University: Foundation Of Technical Republic of Iraq Institutes

The Ministry of Higher Education Institute: Kirkuk Tech. Ins.

& Scientific Research Department: Chemistry Industries

Lecturer name: Dr. Moneeb T. Salman

Academic Status:

Qualification: Ph.D.

Department Chemical Industries Major

Course Title Chemistry Code CHEM1 Course Instructor Dr. Moneeb T. Salman

Course Description: Semester 1 The course involves the study of atomic structure, modern periodic table, chemical bonds and formulae, states of matter, analytical chemistry, chemical reactions, acids and Contact Th. bases and salts, fundamentals of organic chemistry, Hours 1 extraction, chromatography and polymers. (Hour/week) Pr. 2

Total 3

General Goal: The student will become familiar with the fundamentals of chemistry such as: atomic structure, periodic table, fundamentals of analytical and organic chemistry, extraction, chromatography and polymers. Behavioural Objectives: The student will be able to: · Do a general revision of atomic structure and electronic distribution · Study the modern periodic table · Know how to build the chemical formulae of molecules and compounds and their nomenclature and the type of bonds between atoms · Know the three states of matter and their application to some daily observations · Identify acids, bases and salts · Chemical reactions and chemical equilibrium · Know the fundamentals of organic chemistry

1 Topics  The importance of chemistry, its common branches.  Atomic structure and electronic distribution.  Periodic table.  Chemical formulae of molecules and compounds, nomenclature and the types of bonds.  Acids, bases and salts.  Chemical reactions and chemical equilibrium.  Fundamentals of analytical chemistry, qualitative & quantitative chemical analysis, standard solution and indicators.  Units of concentration  Fundamentals of organic chemistry.  Extraction.  Chromatography.  Polymers.

Course Instructor Dr. Moneeb T. Salman

E-mail [email protected]

Title General Chemistry

Course Coordinator

The student will become familiar with the fundamentals of chemistry such as: atomic structure, periodic table, chemical reactions and fundamentals of organic Course Objective chemistry, extraction, chromatographic identification and polymers.

The course involves the study of atomic structure, modern

Course Description periodic table, chemical bonds and formulae, states of matter, chemical reactions, solution, acids and bases and salts, fundamentals of organic chemistry, chromatographic technology and polymerization.

Textbook Nil

2 1- ﺍﻟﻜﻴﻤﻴﺎﺀ ﺍﻟﻌﻀﻮﻳﺔ / ﺍﻟﺠﺰﺀ ﺍﻻﻭﻝ ﻭﺍﻟﺜﺎﻧﻲ – ﺗﺄﻟﻴﻒ ﺍﻟﺪﻛﺘﻮﺭ ﻓﻬﺪ ﻋﻠﻲ ﺣﺴﻴﻦ ﻭﻣﺴﺎﻋﺪﻳﻪ ، ﺍﻟﻄﺒﻌﺔ References ﺍﻻﻭﻟﻰ / ﺑﻐﺪﺍﺩ 1977 - 2- Organic Chemistry , by Morrson & Boy 1 , 3rd edition, 1975, USA - 3- Chemistry of organic compounds , by Noller. Philadelphia, USA . 1951 - 4- Organic Chemistry, Vol. 1 and 2 , by Finar – Longman Group . Ltd. 1973. 5- Practical Organic Chemistry , by Vogel . 3rd edition . published by Longman – Group Ltd .1975

1. 6- Laboratory outlines and note – Book for organic chemistry , C. E . Boord asnd W.R. Bord , 3rd edition. Join Wiley and sons, Inc. New York . 1967. 2. 7- Experimental Organic chemistry chemistry . by John Baldwin. Second edition . Mgraw – Hill , Book Co. New York. 1970 . 3. 8- Qualitative Organic analysis by Kemp. William .Mcgraw Hill . India, 1970 9- Practical Organic Chemistry by Mann and Sannders. Published by Longman , Group Ltd, London. 10- Chemistry of organic compounds by Noller W.B Saunders Co. London 1951.

Term Tests Laboratory Quizzes Project Final Exam

Course Assessment As (30%) As (10%) As (10%) ---- As (50%)

The student will become familiar with the fundamentals of chemistry such as: atomic structure, periodic

table, chemical reactions and fundamentals of organic chemistry and General Notes electrochemistry

3 Course Weekly Outline week Date Topics Covered Lab. Experiment Note Assignments

1 Introduction, branches and importance General Safety of chemistry

2 Atom, Atomic Structure, elements in Recrystallization periodic table.

3 gradation of properties in periodic Separation table, ionization, electro negativity and electronic distribution

4 Bonds, types of bonds, attraction forces between molecules, Distillation double and triple bonds

5 Fundamentals of analytical chemistry, Determination of qualitative & quantitative chemical Melting Point analysis, standard solution and indicators

6 Units of concentration, molarty. Determination of Boiling Point

7 Units of concentration, normaly, Extraction formalty.

8 Units of concentration, mole fraction Preparation of different and percentage with practical solutions examples.

9 Gravimetric quantity analysis, Preparation of Methane gravimetric factor.

10 Fundamentals of Organic Chemistry, Preparation of Ethane importance, hydrocarbons, alkyl halides

11 Alkanes, general formulae and Preparation of nomenclature. Acetylene

4 12 Alkanes, preparation, reactions and Preparation of physical properties. Ethylbromide

13 Alkenes, general formulae and Preparation of Acetone nomenclature.

14 Alkenes, preparation, reactions and Preparation of physical properties. Acetaldehyde

15 Alkynes, general formulae and Preparation of nomenclature. Actylchloride

16 Alkynes, preparation, reactions and Preparation of physical properties. Acetamide

17 Alcohols, general formulae and Preparation of Azoxyl nomenclature, preparation, reactions benzene and physical properties.

18 Ethers, general formulae and Preparation of nomenclature, preparation, reactions Bonozetriazole and physical properties.

19 Phenols, general formulae and Preparation of Phenol nomenclature, preparation, reactions and physical properties.

20 Ketones & Aldehydes, general Preparation of Benzyl formulae and nomenclature, alcohol

21 Preparation, reactions and physical Preparation of benzoic properties. acid from benzaldehyde

22 Carboxylic acids, general formulae Preparation of benzpic and nomenclature, preparation, acid from Toluen reactions and physical properties.

23 Ester, Amine & Amide Preparation of Acetanilide

24 Acids, bases and salts, nomenclature, Preparation of classification and reactions. .Nitrobenzene

5 25 Polymer, polymerization Preparation of n- dinitrobenzene

26 Types of polymerization. Preparation of Aniline

27 Polymers synthesis. Preparation of Ethylacetate

28 Physical properties of polymer. Preparation of Acetic acid

29 Preparation of Oxime Fundamentals of Chromatography, chromatographic techniques,

30 Polymers Column chromatography, Paper chromatography, Thin layer chromatography, Gas chromatography, Liquid chromatography

Detailed Curriculum (Theory) Hrs Contents Behavioral Objectives 1 Introduction, branches of chemistry and The student should : importance Become familiar with the definition and branches of chemistry. 2 2- Atomic Structure: Introduction Become familiar with the Electric nature of atoms constituents of an atom: nucleus, protons, neutrons and electrons. Constituents of atoms Learn how to distribute electrons Oxidation numbers and valence around the nucleus correctly Metals, non metals and noble gases Learn the building stages of the Chemical reactivity periodic table and the various Quantum numbers periods and groups within it. Electronic distribution Learn how to read information Periodic table and gradation of properties in from the table correctly it 1 3- Bonds Learn the symbols of chemical Types of bonds elements and memorize it and will Attraction forces between molecules also Double and triple bonds learn the nomenclature of compounds and the types of chemical bonds and the difference between them

6 Learn the types of physical bonds 1 4- Fundamentals of analytical chemistry,  Learn the qualitative chemical analysis qualitative & quantitative chemical analysis, ( chemical & flame test ) standard solution and indicators  quantitative chemical analysis ( gravimetric & volumetric analysis)  standard solution  indicators 3 5- Units of concentration, molarty, normality,  Become familiar with calculating & formality, mole fraction and percentage determining the concentration of any solution with different unit  Solve an examples by using the equation to calculate the (molarty, normality, formality, mole fraction and percentage ) of solutions. 6- Gravimetric quantity analysis, gravimetric  Become familiar with determination of 1 factor. an analyte based on the mass of a solid  Learn the precipitation, evaporation, crystallization, filtration and sublimation 14 7- Fundamentals of Organic Chemistry Learn the characterization ,properties , General formulae and nomenclature for preparation nomenclature the following function groups: Hydrocarbons( alkanes-alkenes-alkynescyclic  Classification of organic compounds hydrocarbons- alcohols- aldehydes- ketones- o Functional groups carboxylic acids- ethers- amines, amides, o Aliphatic compounds aromatic acids and esters ) o Aromatic compounds

1 8- Acids, bases and salts, Learn and reactions. Learn an acid-base, defintion  reaction, Learn 4 9- Polymer, polymerization, types of Learn the various types of polymerization. Synthesis and physical polymers properties of polymer. 2 10- Fundamentals of Chromatography and its Know the importance of techniques chromatographic separations for qualitative analysis

7 Week No. 1 / Introduction To General Chemistry

Preliminary questions:

1- What Is Chemistry 2- Why We Study Chemistry 3- What Fields of Study Use Chemistry? 4- What are the main branches of chemistry?

What Is Chemistry?

1. The science that studies the composition, properties, and activity of organic and inorganic substances and various elementary forms of matter.

2. Chemical properties, reactions, phenomena, etc.: the chemistry of carbon.

3. a. sympathetic understanding; rapport. b. sexual attraction.

4. The constituent elements of something.

Glossary definition "scientific study of matter, its properties, and interactions with other matter and with energy".

An important point to remember is that chemistry is a science, which means its procedures are systematic and reproducible and its hypotheses are tested using the scientific method. Chemists, scientists who study chemistry, examine the properties and composition of matter and the interactions between substances. Chemistry is closely related to physics and to biology. As is true for other sciences, mathematics is an essential tool for the study of chemistry

Why We Study Chemistry?

