorous definition of metalloids, but the elements having the properties of both metals and nonmetals are called metalloids. The metalloids often form

amphoteric oxides (B, Si, Ge, As and Sb and often behave as semiconductor (B, Si, Ge, As).

Non-Metals

About 22 elements, which are members of p-block, behave as non-metals. They are usually poor conductor. They are found as gases, liquids or solids.

1. Carbon Family

The elements of Group 14, C, Si, Ge, Sn, Pb constitute the carbon family. Carbon is nonmetal, Si and Ge are semimetal, and tin and lead are metals. The chemistry of carbon is vast and studied separately as organic chemistry. This is due to the property of catenation, i.e., ability to carbon chains. The carbon cycle is an important cycle in environment.

2. Nitrogen family

Nitrogen family consists of group 15 elements, viz., N, P, As, Sb, Bi, etc. The first two elements, N and P are very important and their biogeochemical cycles play a very important role in the chemistry of the environment.

3. Oxygen family

Group 16 contain 6 elements, viz., O, S, Se, etc. O and S are nonmetals whereas Se is a metalloid. The importance of oxygen and sulfur cycles in environment shall be presented in a subsequent module.

4. Halogen

The elements of group 17 are known as halogen, viz., F, Cr, Br, I. The halogens have a very strong tendency to pick up one electron to acquire the stable noble gas configuration. These are among the most reactive elements. Chlorine easily undergoes photochemical dissociation and forms Cl atom/free radical, which is highly reactive and responsible for stratospheric ozone depletion. Chlorine is widely used for disinfection drinking water.

5. Noble/Inert/Rare Gases

Group 18 elements such as He, Ne, Ar, Kr, Xe and Rn are known as noble, rare or inert gases. Due to their stable ns², np⁶electronic configuration, they have very little tendency to form compounds with other elements. For this reason they exist in mono-atomic form, e. g., He, Ar, Ne, etc. In air, argon is the third most significant gas (0.9%) after N2 and O2. Radon is a radioactive element having serious environmental concerns as discussed another module.

Chemical Bonding

Chemical bonding is the process by which atoms or molecules combine to form more complex structures, held together by attractive forces. These forces arise from the interactions between electrons and nuclei of different atoms. Chemical bonds can be categorized into several types, including ionic bonds, where electrons are transferred between atoms, covalent bonds, where electrons are shared between atoms, and metallic bonds, which involve a 'sea' of shared electrons among a lattice of metal atoms. The nature of the chemical bond determines the physical and chemical properties of the resulting compounds.

1. Ionic Bonding

Ionic bonds are formed through the transfer of electrons from one atom to another, resulting in the formation of ions. This typically occurs between metals and non-metals. Metals lose electrons to become positively charged cations, while non-metals gain those electrons to become negatively charged anions. The electrostatic attraction between oppositely charged ions creates the ionic bond.

Example Reaction: Formation of Sodium Chloride (NaCl)

Sodium (Na) loses one electron to form a sodium ion (Na⁺): $Na \to {Na}^{+} + {????}^{-}$
Chlorine (Cl) gains one electron to form a chloride ion (Cl⁻): $Cl + {????}^{-} \to {Cl}^{-}$
The Na⁺ and Cl⁻ ions attract each other to form sodium chloride: ${Na}^{+} + {Cl}^{-} \to NaCl$

2. Covalent Bonding

Covalent bonds involve the sharing of electrons between atoms, typically between non-metals. The shared electrons allow each atom to achieve a stable electron configuration, similar to that of noble gases.

Example Reaction: Formation of Water (H₂O)

Hydrogen (H) atoms share electrons with an oxygen (O) atom to form water molecules: $2 H_{2} + O_{2} \to 2 H_{2} O$ In a water molecule, each hydrogen atom shares one electron with the oxygen atom, forming two covalent bonds: $H - O - H$

3. Metallic Bonding

Metallic bonds occur between metal atoms. In metallic bonding, atoms in a metal lattice share a 'sea' of delocalized electrons, which are free to move throughout the structure. This electron sharing gives metals their characteristic properties, such as conductivity, malleability, and ductility.

Example: Metallic Bonding in Copper (Cu)

Copper atoms in a solid state share a pool of electrons that move freely throughout the metal lattice. This delocalization of electrons binds the metal ions together, but it is not represented by a simple reaction equation like ionic or covalent bonding.

Other Types of Bonding

Hydrogen Bonding

Hydrogen bonds are weak bonds that occur when a hydrogen atom, covalently bonded to a highly electronegative atom (such as oxygen or nitrogen), is attracted to another electronegative atom. This type of bond is crucial in biological molecules like DNA and proteins.