Because understanding chemistry helps you to understand the world around you. Cooking is chemistry. Everything you can touch or taste or smell is a chemical. When you study chemistry, you come to understand a bit about how things work. Chemistry isn't secret knowledge, useless to anyone but a scientist. It's the explanation for everyday things, like why laundry detergent works better in hot water or how baking soda works or why not all pain relievers work equally well on a headache. If you know some chemistry, you can make educated choices about everyday products that you use.

8 What Fields of Study Use Chemistry?

You could use chemistry in most fields, but it's commonly seen in the sciences and in medicine. Chemists, physicists, biologists, and engineers study chemistry. Doctors, nurses, dentists, pharmacists, physical therapists, and veterinarians all take chemistry courses. Science teachers study chemistry. Fire fighters and people who make fireworks learn about chemistry. So do truck drivers, plumbers, artists, hairdressers, chefs... the list is extensive.

Chemistry is the science that is concerned with the characterization, composition, and transformations, of matter. And it’s concerned with the forces that hold these structures together. The physical properties of substances are studied also.

Modern chemistry, which emerged late in the eighteenth century, took hundreds of years to develop. Chemistry has gradually developed five principal branches:

1- Organic chemistry: Is the chemistry of the compounds of carbon (except for few that are classified as inorganic compounds). The term organic is a holdover from the time when it was believed that these compounds could be derived only from plant or animal sources.

2- Inorganic chemistry: Is the chemistry of all the elements except carbon. Some simple carbon compounds (for example, the carbonates and carbon dioxide) are traditionally classified as inorganic compounds, since they can be derived from mineral sources.

3- Analytical chemistry: is the branch of chemistry involved with the identification of the composition, both qualitative and quantitative, of substances, studying the properties of materials or developing tools to analyze materials.

4- Physical chemistry: is the branch of chemistry that applies physics to the study of chemistry. The physical principles that underlie the structure of matter and chemical transformations, Quantum mechanics and thermodyamics are examples of physical chemistry.

5- Biochemistry: Is the branch of chemistry concerned with the chemical

9 reactions that occur inside living organisms both plant and animal.

There are several branches of chemistry. Here is a list of the main branches of chemistry, with an overview of what each branch of chemistry studies.

Agrochemistry - This branch of chemistry may also be called agricultural chemistry. It deals with the application of chemistry for agricultural production, food processing, and environmental remediation as a result of agriculture.

Astrochemistry - Astrochemistry is the study of the composition and reactions of the chemical elements and molecules found in the stars and in space and of the interactions between this matter and radiation.

Chemical Engineering - Chemical engineering involves the practical application of chemistry to solve problems.

Chemistry History - Chemistry history is the branch of chemistry and history that traces the evolution over time of chemistry as a science. To some extent, alchemy is included as a topic of chemistry history.

Electrochemistry - Electrochemistry is the branch of chemistry that involves the study of chemical reactions in a solution at the interface between an ionic conductor and an electrical conductor. Electrochemistry may be considered to be the study of electron transfer, particularly within an electrolytic solution.

Environmental Chemistry - Environmental chemistry is the chemistry associated with soil, air, and water and of human impact on natural systems.

Food Chemistry - Food chemistry is the branch of chemistry associated with the chemical processes of all aspects of food. Many aspects of food chemistry rely on biochemistry, but it incorporates other disciplines as well.

General Chemistry - General chemistry examines the structure of matter and the reaction between matter and energy. It is the basis for the other branches of chemistry.

Geochemistry - Geochemistry is the study of chemical composition and chemical processes associated with the Earth and other planets.

10 Green Chemistry - Green chemistry is concerned with processes and products that eliminate or reduce the use or release of hazardous substances. Remediation may be considered part of green chemistry.

Medicinal Chemistry - Medicinal chemistry is chemistry as it applies to pharmacology and medicine.

Nuclear Chemistry - Nuclear chemistry is the branch of chemistry associated with nuclear reactions and isotopes.

Photochemistry - Photochemistry is the branch of chemistry concerned with interactions between light and matter.

Polymer Chemistry - Polymer chemistry or macromolecular chemistry is the branch of chemistry the examines the structure and properties of macromolecules and polymers and finds new ways to synthesize these molecules.

Thermochemistry - Thermochemistry may be considered a type of Physical Chemistry. Thermochemistry involves the study of thermal effects of chemical reactions and the thermal energy exchange between processes.

Theoretical Chemistry - Theoretical chemistry applies chemistry and physics calculations to explain or make predictions about chemical phenomena.

Dimensional questions:

1- Define: chemistry, Inorganic chemistry, Organic chemistry, Biochemistry, Physical chemistry, Analytical chemistry. 2- Determine the branch of chemistry that deals with: a- thermal effects of chemical reactions. b- the structure and properties of macromolecules. c- interactions between light and matter. d- nuclear reactions and isotopes. e- pharmacology and medicine. f- processes and products that eliminate or reduce the use or release of hazardous substances. g- the Earth and other planets.

11 h- the chemical processes of all aspects of food. i- soil, air, and water and of human impact on natural systems. j- the study of chemical reactions in a solution. k- the practical application of chemistry to solve problems. l- the interactions between matter and radiation. m- agricultural production, food processing.

Week No. 2,3 /

ATOM , ATOMIC STRUCTURE AND THE PERIODIC TABLE

ATOM & ATOMIC STRUCTURE

An atom is the defining structure of an element, which cannot be broken by any chemical means. A typical atom consists of a nucleus of protons and neutrons with electrons circling this nucleus. Atom Examples: hydrogen, carbon-14, zinc, cesium, Cl- (a substance can be an atom and an isotope or ion at the same time)

Atom Definition: The smallest particle of an element that can combine with the atoms of other elements to form compounds. Proton: A subatomic particle that has a mass approximately 1.0073µ, carries a one – unit positive charge, and is found in the nucleus of the atom. Electron: A subatomic particle that has a mass approximately 0.00055µ, carries a one – unit negative charge, and is found outside the nucleus in an atom. Neutron: A subatomic particle that has a mass approximately 1.0087µ, is uncharged, and is found in the nucleus of an atom.

Ion: A particle made up of an atom or a group of atoms that bears an electric charge. An ion may have either a positive charge (because one or more

12 electrons have been lost) or negative charge (because one or more electrons have been gained). Atomic number A: is the number of protons in the nucleus of an atom of an element. In an uncharged atom, it is also equal to the number of electrons. Mass number M: The number of neutrons and protons, taken together, in the nucleus of an atom. Atomic symbols: an atom is identified by two numbers, the atomic and the mass numbers:

M

Symbol of element

A

M = number of neutrons + number of protons

Number of protons = A (In an uncharged atom)

Number of neutrons = M – A

Number of electrons = A (In an uncharged atom)

Atomic weight: The average mass of atoms of an element relative to the mass 12 of a C6 taken as exactly 12µ.

Ionization: Is the formation of an ion

Ionization energy (or enthalpy): The standard energy change when one electron is removed from a gaseous atom, molecule or ion.

Electro negativity: The measure of the ability of an atom to attract bonding electrons to itself when it is joined to another atom by a covalent bond.

Electronic structure of atoms:

According to electronic theory atoms consist of positive nucleus surrounding by negatively charged electrons. The sum of the negative charges is equal to the number of positive charges in the nucleus; the latter is called the atomic number.

From spectroscopic studies we know that the electrons are distributed around

13 the nucleus in successive shells or energy level of increasing radius as shown in the fig:

nucleus energy level

1- Each shell of electrons lies at a definite radius; the electrons in each shell are at the same energy.

2- Each shell is labeled with n. The shell closest to the nucleus is labeled n = 1. The higher the value of n, the higher the energy of the electron.

3- Each shell can only hold a fixed number of electrons. (The total number of electrons allowed in a shell is 2n2)

4- Electrons will occupy as low an energy level as possible; this means that if electrons are fed in to the shells of an atom the lowest energy shells are filled first.

For identification each one of these energy levels is assigned a letter and a number called a principal quantum number. The first shell is assigned the letter (k) and the principal quantum number (n) equal to one (1) this shell is of lower energy.

The maximum number of electrons that can possibly occupy a shell is equal to 2n2 Thus for the (k) shell in which (n=1). The maximum number of electrons 2(1)2 = 2 electrons.

Name of p.q. The maximum number shell number of electrons (n) K 1 2(1)2 = 2 L 2 2(2)2 = 8 M 3 2(3)2 = 18 N 4 2(4)2 = 32 O 5 2(5)2 = 50 P 6 2(6)2 = 72

14 Q 7 2(7)2 = 98

Sublevel (sub shell) of energy:

All electrons in the same shell possess the same energy. More detailed experiments show that with the exception of electrons in the first shell this is not the case, and the main energy levels of atoms possess sublevels of energy. The sublevels include the four types s, p, d and f :

Sublevel contain electrons No. of orbitals type up to s 2 1 p 6 3 d 10 5 f 14 7

Most of the physical and chemical properties of atoms, and hence of all matter, are determined by the nature of the electron cloud enclosing the nucleus.

S-orbital Px Py Pz

Electronic cloud representation of s and p orbitals

15 Electron configurations of atoms:

There is no more than two electrons may occupy the same orbital in an opposite spins (clockwise or counter clockwise). Electrons with anti parallel spins are represented by arrows pointing in opposite direction .

 a list showing how many electrons are in each orbital or subshell in an atom or ion  subshell notation: list subshells of increasing energy, with number of electrons in each subshell as a superscript o examples o 1s2 2s2 2p5 means "2 electrons in the 1s subshell, 2 electrons in the 2s subshell, and 5 electrons in the 2p subshell" o 1s2 2s2 2p6 3s2 3p3 is an electron configuration with 15 electrons total; 2 electrons have n=1 (in the 1s subshell); 8 electrons have n=2 (2 in the 2s subshell, and 6 in the 2p subshell); and 5 electrons have n=3 (2 in the 3s subshell, and 3 in the 3p subshell).  ground state configurations fill the lowest energy orbitals first

16 Electron configurations of the first 11 elements, in subshell notation. Notice how configurations can be built by adding one electron at a time. atom A ground state electronic configuration

H 1 1s1

He 2 1s2

Li 3 1s2 2s1

Be 4 1s2 2s2

B 5 1s2 2s2 2p1

C 6 1s2 2s2 2p2

N 7 1s2 2s2 2p3

O 8 1s2 2s2 2p4

F 9 1s2 2s2 2p5

Ne 10 1s2 2s2 2p6

Na 11 1s2 2s2 2p6 3s1

Electron Configuration of Nickel

Electrons surround the nucleus of an atom in patterns of shells and sub shells. In this table showing the electron configuration of a nickel atom, the large numbers (1, 2, 3, 4) indicate shells of electrons (shown as small spheres), the letters (s, p, d) indicate sub shells within these shells, and the exponents indicate the number of electrons present in each sub shell. Sub shells may be further divided into orbitals. Each orbital can contain two electrons, and orbitals are designated in the table by horizontal bars connecting pairs of

17 electrons. The small up and down arrows indicate the direction of each electron’s spin. Electrons that occupy the same

orbital always have opposite spins. If all the electrons were stripped away from an atom of nickel (that is, the atom was totally ionized) and electrons were allowed to return one at a time, the electrons would fill up the slots indicated on the chart from left to right, top to bottom. Electrons do not always fill all the sub shells of a shell before beginning to fill the next shell. The s sub shell of shell 4, for example, actually fills before the d sub shell of shell 3 (shown as the lowest row in this chart).

Writing electron configurations

 strategy: start with hydrogen, and build the configuration one electron at a time. 1. fill sub shells in order by counting across periods, from hydrogen up to the element of interest:

18 2. rearrange subshells (if necessary) in order of increasing n & l  examples: Give the ground state electronic configurations for:

o Al o Fe o Ba o Hg watch out for d & f block elements; orbital interactions cause exceptions to the Aufbau principle o half-filled and completely filled d and f subshells have extra stability

Know these exceptions to the Aufbau principle in the 4th period. (There are many others at the bottom of the table, but don't worry about them now.) configuration predicted by true ground state exception the Aufbau principle configuration

Cr 1s2 2s2 2p6 3s2 3p6 3d4 4s2 1s2 2s2 2p6 3s2 3p6 3d5 4s1

Cu 1s2 2s2 2p6 3s2 3p6 3d9 4s2 1s2 2s2 2p6 3s2 3p6 3d10 4s1

Electron configurations including spin

 unpaired electrons give atoms (and molecules) special magnetic and chemical properties  when spin is of interest, count unpaired electrons using orbital box diagrams  drawing orbital box diagrams 1. write the electron configuration in subshell notation 2. draw a box for each orbital. . Remember that s, p, d, and f subshells contain 1, 3, 5, and 7 degenerate orbitals, respectively. . Remember that an orbital can hold 0, 1, or 2 electrons only, and if there are two electrons in the orbital, they must have opposite (paired) spins (Pauli principle)

19 3. within a subshell (depicted as a group of boxes), spread the electrons out and line up their spins as much as possible (Hund's rule)  the number of unpaired electrons can be counted experimentally

o configurations with unpaired electrons are attracted to magnetic fields (paramagnetism) o configurations with only paired electrons are weakly repelled by magnetic fields (diamagnetism)

Periodic Table: Is the arrangement of the atoms of elements according to increasing of atomic number. The columns of elements in the table are called groups (or families) and elements in the same group often have similar properties. There are 116 elements although elements 110-116 are as yet unnamed. These 116 elements are organized in the periodic table:

The modern chemical symbols were introduced by Berzelius. The horizontal rows of elements in the table are called "periods" and the vertical columns of elements are called "groups" (1A, 2A 3B etc.). There are three general classes of elements distinguished by their physical properties: the metals (generally shiny and conduct electricity), the nonmetals (not shiny, sometimes gasses at STP and poor conductors of electricity) and the metalloids (properties in between those of metals and nonmetals.). Some groups have special names:

Group 1A: Alkali metals Group 2A: Alkali earth metals

Groups 3B-2B: Transition metals Group 7A: Halogens

Group 8A: Noble gases

Elements within a group share similar chemical properties. Other chemical and physical properties of the elements can be deduced from their position in the periodic table. The structure of the periodic table and thus their chemical and physical properties is directly related to their atomic structure.

*Hydrogen does not really belong in any group of the table. So it is usually placed at the top, on its own.

20 1 2 H He

3 4 5 6 7 8 9 10 Li Be B C N O F Ne

11 12 13 14 15 16 17 18 Na Mg Al Si P S Cl Ar

19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

55 56 * 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 Cs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

87 88 ** 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118 Fr Ra Rf Db Sg Bh Hs Mt Ds Rg Cp Uut Uuq Uup Uuh Uus Uuo

57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu

89 90 91 92 93 94 95 96 97 98 99 100 101 102 103 Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

Alkali Alkaline Basic Metal Earth Metal

Noble Non Halogen Gas Metal

Rare Earth

Classification of elements

 features of the periodic table o Periods are horizontal rows on the table. o Groups (or families) are columns on the table. . elements in the same group are called congeners. They have similar chemical properties. o Blocks are regions on the table.  important groups:

21 o alkali metals (Group 1A, first column ) . soft, extremely reactive metals . react with cold water to form hydrogen gas . form +1 ions o alkaline earth metals (Group 2A, second column): . soft, reactive metals . compounds are a major component of earth's crust . form +2 ions o halogens (Group 7A, next-to-last column): . poisonous and extremely reactive nonmetals . fluorine and chlorine are yellow-green gases . bromine is a volatile red-brown liquid . iodine is a volatile blue black solid . all form -1 ions o noble gases (Group 0, last column) . all are monatomic gases . a. k. a. inert gases; almost completely unreactive  Important blocks: o transition metals are the elements in the region from the third to twelfth columns. . hard, dense metals . less reactive than Group IA and IIA o rare earth metals are the elements in the annex at the bottom of the table. . lanthanides (annex, top row) . actinides (annex, bottom row) o main group elements are all elements except the transition and rare earth metals. . group numbers end with "A" o metals, nonmetals, and metalloids (semimetals) . metallic properties . luster . malleability: can be hammered into thin sheets . ductility: can be drawn into wire . conduct heat and electricity well

The chemical bond

22  Electrons are the glue that holds groups of atoms together.

 valence electrons are the electrons in an unfilled outer shell.  Since these are the outermost electrons, these are the ones that bringing another atom close by. These are the ones that are involved in bonding.  Define: a chemical bond is the result of a redistribution of electrons that leads to a more stable configuration between two or more atoms.

Types of Chemical bonds

 We can classify chemical bonding into six major types: ionic, covalent, metallic, polar, hydrogen and van der Waals.

Ionic

o Ionic bonding occurs between a pair of atoms when one of the atoms gives up its valence electrons to the other. The result is that both atoms have filled shells. Both atoms also end up with a charge, one negative, and the other positive. We call the positive charged atom a cation, and the negatively charged one, an anion.

o A classic example of ionic bonding is between Na and Cl. Na is a silvery metal. It has 1 valence electron. Cl is a yellow-green gas, and it needs 1 electron to fill its valence shell. If you put the gas

23 and the metal together, then they will burn as electrons are exchanged. The metal dissolves and the gas disappears. The ions now have opposite charges and are attracted to each other by electrostatic forces. They form a crystal with the rock salt structure.

o Ionic bonds commonly form between atoms of the

1st column and the 7th, and between the 2nd column

and the 6th. E.g. NaCl, MgO

o Materials composed of ionic bonds have distinctive properties. E.g.CaCl2

Covalent

24 o Covalent bonds often form between atoms with too many electrons in their valence shells to give away, o but not enough to easily fill. Thus they share electrons with their neighbors, in such a way that including the shared electrons the shells are full.

o E.g. H2. N2, O2, diamond. o These are stronger bonds than either of the other two types. This is because the electrons are shared. o The carbon bond usually forms covalent bonds. The possible arrangements are many. Life is based upon carbon-carbon bonding. The branch of science called organic chemistry is the study of carbon-based molecules.

25 o Your body is based upon carbon bonding. o So the covalent bond is considered the most important bond.  Ionic Bonding o Some atoms gain electrons to become anions o Others lose electrons to become cations o Ions are attracted by their opposing charges o Electrical Neutrality Maintained o Most Important Bonding in Rocks and Minerals  Covalent Bonding o Electrons share electrons to fill incomplete shells

Metallic

o Metallic bonding occurs between atoms that have a small number of electrons in their valence shells. They give the electrons up, not just to one other atom, but to the complete group of atoms. We think of the electrons as becoming loose. Positively charged atoms sitting in a sea of electrons. We often call it electron-gas. E.g. sodium.

26 This arrangement provides metals with many of their characteristic physical properties. The shiny luster is because with so many free electrons.

A. Outermost electrons wander freely through metal. Metal consists of cations held together by negatively-charged electron "glue." B. Free electrons can move rapidly in response to electric fields, hence metals are a good conductor of electricity. C. Free electrons can transmit kinetic energy rapidly, hence metals are good conductors of heat. D. The layers of atoms in metal are hard to pull apart because of the electrons holding them together, hence metals are tough. But individual atoms are not

27 held to any other specific atoms; hence atoms slip easily past one another. Metallic Bonding is the basis of our industrial civilization.

Hydrogen Bonding

Hydrogen Bonding is Geologically Important

A. Water molecules are asymmetrical. The positively-charged portions of one are attracted to the negatively-charged parts of another. It takes a lot of energy to pull them apart. Hence: *Water melts and boils at unusually high temperatures for such a light molecule. *Water has a high heat capacity. *It takes a lot of energy to melt ice and vaporize water. *Thus water is the principal heat reservoir on the Earth.

B. The asymmetrical charge distribution on a water molecule makes it very effective in dissolving ionically-bonded materials. However, it is not an effective solvent of covalently bonded materials (oil and water don't mix). Hence: *Water is very effective at weathering rocks and minerals. It is the closest thing to a universal solvent. *Water is very effective at transporting ions and dissolved nutrients in the human body. *Water is not an effective solvent of organic molecules. Thus we do not dissolve in our own cell fluids. Nifty feature.

28 C. When water freezes, it assumes a very open structure and actually expands. Most materials shrink when they freeze and sink in their liquid phases. Implications: *If ice sank like most frozen solids, it would accumulate at the bottoms of frozen lakes and seas. Most of the world's water would be ice. *Expansion of ice in rocks is a powerful weathering agent.