Example: Hydrogen Bonding in Water (H₂O)

The hydrogen atoms of one water molecule are attracted to the oxygen atom of another water molecule, forming hydrogen bonds: $H_{2} O \dots H_{2} O$

Van der Waals Forces

Van der Waals forces are weak intermolecular forces that arise from temporary dipoles in molecules. They include dispersion forces and dipole-dipole interactions.

Example: Interaction Between Noble Gas Atoms

Inert gases like argon (Ar) exhibit Van der Waals forces due to temporary dipoles: $Ar \dots Ar$

Chemical reactions and equations

Chemical reactions are processes in which substances, known as reactants, undergo chemical changes to form new substances, called products. These reactions involve the breaking and forming of chemical bonds, resulting in the transformation of substances with different chemical properties. The representation of these chemical reactions using symbols and formulas is known as chemical equations.

Types of Chemical Reactions

Combination (Synthesis) Reactions
- Two or more reactants combine to form a single product.
- Example: $2 H_{2} + O_{2} \to 2 H_{2} O$
Decomposition Reactions
- A single compound breaks down into two or more simpler substances.
- Example: $2 H_{2} O_{2} \to 2 H_{2} O + O_{2}$
Single Displacement (Replacement) Reactions
- One element replaces another element in a compound.
- Example: $Zn + 2 HCl \to {ZnCl}_{2} + H_{2}$
Double Displacement (Metathesis) Reactions
- The ions of two compounds exchange places in an aqueous solution to form two new compounds.
- Example: ${AgNO}_{3} + NaCl \to AgCl + {NaNO}_{3}$
Combustion Reactions
- A substance combines with oxygen, releasing energy in the form of light and heat.
- Example: ${CH}_{4} + 2 O_{2} \to {CO}_{2} + 2 H_{2} O$
Redox (Oxidation-Reduction) Reactions
- Involves the transfer of electrons between two species, leading to changes in their oxidation states.
- Example: $2Na + {Cl}_{2} \to 2 NaCl$

Chemical Equations

Chemical equations use chemical symbols and formulas to represent reactants and products in a chemical reaction. A balanced chemical equation has the same number of atoms of each element on both sides, ensuring the law of conservation of mass is obeyed.

Parts of a Chemical Equation:

Reactants: Substances that undergo change (left side of the equation).
Products: New substances formed (right side of the equation).
Coefficients: Numbers placed before formulas to balance the equation.
Arrow (→): Indicates the direction of the reaction, from reactants to products.

Example of a Balanced Equation: $N_{2} + 3 H_{2} \to 2 {NH}_{3}$

Steps to Balance a Chemical Equation:

Write the unbalanced equation. $Fe + O_{2} \to {Fe}_{2} O_{3}$
Count the number of atoms of each element on both sides.
- Reactants: 1 Fe, 2 O
- Products: 2 Fe, 3 O
Use coefficients to balance each element.
- Balance Fe: $4 Fe + 3 O_{2} \to 2 {Fe}_{2} O_{3}$
Check to ensure the equation is balanced.
- Reactants: 4 Fe, 6 O
- Products: 4 Fe, 6 O

Types of Equations

Word Equations
- Describes the reactants and products in words.
- Example: $Iron + Oxygen \to Iron(III) oxide$
Formula Equations
- Uses chemical symbols and formulas.
- Example: $4 Fe + 3 O_{2} \to 2 {Fe}_{2} O_{3}$
Ionic Equations
- Shows the ions involved in the reaction, often used for reactions in aqueous solutions.
- Example: ${Ag}^{+} + {Cl}^{-} \to AgCl(s)$
Net Ionic Equations
- Shows only the species that actually change during the reaction.
- Example: ${Ag}^{+} (???? ????) + {Cl}^{-} (???? ????) \to AgCl(s)$

Important Concepts

Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.
Stoichiometry: The calculation of reactants and products in chemical reactions.
Reaction Conditions: Factors such as temperature, pressure, and catalysts that affect the rate and outcome of a reaction.

Classes of organic compounds

This table provides an overview of various classes of organic compounds, their functional groups, and examples of each class.