Polar bonds

o This is a different kind of bond. It is not due to sharing of electrons, but results from the fact that some molecules distribute electrons in such a way that they have a region that seems charged. I.e. they become polar. o E.g. water, H2O. The side of the molecule with the hydrogens is slightly positive, the side with the O is slightly negative. The complete molecule is neutral. This means that the water molecule is stable, yet it can appear to be charged, depending upon its orientation. This is why water can dissolve so many things, e.g. water dissolves NaCl.

Van der Waals

o Dynamics polar bonding. E.g. He2

29 o E.g. clays. Sheets of strong covalent and ionic bonded atoms held together by weak Van der Waals forces. E.g. graphite. Like stacks of paper.

Summary of Bonding

 Ionic bonding holds rocks and minerals together  Covalent bonding holds people and other organisms together  Metallic bonding holds civilization together  Hydrogen bonding gives water its heat-retaining and solvent properties  Polar, hydrogen bonds and Van der Waals forces are weak.

If we want to determine the type of bonding in the compounds (H2O- HCL-

CaCL2 and CH4) we may use the following rules :

The rule: If the different in electro negativity between two atoms sharing a bond is between:

Electro negativity The type of bonding 0 – 0.6 the bond is covalent bond 0.7 – 1.6 the bond is polar covalent bond. 1.7 – higher the bond is ionic bond Example:

Compounds E. N The difference Type of bond

H2O H=2.1 O= 3.5 1.4 polar covalent HCL H=2.1 CL=3.0 0.9 polar covalent

CaCL2 CL=3.0 Ca= 1 2 ionic bond

CH4 C = 2.5 H=2.1 0.4 covalent bond

30 If the electronegativity values of two atoms are: *similar... Metallic bonds form between two metal atoms. Covalent bonds form between two non-metal atoms. Nonpolar covalent bonds form when the electronegativity values are very similar. Polar covalent bonds form when the electronegativity values are a little further apart. *different... Ionic bonds are formed.

Lewis Structures or Electron Dot Structures:

A Lewis structure is a type of shorthand notation. Atoms are written using their element symbols. Lines are drawn between atoms to indicate chemical bonds. Single lines are single bonds. Double lines are double bonds. Triple lines are triple bonds. (Sometimes pairs of dots are used instead of lines, but this is uncommon. Steps to Drawing a Lewis Structure

1. Pick a Central Atom

2.Count Electrons

3.Place Electrons around Atoms

Q1: The electrons in a nonpolar covalent bond are:

(gained, lost, shared equally, shared unequally) Q2: What is the charge on the ions formed by the alkaline earth metals?

(+1, +2, -1, -2)

Q:What is the most correct name for the ionic compound formed by Fe2+ and Cl-?

31 (iron chloride, iron (I) chloride, iron (II) chloride, iron (III) chloride)

Q4: What type of bonds are formed in N2O4 and what is the name of this compound?

(covalent, dinitrogen tetroxide /covalent, nitrogen tetroxide/ionic, nitrogen oxide/ ionic, dinitrogen oxide)

Q5: The bond between sulfur (electronegativity value 2.5) and chlorine (electronegativity value 3.0) would be: (not formed, ionic, polar covalent, nonpolar covalent)

Q6: What is the formula for the ion which has 17 protons and 18 electrons? (Cl+, Cl-, Ar+, Ar- )

Q7: Ionic compounds may contain polyatomic ions, which consist of groups of atoms having an electrical charge. An example is magnesium nitrate. The formula of magnesium nitrate is: [ MgNO3 , Mg2NO3, Mg(NO3)2, Mg2(NO3)3 ]

Q8: What is the formula of phosphorus trichloride? ( KCl, KCl3, P3Cl, PCl3 )

Q9: How many electrons are gained/lost by magnesium and what is the charge on the ion that it forms?

A - loses 2 electrons to form a magnesium ion with a 2- charge b - gains 2 electrons to form a magnesium ion with a 2- charge c - loses 2 electrons to form a magnesium ion with a 2+ charge d - gains 2 electrons to form a magnesium ion with a 2+ charge

Q10: The electron-dot structure of carbon has how many dots? (2, 4, 6, 8)

32 Week No. 5, 6 / ANALYTICAL CHEMISTRY

Analytical Chemistry: is the study of the separation, identification, and quantification of the chemical components of natural and artificial materials. Qualitative analysis gives an indication of the identity of the chemical species in the sample and quantitative analysis determines the amount of one or more of these components. The separation of components is often performed prior to analysis. Analytical methods can be separated into classical and instrumental. Classical methods (also known as wet chemistry methods) use separations such as precipitation, extraction, and distillation and qualitative analysis by color, odor, melting point or by measurement of weight or volume. Instrumental methods use an apparatus to measure physical quantities of the analyte such as light absorption, fluorescence, or conductivity. The separation of materials is accomplished using chromatography or electrophoresis methods. Types of chemical analysis: The processes that applied on a substance to identify its constituents and their percentages, divided in to two types are:

1- Qualitative chemical analysis: It used for identification a substance constituents and the way of their combinations to separate and detect cations and anions in a sample substance. 1) Chemical tests 2) Flame test First, ions are removed in groups from the initial aqueous solution. After each group has been separated, then testing is conducted for the individual ions in each group. Here is a common grouping of cations: + 2+ 2+ Group 1: Ag , Hg2 , Pb Precipitated in 1 M HCl Group 2: Bi3+, Cd2+, Cu2+, Hg2+, (Pb2+), Sb3+ and Sb5+, Sn2+ and Sn4+ Precipitated

in 0.1 M H2S solution at pH 0.5 Group 3: Al3+, (Cd2+), Co2+, Cr3+, Fe2+ and Fe3+, Mn2+, Ni2+, Zn2+ Precipitated in

0.1 M H2S solution at pH 9 2+ 2+ + 2+ + + 2+ 2+ 2+ Group 4: Ba , Ca , K , Mg , Na , NH4 , Ba , Ca , and Mg are precipitated in

33 0.2 M (NH4)2CO3 solution at pH 10; the other ions are soluble Many reagents are used in qualitative analysis, but only a few are involved in nearly every group procedure. The four most commonly used reagents are 6M

HCl, 6M HNO3, 6M NaOH, 6M NH3. Understanding the uses of the reagents is helpful when planning an analysis. Quantitative chemical analysis: It used to determine the amount of substance constituents in a certain weight of the same substance and the percentage of each constituent. Solution: A homogeneous mixture of two or more pure substances that evenly distributed in each other. Solute: A component of a solution that is present in an amount is smaller than the amount of the solvent. The solute is said to be dissolved in the solvent. Solvent: The component of a solution that is present in the largest amount or that determines the physical state of the solution. Standard solution: A solution that has a known concentration of solute. The standard solutions are used to find the concentrations of substances which have unknown concentrations, by using standardization of the solution of the substance with a standard solution. Indicators: A substance used in a titration that signals the end of the titration by changing color. Indicators are organic compounds (acid or base) with weak ionization, and the color of its molecule differs than the color of its ions. *The acidic indicator do not ionize in acidic medium so the color of its molecule will appear, but ionize in a basic medium, so the color of its ion will appear. *The basic indicator does not ionize in a basic medium but ionized in acidic medium.*The most important indicators and the range of its color change are:-

Indicator Color in Type Range of color change Basic Acidic medium medium Methyl orange Yellow Red Weak base Between PH 3.1-4 Phenophthalene purple Colorless Weak acid Between PH 8.3-10 Letmous paper Blue red Weak acid Between PH 5-8

Gravimetric analysis : Gravimetric analysis involves determining the amount of material present by

34 weighing the sample before and/or after some transformation. A common example used in undergraduate education is the determination of the amount of water in a hydrate by heating the sample to remove the water such that the difference in weight is due to the loss of water.

Week No. 7, 8, 9 / Concentration of a chemical solution

The concentration of a chemical solution refers to the amount of solute that is dissolved in a solvent. We normally think of a solute as a solid that is added to a solvent (e.g., adding table salt to water), but the solute could just as easily exist in another phase. For example, if we add a small amount of ethanol to water, then the ethanol is the solute and the water is the solvent. If we add a smaller amount of water to a larger amount of ethanol, then the water could be the solute! Units of Concentration: Once you have identified the solute and solvent in a solution, you are ready to determine its concentration. Concentration may be expressed several different ways, using percent composition by mass, mole fraction, molarity, molality, or normality. 1. Percent Composition by Mass or volume or both (%) This is the mass of the solute divided by the mass of the solution (mass of solute plus mass of solvent), multiplied by 100. Example: Determine the percent composition by mass of a 100 g salt solution which contains 20 g salt. Solution: 20 g NaCl / 100 g solution x 100 = 20% NaCl solution

2- Mole Fraction (X) This is the number of moles of a compound divided by the total number of moles of all chemical species in the solution. Keep in mind, the sum of all mole fractions in a solution always equals 1. Wt No. of moles (n) =------

M. wt

nA

XA = ------

nA + nB + nC +……ect XA = is the mole fraction of A, nA,nB,nC …. are the number of mole of

35 A,B,C,…….The sum of the mole fractions of all the components present in the solution must equal 1.

XA+XB+XC+……….= 1

Ex1: A gaseous solution contains 2g of He and 4g of O2.

What are the mole fractions Of He and O2 in the solution? (molecular weight Helium = 4;

molecular weight of O2 = 16)

Sol: we first find the number of moles of each component present in the solution

Ex2: What are the mole fractions of the components of the solution formed when 92 g glycerol is mixed with 90 g water? (molecular weight water = 18; molecular weight of glycerol = 92) Sol.: 90 g water = 90 g x 1 mol / 18 g = 5 mol water 92 g glycerol = 92 g x 1 mol / 92 g = 1 mol glycerol total mol = 5 + 1 = 6 mol

xwater = 5 mol / 6 mol = 0.833

x glycerol = 1 mol / 6 mol = 0.167 It's a good idea to check your math by making sure the mole fractions add up to 1:

xwater + xglycerol = 0.833 + 0.167 = 1.000

3- Molarity (M) is probably the most commonly used unit of concentration. It is the number of moles of solute per liter of solution (not necessarily the same as the volume of solvent!). Moles = Molarity x Volume (i.e., moles= M x V) Wt 1000

(molarity) =------X------

M.WT V

Example:

What is the molarity of a solution made when water is added to 11 g CaCl2 to make 100 mL of solution?

Solution:11 g CaCl2 / (110 g CaCl2 / mol CaCl2) = 0.10 mol CaCl2 100 mL x 1 L / 1000 mL = 0.10 L molarity = 0.10 mol / 0.10 L molarity = 1.0 M

4. Molality (m) Molality is the number of moles of solute per kilogram of solvent. Because the density of water at 25°C is about 1 kilogram per liter, molality is approximately equal to molarity for dilute aqueous solutions at this temperature. This is a useful approximation, but remember that it is only an approximation and doesn't apply when the solution is at a different

36 temperature, isn't dilute, or uses a solvent other than water.

Wt 1000

(mole/kg) m = ------X ------

M.WT V

Example: What is the molality of a solution of 10 g NaOH in 500 g water? Solution: 10 g NaOH / (40 g NaOH / 1 mol NaOH) = 0.25 mol NaOH 500 g water x 1 kg / 1000 g = 0.50 kg water molality = 0.25 mol / 0.50 kg molality = 0.05 M / kg molality = 0.50 m

5. Normality (N) Normality is equal to the gram equivalent weight of a solute per liter of solution. A gram equivalent weight or equivalent is a measure of the reactive capacity of a given molecule. Normality is the only concentration unit that is reaction dependent. Wt 1000

N= ------X------

Eq.WT V

Example: 1 M sulfuric acid (H2SO4) is 2 N for acid-base reactions because each mole of sulfuric acid provides 2 moles of H+ ions. On the other hand, 1 M sulfuric acid is 1 N for sulfate precipitation, since 1 mole of sulfuric acid provides 1 mole of sulfate ions.

Dilutions You dilute a solution whenever you add solvent to a solution. Adding solvent results in a solution of lower concentration. You can calculate the concentration of a solution following a dilution by applying one of the equations: Initial Conc. X Initial Vol = Final Conc X Final Vol Mi x Vi = Mf x Vf or

N1 X V1 = N2 X V2 where M is molarity, N is normality, V is volume, and the subscripts i and f refer to the initial and final values. Example: How many millilieters of 5.5 M NaOH are needed to prepare 300 mL of 1.2 M NaOH? Solution:

5.5 M x V1 = 1.2 M x 0.3 L V1 = 1.2 M x 0.3 L / 5.5 M V1 = 0.065 L V1 = 65 mL So, to prepare the 1.2 M NaOH solution, you pour 65 mL of 5.5 M NaOH into your container and add water to get 300 mL final volume. More Sample Concentration Calculations Most laboratories keep stock

37 solutions of common or frequently used solutions of high concentration. These stock solutions are used for dilutions. Problem: Calculate the amount of 1 M NaOH aqueous solution needed to make 100 mL of 0.5 M NaOH aqueous solution. Formula needed: M = m/V where M = molarity of solution in mol/liter m = number of moles of solute V = volume of solvent in liters Step 1:Calculate the number of moles of NaOH needed for 0.5 M NaOH aqueous solution. M = m/V 0.5 mol/L = m/(0.100 L) solve for m: m = 0.5 mol/L x 0.100 L = 0.05 mol NaOH. Step 2: Calculate the volume of 1 M NaOH aqueous solution that gives that gives the number of moles of NaOH from step 1. M = m/V V=m/M V = (0.05 moles NaOH)/(1 mol/L) V = 0.05 L or 50 mL Answer: 50 mL of 1 M NaOH aqueous solution is needed to make 100 mL of 0.5 M NaOH aqueous solution. Common ways to express the concentrations of solutions:-The concentration of a solute in a solution can be expressed in several different ways are: Problem: Calculate the amount of 1 M NaOH aqueous solution needed to make 100 mL of 0.5 M NaOH aqueous solution. Formula needed: M = m/V where M = molarity of solution in mol/liter m = number of moles of solute V = volume of solvent in liters Step 1: Calculate the number of moles of NaOH needed for 0.5 M NaOH aqueous solution. M = m/V 0.5 mol/L = m/(0.100 L) solve for m: m = 0.5 mol/L x 0.100 L = 0.05 mol NaOH. Step 2: Calculate the volume of 1 M NaOH aqueous solution that gives that gives the number of moles of NaOH from step 1. M = m/V V=m/M V = (0.05 moles NaOH)/(1 mol/L) V = 0.05 L or 50 mL Answer: 50 mL of 1 M NaOH aqueous solution is needed to make 100 mL of 0.5 M NaOH aqueous solution.

38

Week No. 10 / ORGANIC CHEMISTRY, Hydrocarbon

In organic chemistry, a hydrocarbon is an organic compound consisting entirely of hydrogen and carbon. With relation to chemical terminology, aromatic hydrocarbons, alkanes, alkenes and alkynes-based compounds composed entirely of carbon and hydrogen are referred to as "pure" hydrocarbons, whereas other hydrocarbons with bonded compounds or impurities of sulfur or nitrogen, are referred to as "impure", and remain somewhat erroneously referred to as hydrocarbons. Hydrocarbons are referred to as consisting of a "backbone" or "skeleton" composed entirely of carbon and hydrogen and other bonded compounds, and have a functional group that generally facilitates combustion. The majority of hydrocarbons found naturally occur in crude oil, where decomposed organic matter provides an abundance of carbon and hydrogen which, when bonded, can catenate to form seemingly limitless chains.

Types of hydrocarbons

The classifications for hydrocarbons defined by IUPAC nomenclature of organic chemistry (International union of pure and Applied chemistry) are as follows:

1. Saturated hydrocarbons (alkanes) are the most simple of the hydrocarbon species and are composed entirely of single bonds and are saturated with hydrogen. The general formula for saturated hydrocarbons is CnH2n+2 (assuming non-cyclic structures). Saturated hydrocarbons are the basis of petroleum fuels and are either found as linear or branched species. Hydrocarbons with the same molecular formula but different structural formulae are called structural isomers. Unsaturated hydrocarbons have one or more double or triple bonds between carbon atoms. Those with one double bond are called alkenes, with the formula CnH2n (assuming non-cyclic structures). Those containing triple bonds are called alkynes, with general formula CnH2n-2.

39 2. Cycloalkanes are hydrocarbons containing one or more carbon rings to which hydrogen atoms are attached. The general formula for a saturated hydrocarbon containing one ring is CnH2n. 3. Aromatic hydrocarbons, also known as arenes, are hydrocarbons that have at least one aromatic ring.

Hydrocarbons can be gases (e.g. methane and propane), liquids (e.g. hexane and benzene), waxes or low melting solids (e.g. paraffin wax and naphthalene) or polymers (e.g. polyethylene, polypropylene and polystyrene).

Week No. 11, 12 / Alkanes

Alkanes, also known as paraffins, are chemical compounds that consist only of the elements carbon (C) and hydrogen (H) (i.e., hydrocarbons), wherein these atoms are linked together exclusively by single bonds (i.e., they are saturated compounds) without any cyclic structure (i.e. loops). Each carbon atom must have 4 bonds (either C-H or C-C bonds), and each hydrogen atom must be joined to a carbon atom (H-C bonds). A series of linked carbon atoms is known as the carbon skeleton or carbon backbone. In general, the number of carbon atoms is often used to define the size of the alkane (e.g., C2-alkane). An alkyl group is NOT a functional group and does not show any characteristic reactions since it consists solely of strong and non-polar carbon-carbon single and carbon-hydrogen single bonds. It is important to note that free radical substitution is NOT a characteristic reaction of alkanes and alkyl groups. Saturated hydrocarbons can be linear (general formula CnH2n + 2) wherein the carbon atoms are joined in a snake-like structure, branched (general formula CnH2n + 2, n > 3) wherein the carbon backbone splits off in one or more directions, or cyclic (general formula CnH2n, n > 2) wherein the carbon backbone is linked so as to form a loop. According to the definition by IUPAC, the former two are alkanes, whereas the third group is called cycloalkanes. Saturated hydrocarbons can also combine any of the linear, cyclic (e.g., polycyclic) and branching structures, and they are still alkanes (no general formula) as long as they are acyclic (i.e., having no loops). The simplest possible alkane (the parent molecule) is methane, CH4. There is no limit to the number of carbon atoms that can be linked together, the only limitation being that the molecule is acyclic, is saturated, and is a hydrocarbon. Saturated oils and waxes are examples of larger alkanes where the number of carbons in the carbon backbone tends to be greater than 10. Alkanes are not very reactive and have little biological activity. Alkanes can be viewed as a molecular scaffold upon which can be hung the interesting biologically active/reactive portions (functional groups) of the molecule. IUPAC naming conventions can be used to produce a systematic name.

40 The key steps in the naming of more complicated branched alkanes are as follows:

 Identify the longest continuous chain of carbon atoms  Name this longest root chain using standard naming rules  Name each side chain by changing the suffix of the name of the alkane from "-ane" to "-yl"  Number the root chain so that sum of the numbers assigned to each side group will be as low as possible  Number and name the side chains before the name of the root chain  If there are multiple side chains of the same type, use prefixes such as "di-" and "tri-" to indicate it as such, and number each one.

Comparison of nomenclatures for three isomers of C5H12

Common name n-pentane isopentane neopentane

IUPAC name pentane 2-methylbutane 2,2-dimethylpropane

Structure

Boiling point

Melting (blue) and boiling (pink) points of the first 14 n-alkanes in °C.

41 Under standard conditions, from CH4 to C4H10 alkanes are gaseous; from C5H12 to C17H36 they are liquids; and after C18H38 they are solids. A straight-chain alkane will have a boiling point higher than a branched-chain alkane due to the greater surface area in contact, thus the greater van der Waals forces, between adjacent molecules. For example, compare isobutane and n-butane, which boil at -12 and 0 °C, and 2,2-dimethylbutane and 2,3-dimethylbutane which boil at 50 and 58 °C, respectively.[6] For the latter case, two molecules 2,3- dimethylbutane can "lock" into each other better than the cross-shaped 2,2- dimethylbutane, hence the greater van der Waals forces.

On the other hand, cycloalkanes tend to have higher boiling points than their linear counterparts due to the locked conformations of the molecules, which give a plane of intermolecular contact.