Class of Organic Compounds	Functional Group	Example
Hydrocarbons	-	Methane (CH₄), Ethylene (C₂H₄), Acetylene (C₂H₂)
Alcohols	Hydroxyl (-OH)	Methanol (CH₃OH), Ethanol (C₂H₅OH), Isopropanol (C₃H₇OH)
Aldehydes	Carbonyl (C=O) at end	Formaldehyde (CH₂O), Propanal (CH₃CH₂CHO)
Ketones	Carbonyl (C=O) within	Acetone (CH₃COCH₃), Butanone (CH₃COCH₂CH₃)
Carboxylic Acids	Carboxyl (COOH)	Formic acid (HCOOH), Acetic acid (CH₃COOH), Benzoic acid (C₆H₅COOH)
Esters	Carboxylate (COO-)	Methyl acetate (CH₃COOCH₃), Ethyl butanoate (CH₃CH₂COOCH₂CH₃)
Amines	Amino (NH₂)	Methylamine (CH₃NH₂), Dimethylamine (CH₃NHCH₃), Trimethylamine (N(CH₃)₃)
Ethers	Oxygen between two groups	Dimethyl ether (CH₃OCH₃), Diethyl ether (CH₃CH₂OCH₂CH₃), Anisole (CH₃OC₆H₅)
Halogenoalkanes	Halogen (F, Cl, Br, I)	Chloromethane (CH₃Cl), Bromoethane (CH₃CH₂Br), Iodoform (CHI₃)

Catalytic processes

Catalysis is a process by which the rate of a chemical reaction is increased by the presence of a substance called a catalyst. Catalysts work by providing an alternative reaction pathway with lower activation energy, thus enabling the reaction to proceed more rapidly. Catalytic processes play a crucial role in various industries, including petrochemicals, pharmaceuticals, and environmental remediation.

1. Homogeneous Catalysis

Definition: Homogeneous catalysis occurs when the catalyst is present in the same phase as the reactants (e.g., all are in the liquid phase or all are in the gas phase).
Mechanism: The catalyst forms an intermediate complex with the reactants, facilitating the reaction by lowering the activation energy.
Examples:
- Hydrogenation of alkenes using Wilkinson's catalyst (RhCl(PPh₃)₃).
- Decomposition of hydrogen peroxide catalyzed by iron(II) ions.

2. Heterogeneous Catalysis

Definition: Heterogeneous catalysis occurs when the catalyst is in a different phase from the reactants (e.g., solid catalyst with gas or liquid reactants).
Mechanism: The reactants adsorb onto the surface of the catalyst, where the reaction takes place, and then desorb as products.
Examples:
- Haber-Bosch process for ammonia synthesis using an iron catalyst.
- Fischer-Tropsch synthesis for hydrocarbon production from synthesis gas using metal catalysts such as cobalt or iron supported on alumina.

3. Enzyme Catalysis

Definition: Enzymes are biological catalysts that accelerate chemical reactions in living organisms.
Mechanism: Enzymes bind to specific substrates at their active sites, facilitating the conversion of substrates into products.
Examples:
- Digestive enzymes such as amylase, protease, and lipase catalyze the breakdown of carbohydrates, proteins, and fats in the digestive system.
- DNA polymerase catalyzes the polymerization of nucleotides during DNA replication.

4. Acid-Base Catalysis

Definition: Acid-base catalysis involves the use of acids or bases as catalysts to accelerate chemical reactions by proton donation or acceptance.
Mechanism: Acids donate protons to reactants, while bases accept protons from reactants, altering their reactivity.
Examples:
- Esterification reaction, where a carboxylic acid reacts with an alcohol in the presence of an acid catalyst (e.g., sulfuric acid) to form an ester.
- Transesterification reaction, where an alcohol reacts with an ester in the presence of a base catalyst (e.g., sodium methoxide) to form another ester.

5. Photocatalysis

Definition: Photocatalysis involves the use of a catalyst that is activated by light to initiate chemical reactions.
Mechanism: The catalyst absorbs photons and generates electron-hole pairs, which participate in redox reactions to drive the desired reaction.
Examples:
- Photocatalytic degradation of organic pollutants in water using semiconductor materials such as titanium dioxide (TiO₂).
- Photocatalytic water splitting for hydrogen production using photocatalysts like ruthenium-doped titanium dioxide.

Catalytic processes are essential for improving reaction efficiency, selectivity, and sustainability in chemical synthesis and environmental remediation. They enable the production of valuable chemicals, fuels, and pharmaceuticals while minimizing energy consumption and waste generation.

Solubility and Electrochemistry

Solubility

Solubility refers to the ability of a substance (solute) to dissolve in a solvent to form a homogeneous mixture called a solution. It is typically expressed as the maximum amount of solute that can dissolve in a given amount of solvent at a specified temperature and pressure. Solubility depends on various factors, including temperature, pressure, polarity, and the nature of the solute and solvent.