Melting point

The melting points of the alkanes follow a similar trend to boiling points for the same reason as outlined above. That is, (all other things being equal) the larger the molecule the higher the melting point. There is one significant difference between boiling points and melting points. Solids have more rigid and fixed structure than liquids. This rigid structure requires energy to break down. Thus the stronger better put together solid structures will require more energy to break apart. For alkanes, this can be seen from the graph above (i.e., the blue line). The odd-numbered alkanes have a lower trend in melting points than even numbered alkanes. This is because even numbered alkanes pack well in the solid phase, forming a well-organised structure, which requires more energy to break apart. The odd-number alkanes pack less well and so the "looser" organised solid packing structure requires less energy to break apart.

The melting points of branched-chain alkanes can be either higher or lower than those of the corresponding straight-chain alkanes, again depending on the ability of the alkane in question to packing well in the solid phase: This is particularly true for isoalkanes (2-methyl isomers), which often have melting points higher than those of the linear analogues.

Conductivity

Alkanes do not conduct electricity, nor are they substantially polarized by an electric field. For this reason they do not form hydrogen bonds and are insoluble in polar solvents such as water. Alkanes are said to be hydrophobic in that they repel water. Their solubility in nonpolar solvents is relatively good. The density of the alkanes usually increases with increasing number of carbon atoms, but remains

42 less than that of water. Hence, alkanes form the upper layer in an alkane-water mixture.

Reactions with oxygen (combustion reaction)

All alkanes react with oxygen in a combustion reaction, although they become increasingly difficult to ignite as the number of carbon atoms increases. The general equation for complete combustion is:

CnH2n+2 + (1.5n+0.5)O2 → (n+1)H2O + nCO2

In the absence of sufficient oxygen, carbon monoxide or even soot can be formed, as shown below:

CnH(2n+2) + ½ nO2 → (n+1)H2O + nCO for example methane: 2CH4 + 3O2 → 2CO + 4H2O

CH4 + O2 → C + 2H2O

Reactions with halogens

Alkanes react with halogens in a so-called free radical halogenation reaction. The hydrogen atoms of the alkane are progressively replaced by halogen atoms. Free-radicals are the reactive species that participate in the reaction, which usually leads to a mixture of products. The reaction is highly exothermic, and can lead to an explosion.

Cracking: Cracking breaks larger molecules into smaller ones. This can be done with a thermal or catalytic method.

Week No. 13, 14 / ALKENE alkene, olefin, or olefine is an unsaturated chemical compound containing at least one carbon-to-carbon double bond.[1] The simplest acyclic alkenes, with only one double bond and no other functional groups, form a homologous series of hydrocarbons with the general formula CnH2n.

The simplest alkene is ethylene (C2H4), which has the International Union of Pure and Applied Chemistry (IUPAC) name ethene. Alkenes are also called olefins (an archaic synonym, widely used in the petrochemical industry). Aromatic compounds are often drawn as cyclic alkenes, but their structure and properties are different and they are not considered to be alkenes.

43 Alkenes react in many addition reactions, which occur by opening up the double-bond.

 Catalytic hydrogenation: Hydrogenation of alkenes produces the corresponding alkanes. The reaction is carried out under pressure in the presence of a metallic catalyst. Common industrial catalysts are based on platinum, nickel or palladium. For laboratory syntheses, Raney nickel (an alloy of nickel and aluminium)is often employed. The simplest example of this reaction is the catalytic hydrogenation of ethylene to yield ethane:

CH2=CH2 + H2 → CH3-CH3

 elElectrophilic addition: Most addition reactions to alkenes follow the mechanism of ectrophilic addition :

 Electrophilic halogenation: Addition of elemental bromine or chlorine to alkenes yields vicinal dibromo- and dichloroalkanes, respectively. The decoloration of a solution of bromine in water is an analytical test for the presence of alkenes:

CH2=CH2 + Br2 → BrCH2-CH2Br

Oxidation

Alkenes are oxidized with a large number of oxidizing agents.

 In the presence of oxygen, alkenes burn with a bright flame to produce carbon dioxide and water.  Catalytic oxidation with oxygen or the reaction with percarboxylic acids yields epoxides  Reaction with ozone in ozonolysis leads to the breaking of the double bond, yielding two aldehydes or ketones

R1-CH=CH-R2 + O3 → R1-CHO + R2-CHO + H2O This reaction can be used to determine the position of a double bond in an unknown alkene.

44 Polymerization

Polymerization of alkenes is a reaction that yields polymers of high industrial value at great economy, such as the plastics polyethylene and polypropylene. To form the root of the IUPAC names for alkenes, simply change the -an- infix of the parent to -en-. For example, CH3-CH3 is the alkane ethANe. The name of CH2=CH2 is therefore ethENe.

In higher alkenes, where isomers exist that differ in location of the double bond, the following numbering system is used:

1. Number the longest carbon chain that contains the double bond in the direction that gives the carbon atoms of the double bond the lowest possible numbers. 2. Indicate the location of the double bond by the location of its first carbon. 3. Name branched or substituted alkenes in a manner similar to alkanes. 4. Number the carbon atoms, locate and name substituent groups, locate the double bond, and name the main chain.

Naming substituted hex-1-enes

Cis-Trans notation

In the specific case of disubstituted alkenes where the two carbons have one substituent each, Cis-trans notation may be used. If both substituents are on the same side of the bond, it is defined as (cis-). If the substituents are on either side of the bond, it is defined as (trans-).

The difference between cis- and trans- isomers

45 E,Z notation

When an alkene has more than one substituent (especially necessary with 3 or 4 substituents), the double bond geometry is described using the labels E and Z. These labels come from the German words "entgegen," meaning "opposite," and "zusammen," meaning "together." Alkenes with the higher priority groups (as determined by CIP rules) on the same side of the double bond have these groups together and are designated Z. Alkenes with the higher priority groups on opposite sides are designated E. A mnemonic to remember this: Z notation has the higher priority groups on "ze zame zide."

The difference between E and Z isomers

Week No. 15, 16 / ALKYNE

Alkynes are hydrocarbons that have a triple bond between two carbon atoms, with the formula CnH2n-2. Alkynes are traditionally known as acetylenes, although the name acetylene also refers specifically to C2H2, known formally (but rarely) as ethyne using IUPAC nomenclature. Like other hydrocarbons, alkynes are generally hydrophobic but tend to be more reactive.

Alkynes are characteristically more unsaturated than alkenes. Thus they add two equivalents of bromine whereas an alkene adds only one equivalent. Other reactions are listed below. Alkynes are usually more reactive than alkenes. They show greater tendency to polymerize or oligomerize than alkenes do. The resulting polymers, called polyacetylenes (which do not contain alkyne units) are conjugated and can exhibit semiconducting properties.

Addition of hydrogen, halogens, and related reagents

Alkynes characteristically undergo reactions that show that they are "doubly unsaturated," meaning that each alkyne unit is capable of adding two equivalents of H2, halogens or related HX reagents (X = halide, pseudohalide, etc.). Depending on catalysts and conditions, alkynes add one or two equivalents of hydrogen. Hydrogenation to the alkene is usually more desirable since alkanes are less useful:

46 RC≡CR' + H2 → cis-RCH=CR'H

The largest scale application of this technology is the conversion of acetylene to ethylene catalysts in refineries. The steam cracking of alkanes affords a few percent acetylene, which is selectively hydrogenated in the presence of a palladium/silver catalyst. For more complex alkynes, the Lindlar catalyst is widely recommended to avoid formation of the alkane, for example in the conversion of phenylacetylene to styrene.[3]

Similarly, halogenation of alkynes gives the vinyl dihalides or alkyl

tetrahalides: RC≡CR' + 2 Br2 → RCBr2CRBr2

The addition of nonpolar E-H bonds across C≡C is general for silanes, boranes, and related hydrides. The hydroboration of alkynes gives vinylic boranes which oxidize to the corresponding aldehyde or ketone. Hydrohalogenation gives the corresponding vinyl halides or alkyl dihalides, again depending on the number of equivalents of HX added. The addition of water to alkynes is a related reaction except the initial enol intermediate converts to the ketone or aldehyde. Illustrative is the hydration of [4] phenylacetylene gives acetophenone, and the (Ph3P)AuCH3-catalyzed hydration of 1,8-nonadiyne to 2,8-nonanedione:[5]

PhC≡CH + H2O → PhCOCH3

HC≡CC6H12C≡CH + 2H2O → CH3COC6H12COCH3

Cycloadditions and oxidation

Alkynes undergo diverse cycloaddition reactions. Most notable is the Diels- Alder reaction with 1,3-dienes to give 1,4-cyclohexadienes. This general reaction has been extensively developed and electrophilic alkynes are especially effective dienophiles. The "cycloadduct" derived from the addition of alkynes to 2-pyrone eliminates carbon dioxide to give the aromatic compound. Other specialized cycloadditions include multicomponent reactions such as alkyne trimerisation to give aromatic compounds and the [2+2+1]cycloaddition of an alkyne, alkene and carbon monoxide in the Pauson–Khand reaction. Non- carbon reagents also undergo cyclization, e.g. Azide alkyne Huisgen cycloaddition to give triazoles. Cycloaddition processes involving alkynes are often catalyzed by metals, e.g. enyne metathesis and alkyne metathesis, which

47 allows the scrambling of carbyne (RC) centers: RC≡CR + R'C≡CR' 2 RC≡CR'

Oxidative cleavage of alkynes proceeds via cycloaddition to metal oxides. Most famously, potassium permanganate converts alkynes to a pair of carboxylic acids.

THE NAMES OF ORGANIC COMPOUNDS

This lecture explains how to write the formula for an organic compound given its name - and vice versa. It covers alkanes, cycloalkanes, alkenes, simple compounds containing halogens, alcohols, aldehydes and ketones.

There are two skills you have to develop in this area:

 You need to be able to translate the name of an organic compound into its structural formula.  You need to be able to name a compound from its given formula.

Cracking the code

A modern organic name is simply a code. Each part of the name gives you some useful information about the compound.

For example, to understand the name 2-methylpropan-1-ol you need to take the name to pieces.

The prop in the middle tells you how many carbon atoms there are in the longest chain (in this case, 3). The an which follows the "prop" tells you that there aren't any carbon-carbon double bonds.

The other two parts of the name tell you about interesting things which are happening on the first and second carbon atom in the chain. Any name you are likely to come across can be broken up in this same way.

Counting the carbon atoms

You will need to remember the codes for the number of carbon atoms in a chain up to 6 carbons. There is no easy way around this - you have got to learn them. If you don't do this properly, you won't be able to name anything!

code no of carbons

48 meth 1 eth 2 prop 3 but 4 pent 5 hex 6 Hep 7 oct 8 non 9 dec 10 undec 11 dodec 12

Types of carbon-carbon bonds

Whether or not the compound contains a carbon-carbon double bond is shown by the two letters immediately after the code for the chain length.

code means an Contains a carbon-carbon single bonds en contains a carbon-carbon double bond yn contains a carbon-carbon triple bond

For example, butane means four carbons in a chain with no double bond.

Propene means three carbons in a chain with a double bond between two of the carbons.

Molecular and Structural Formula

Molecular formula: A chemical formula for a molecular substance that gives the number and type of each atom present in a molecule of the substance. For example the formula H2O indicates that a molecule of water contains two atoms of hydrogen and one atom of oxygen.

Structural formula: A chemical formula for a molecule in which a separate symbol is used to indicate each atom and dashes are used to show how these atoms are joined together.

49 Alkyl groups

Compounds like methane, CH4, and ethane, CH3CH3, are members of a family of compounds called alkanes. If you remove a hydrogen atom from one of these you get an alkyl group.

For example:

 A methyl group is CH3.  An ethyl group is CH3CH2.

These groups must, of course, always be attached to something else.

Types of compounds

The alkanes

Example 1: Write the structural formula for 2-methylpentane.

Start decoding the name from the bit that counts the number of carbon atoms in the longest chain - pent counts 5 carbons.

Are there any carbon-carbon double bonds? No - an tells you there aren't any.

Now draw this carbon skeleton:

Put a methyl group on the number 2 carbon atom:

Does it matter which end you start counting from? No - if you counted from the other end, you would draw the next structure. That's exactly the same as the first one, except that it has been flipped over.

Finally, all you have to do is to put in the correct number of hydrogen atoms on each carbon so that each carbon is forming four bonds.

50 If you had to name this yourself:

 Count the longest chain of carbons that you can find. Don't assume that you have necessarily drawn that chain horizontally. 5 carbons means pent.  Are there any carbon-carbon double bonds? No - therefore pentane.  There's a methyl group on the number 2 carbon - therefore 2-methylpentane. Why the number 2 as opposed to the number 4 carbon? In other words, why do we choose to number from this particluar end? The convention is that you number from the end which produces the lowest numbers in the name - hence 2- rather than 4-.

Example 2: Write the structural formula for 2,3-dimethylbutane.

Start with the carbon backbone. There are 4 carbons in the longest chain (but) with no carbon-carbon double bonds (an).

This time there are two methyl groups (di) on the number 2 and number 3 carbon atoms.

Completing the formula by filling in the hydrogen atoms gives:

Example 3: Write the structural formula for 2,2-dimethylbutane.

This is exactly like the last example, except that both methyl groups are on the same carbon atom. Notice that the name shows this by using 2,2- as well as di. The structure is worked out as before:

51 Example 4: Write the structural formula for 3-ethyl-2-methylhexane. hexan shows a 6 carbon chain with no carbon-carbon double bonds.

This time there are two different alkyl groups attached - an ethyl group on the number 3 carbon atom and a methyl group on number 2.

Filling in the hydrogen atoms gives:

Note: Once again it doesn't matter whether the ethyl and methyl groups point up or down. You might also have chosen to start numbering from the right-hand end of the chain. These would all be perfectly valid structures. All you would have done is to rotate the whole molecule in space, or rotate it around particular bonds.

If you had to name this yourself:

How do you know what order to write the different alkyl groups at the beginning of the name? The convention is that you write them in alphabetical order - hence ethyl comes before methyl which in turn comes before propyl.

The cycloalkanes

In a cycloalkane the carbon atoms are joined up in a ring - hence cyclo.

Example: Write the structural formula for cyclohexane. hexan shows 6 carbons with no carbon-carbon double bonds. cyclo shows that they are in a ring. Drawing the ring and putting in the correct number of hydrogens to satisfy the bonding

52 requirements of the carbons gives:

The alkenes

Example 1: Write the structural formula for propene. prop counts 3 carbon atoms in the longest chain. en tells you that there is a carbon-carbon double bond. That means that the carbon skeleton looks like this:

Putting in the hydrogens gives you:

Example 2: Write the structural formula for but-1-ene. but counts 4 carbon atoms in the longest chain and en tells you that there is a carbon-carbon double bond. The number in the name tells you where the double bond starts.

No number was necessary in the propene example above because the double bond has to start on one of the end carbon atoms. In the case of butene, though, the double bond could either be at the end of the chain or in the middle - and so the name has to code for the its position.

The carbon skeleton is:

And the full structure is:

Incidentally, you might equally well have decided that the right-hand carbon was the number 1 carbon, and drawn the structure as:

Example 3: Write the structural formula for 3-methylhex-2-ene.

53 The longest chain has got 6 carbon atoms (hex) with a double bond starting on the second one (-2-en).

But this time there is a methyl group attached to the chain on the number 3 carbon atom, giving you the underlying structure:

Adding the hydrogens gives the final structure:

Be very careful to count the bonds around each carbon atom when you put the hydrogens in. It would be very easy this time to make the mistake of writing an H after the third carbon - but that would give that carbon a total of 5 bonds.

Compounds containing halogens

Example 1: Write the structural formula for 1,1,1-trichloroethane.

This is a two carbon chain (eth) with no double bonds (an). There are three chlorine atoms all on the first carbon atom.

Example 2: Write the structural formula for 2-bromo-2-methylpropane.

First sort out the carbon skeleton. It's a three carbon chain with no double bonds and a methyl group on the second carbon atom.

Draw the bromine atom which is also on the second carbon.

And finally put the hydrogen atoms in.

54 If you had to name this yourself:

Notice that the whole of the hydrocarbon part of the name is written together - as methylpropane - before you start adding anything else on to the name.

Example 2: Write the structural formula for 1-iodo-3-methylpent-2-ene.

This time the longest chain has 5 carbons (pent), but has a double bond starting on the number 2 carbon. There is also a methyl group on the number 3 carbon.

Now draw the iodine on the number 1 carbon.

Giving a final structure:

Alcohols

All alcohols contain an -OH group. This is shown in a name by the ending ol.

Example 1: Write the structural formula for methanol.

This is a one carbon chain with no carbon-carbon double bond (obviously!). The ol ending shows it's an alcohol and so contains an -OH group.

Example 2: Write the structural formula for 2-methylpropan-1-ol.

55 The carbon skeleton is a 3 carbon chain with no carbon-carbon double bonds, but a methyl group on the number 2 carbon.

The -OH group is attached to the number 1 carbon.

The structure is therefore:

Example 3: Write the structural formula for ethane-1,2-diol.

This is a two carbon chain with no double bond. The diol shows 2 -OH groups, one on each carbon atom.

Note: There's no particular significance in the fact that this formula has the carbon chain drawn vertically. If you draw it horizontally, unless you stretch the carbon- carbon bond a lot, the -OH groups look very squashed together. Drawing it vertically makes it look tidier!

Aldehydes: All aldehydes contain the group:

If you are going to write this in a condensed form, you write it as -CHO - never as - COH, because that looks like an alcohol. The names of aldehydes end in al.

Example 1: Write the structural formula for propanal.

56 This is a 3 carbon chain with no carbon-carbon double bonds. The al ending shows the presence of the -CHO group. The carbon in that group counts as one of the chain.

Example 2: Write the structural formula for 2-methylpentanal. This time there are 5 carbons in the longest chain, including the one in the -CHO group. There aren't any carbon-carbon double bonds. A methyl group is attached to the number 2 carbon. Notice that in aldehydes, the carbon in the -CHO group is always counted as the number 1 carbon.

Ketones:

Ketones contain a carbon-oxygen double bond just like aldehydes, but this time it's in the middle of a carbon chain. There isn't a hydrogen atom attached to the group as there is in aldehydes. Ketones are shown by the ending one.

Example 1: Write the structural formula for propanone.

This is a 3 carbon chain with no carbon-carbon double bond. The carbon-oxygen double bond has to be in the middle of the chain and so must be on the number 2 carbon.

Ketones are often written in this way to emphasise the carbon-oxygen double bond.

Example 2: Write the structural formula for pentan-3-one.

This time the position of the carbon-oxygen double bond has to be stated because there is more than one possibility. It's on the third carbon of a 5 carbon chain with no carbon- carbon double bonds. If it was on the second carbon, it would be pentan-2-one.

57 This could equally well be written:

Carboxylic acids

Carboxylic acids contain the -COOH group, which is better written out in full as:

Carboxylic acids are shown by the ending oic acid. When you count the carbon chain, you have to remember to include the carbon in the -COOH group. That carbon is always thought of as number 1 in the chain.

Example 1: Write the structural formula for 3-methylbutanoic acid.

This is a four carbon acid with no carbon-carbon double bonds. There is a methyl group on the third carbon (counting the -COOH carbon as number 1).

Example 2: Write the structural formula for 2-hydroxypropanoic acid.

The hydroxy part of the name shows the presence of an -OH group. Normally, you would show that by the ending ol, but this time you can't because you've already got another ending. You are forced into this alternative way of describing it.

The old name for 2-hydroxypropanoic acid is lactic acid. That name sounds more friendly, but is utterly useless when it comes to writing a formula for it. In the old days, you would have had to learn the formula rather than just working it out should you need

58 it.

Example 3: Write the structural formula for 2-chlorobut-3-enoic acid.

This time, not only is there a chlorine attached to the chain, but the chain also contains a carbon-carbon double bond (en) starting on the number 3 carbon (counting the -COOH carbon as number 1).

Salts of carboxylic acids

Example: Write the structural formula for sodium propanoate.

This is the sodium salt of propanoic acid - so start from that. Propanoic acid is a three carbon acid with no carbon-carbon double bonds.