Factors Affecting Solubility
- Temperature: Generally, solubility increases with temperature for most solid solutes but decreases for gases.
- Pressure: The solubility of gases in liquids typically increases with pressure.
- Nature of Solute and Solvent: Polar solutes tend to dissolve in polar solvents, while nonpolar solutes dissolve in nonpolar solvents.
- Molecular Size and Structure: Smaller molecules with fewer intermolecular forces tend to be more soluble.
Types of Solutions
- Saturated Solution: Contains the maximum amount of solute dissolved in the solvent at a given temperature.
- Unsaturated Solution: Contains less solute than the maximum that could dissolve at a given temperature.
- Supersaturated Solution: Contains more solute than the maximum that could dissolve at a given temperature, achieved through careful manipulation.
Solubility Rules
- Used to predict the solubility of ionic compounds in water based on the properties of their ions.
- For example, most nitrate (NO₃⁻), acetate (CH₃COO⁻), and alkali metal (Group 1) compounds are soluble in water.

Electrochemistry

Electrochemistry deals with the study of chemical reactions involving the transfer of electrons between reactants. It encompasses redox reactions, electrolysis, electrochemical cells, and corrosion.

Redox Reactions
- Involve the transfer of electrons from one substance (reducing agent) to another (oxidizing agent).
- Reduction: Gain of electrons by a species (reduction half-reaction).
- Oxidation: Loss of electrons by a species (oxidation half-reaction).
Electrochemical Cells
- Consist of two half-cells connected by a conductive pathway (salt bridge or porous barrier).
- Types include galvanic (voltaic) cells, where spontaneous redox reactions generate electrical energy, and electrolytic cells, where non-spontaneous redox reactions are driven by an external electrical energy source.
Electrolysis
- Electrolysis is the process of using electrical energy to drive a non-spontaneous redox reaction.
- Common applications include electroplating, electrolytic refining of metals, and the production of chemicals like chlorine and sodium hydroxide.
Electrode Potentials
- Electrode potential (E) is a measure of the tendency of an electrode to lose or gain electrons.
- Standard electrode potentials (E°) are measured under standard conditions and are used to calculate the standard cell potential (E°cell) of electrochemical cells.
Corrosion
- Corrosion is the deterioration of metals due to chemical or electrochemical reactions with the environment.
- Types of corrosion include galvanic corrosion, where two dissimilar metals in contact form a galvanic cell, and electrolytic corrosion, where metals corrode due to exposure to electrolytes.

Chemical kinetics and Chemical equilibrium

Chemical Kinetics

Chemical kinetics is the branch of chemistry concerned with the study of the rates of chemical reactions and the factors that affect reaction rates. It provides insight into the mechanisms by which reactions occur and allows scientists to predict reaction rates under various conditions.

Reaction Rate
The rate of a chemical reaction is the change in concentration of reactants or products per unit time.
It is typically expressed as the change in concentration per unit time, often in units such as moles per liter per second (M/s).
Rate Laws
Rate laws describe the relationship between the rate of a reaction and the concentrations of reactants.
Reaction Mechanisms
Reaction mechanisms describe the series of elementary steps by which a reaction occurs.
Each elementary step involves the collision and transformation of reactant molecules or ions.
The overall reaction rate depends on the rate-limiting step, which is often the slowest step in the mechanism.
Factors Affecting Reaction Rate
Concentration: Generally, an increase in reactant concentration leads to an increase in reaction rate.
Temperature: Higher temperatures typically result in faster reaction rates due to increased molecular motion.
Catalysts: Catalysts increase reaction rates by providing an alternative reaction pathway with lower activation energy.
Surface Area: For heterogeneous reactions, a larger surface area of reactants typically leads to faster reaction rates.

Chemical Equilibrium

Chemical equilibrium occurs in a reversible reaction when the rates of the forward and reverse reactions are equal, resulting in the concentrations of reactants and products remaining constant over time. It is a dynamic state where reactions continue to occur, but there is no net change in the overall composition.

Equilibrium Constant
The equilibrium constant (K) is a quantitative measure of the extent of a reaction at equilibrium.
Le Chatelier's Principle
Le Chatelier's Principle states that if a system at equilibrium is subjected to a change in concentration, pressure, or temperature, the system will shift to counteract the change and restore equilibrium.
For example, if the concentration of a reactant is increased, the equilibrium will shift towards the product side to partially offset the increase.
Types of Equilibria
Homogeneous Equilibrium: Involves reactants and products in the same phase (e.g., gas phase or aqueous solution).
Heterogeneous Equilibrium: Involves reactants and products in different phases (e.g., gas-solid or liquid-solid).
Solubility Equilibrium: Occurs in saturated solutions where the dissolution and precipitation of a solid solute reach equilibrium.
Applications
Equilibrium concepts are applied in various fields, including chemical synthesis, industrial processes, environmental chemistry, and biological systems.
For example, the Haber-Bosch process for ammonia synthesis and the production of ethanol by fermentation rely on achieving chemical equilibrium.

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