When the carboxylic acids form salts, the hydrogen in the -COOH group is replaced by a metal. Sodium propanoate is therefore:

Notice that there is an ionic bond between the sodium and the propanoate group. Whatever you do, don't draw a line between the sodium and the oxygen. That would represent a covalent bond. It's wrong, and makes you look very incompetent in an exam!

In a shortened version, sodium propanoate would be written CH3CH2COONa or, if you wanted to emphasise the ionic nature, as CH3CH2COO- Na+.

Note: The confusing thing about these salts (and even more so for the esters that are coming up next) is that they are named the wrong way round. In the formula, the sodium is at the end, but appears first in the name. Why?

59 Salts are always named with the metal first - think of sodium chloride or potassium iodide. So for consistency you would need to reverse the formula of

sodium propanoate - NaOOCCH2CH3. But if you reverse the formula, you can't see immediately that it is related to propanoic acid. So you learn to live with the inconsistency.

Esters

Esters are one of a number of compounds known collectively as acid derivatives. In these the acid group is modified in some way. In an ester, the hydrogen in the -COOH group is replaced by an alkyl group (or possibly some more complex hydrocarbon group).

Example 1: Write the structural formula for methyl propanoate.

An ester name has two parts - the part that comes from the acid (propanoate) and the part that shows the alkyl group (methyl).

Start by thinking about propanoic acid - a 3 carbon acid with no carbon-carbon double bonds.

The hydrogen in the -COOH group is replaced by an alkyl group - in this case, a methyl group.

Ester names are confusing because the name is written backwards from the way the structure is drawn. There's no way round this - you just have to get used to it!

In the shortened version, this formula would be written CH3CH2COOCH3.

Example 2: Write the structural formula for ethyl ethanoate.

60 This is probably the most commonly used example of an ester. It is based on ethanoic acid ( hence, ethanoate) - a 2 carbon acid. The hydrogen in the -COOH group is replaced by an ethyl group.

Make sure that you draw the ethyl group the right way round. A fairly common mistake is to try to join the CH3 group to the oxygen. If you count the bonds if you do that, you will find that both the CH3 carbon and the CH2 carbon have the wrong number of bonds.

Acyl chlorides (acid chlorides)

An acyl chloride is another acid derivative. In this case, the -OH group of the acid is replaced by -Cl. All acyl chlorides contain the -COCl group:

Example: Write the structural formula for ethanoyl chloride.

Acyl chlorides are shown by the ending oyl chloride. So ethanoyl chloride is based on a 2 carbon chain with no carbon-carbon double bonds and a -COCl group. The carbon in that group counts as part of the chain. In a longer chain, with side groups attached, the -COCl carbon is given the number 1 position.

Acid anhydrides

Another acid derivative! An acid anhydride is what you get if you dehydrate an acid - that is, remove water from it.

Example: Write the structural formula for propanoic anhydride.

These are most easily worked out by writing it down on a scrap of paper in the following

61 way:

Draw two molecules of acid arranged so that the -OH groups are next to each other. Tweak out a molecule of water - and then join up what's left. In this case, because you want propanoic anhydride, you draw two molecules of propanoic acid.

Amides

Yet another acid derivative! Amides contain the group -CONH2 where the -OH of an acid is replaced by -NH2.

Example: Write the structural formula for propanamide.

This is based on a 3 carbon chain with no carbon-carbon double bonds. At the end of the chain is a -CONH2 group. The carbon in that group counts as part of the chain.

Nitriles

Nitriles contain a -CN group, and used to be called cyanides.

Example 1: Write the structural formula for ethanenitrile.

The name shows a 2 carbon chain with no carbon-carbon double bond. nitrile shows a -CN group at the end of the chain. As with the previous examples involving acids and acid derivatives, don't forget that the carbon in the -CN group counts as part of the chain.

The old name for this would have been methyl cyanide. You might think that that's easier, but as soon as the chain gets more complicated, it doesn't work - as the next example

62 shows.

Example 2: Write the structural formula for 2-hydroxypropanenitrile.

Here we've got a 3 carbon chain, no carbon-carbon double bonds, and a -CN group on the end of the chain. The carbon in the -CN group counts as the number 1 carbon. On the number 2 carbon there is an -OH group (hydroxy). Notice that you can't use the ol ending because you've already got a nitrile ending.

Primary amines

A primary amine contains the group -NH2 attached to a hydrocarbon chain or ring. You can think of amines in general as being derived from ammonia, NH3. In a primary amine, one of the hydrogens has been replaced by a hydrocarbon group.

Example 1: Write the structural formula for ethylamine.

In this case, an ethyl group is attached to the -NH2 group.

This name (ethylamine) is fine as long as you've only got a short chain where there isn't any ambiguity about where the -NH2 group is found. But suppose you had a 3 carbon chain

- in this case, the -NH2 group could be on an end carbon or on the middle carbon. How you get around that problem is illustrated in the next example.

Example 2: Write the structural formula for 2-aminopropane.

The name shows a 3 carbon chain with an amino group attached to the second carbon.

amino shows the -NH2 group.

Ethylamine (example 1 above) could equally well have been called aminoethane.

63 Secondary and tertiary amines

You are only likely to come across simple examples of these. In a secondary amine, two of the hydrogen atoms in an ammonia molecule have been replaced by hydrocarbon groups. In a tertiary amine, all three hydrogens have been replaced.

Example 1: Write the structural formula for dimethylamine.

In this case, two of the hydrogens in ammonia have been replaced by methyl groups.

Example 2: Write the structural formula for trimethylamine.

Here, all three hydrogens in ammonia have been replaced by methyl groups.

Amino acids

An amino acid contains both an amino group, -NH2, and a carboxylic acid group, -COOH, in the same molecule. As with all acids the carbon chain is numbered so that the carbon in the -COOH group is counted as number 1.

Example: Write the structural formula for 2-aminopropanoic acid.

This has a 3 carbon chain with no carbon-carbon double bonds. On the second carbon

(counting the -COOH carbon as number 1) there is an amino group, -NH2.

64 THE NAMES OF AROMATIC COMPOUNDS

This page looks at the names of some simple aromatic compounds. An aromatic compound is one which contains a benzene ring. It assumes that you are reasonably confident about naming compounds containing chains of carbon atoms (aliphatic compounds).

Naming aromatic compounds isn't quite so straightforward as naming chain compounds. Often, more than one name is acceptable and it's not uncommon to find the old names still in use as well.

Background

The benzene ring

All aromatic compounds are based on benzene, C6H6, which has a ring of six carbon atoms and has the symbol:

Each corner of the hexagon has a carbon atom with a hydrogen attached.

The phenyl group

Remember that you get a methyl group, CH3, by removing a hydrogen from methane, CH4.

You get a phenyl group, C6H5, by removing a hydrogen from a benzene ring, C6H6. Like a methyl or an ethyl group, a phenyl group is always attached to something else.

Aromatic compounds with only one group attached to the benzene ring

Cases where the name is based on benzene chlorobenzene

This is a simple example of a halogen attached to the benzene ring. The name is self-

65 obvious.

The simplified formula for this is C6H5Cl. You could therefore (although you never do!) call it phenyl chloride. Whenever you draw a benzene ring with one other thing attached to it, you are in fact drawing a phenyl group. In order to attach something else, you have to remove one of the existing hydrogen atoms, and so automatically make a phenyl group. nitrobenzene

The nitro group, NO2, is attached to a benzene ring.

The simplified formula for this is C6H5NO2. methylbenzene

Another obvious name - the benzene ring has a methyl group attached. Other alkyl side- chains would be named similarly - for example, ethylbenzene. The old name for methylbenzene is toluene, and you may still meet that.

The simplified formula for this is C6H5CH3.

(chloromethyl)benzene

A variant on this which you may need to know about is where one of the hydrogens on the

CH3 group is replaced by a chlorine atom. Notice the brackets around the (chloromethyl) in the name. This is so that you are sure that the chlorine is part of the methyl group and not somewhere else on the ring.

66 If more than one of the hydrogens had been replaced by chlorine, the names would be (dichloromethyl)benzene or (trichloromethyl)benzene. Again, notice the importance of the brackets in showing that the chlorines are part of the side group and not directly attached to the ring. benzoic acid (benzenecarboxylic acid)

Benzoic acid is the older name, but is still in common use - it's a lot easier to say and write than the modern alternative! Whatever you call it, it has a carboxylic acid group, -COOH, attached to the benzene ring.

Cases where the name is based on phenyl

Remember that the phenyl group is a benzene ring minus a hydrogen atom - C6H5. If you draw a benzene ring with one group attached, you have drawn a phenyl group. phenylamine

Phenylamine is a primary amine and contains the -NH2 group attached to a benzene ring.

The old name for phenylamine is aniline, and you could also reasonably call it aminobenzene.

phenylethene

This is an ethene molecule with a phenyl group attached. Ethene is a two carbon chain with

67 a carbon-carbon double bond. Phenylethene is therefore:

The old name for phenylethene is styrene - the monomer from which polystyrene is made. phenylethanone

This is a slightly awkward name - take it to pieces. It consists of a two carbon chain with no carbon-carbon double bond. The one ending shows that it is a ketone, and so has a C=O group somewhere in the middle. Attached to the carbon chain is a phenyl group. Putting that together gives:

phenyl ethanoate

This is an ester based on ethanoic acid. The hydrogen atom in the -COOH group has been replaced by a phenyl group.

Note: If you aren't happy about naming esters, follow this link before you go on. phenol

Phenol has an -OH group attached to a benzene ring and so has a formula C6H5OH.

68 Aromatic compounds with more than one group attached to the benzene ring

Numbering the ring

Any group already attached to the ring is given the number 1 position. Where you draw it on the ring (at the top or in any other position) doesn't matter - that's just a question of rotating the molecule a bit. It's much easier, though, to get in the habit of drawing your main group at the top.

The other ring positions are then numbered from 2 to 6. You can number them either clockwise or anti-clockwise. As with chain compounds, you number the ring so that the name you end up with has the smallest possible numbers in it. Examples will make this clear.

Some simple examples

Substituting chlorine atoms on the ring

Look at these compounds:

All of these are based on methylbenzene and so the methyl group is given the number 1 position on the ring.

Why is it 2-chloromethylbenzene rather than 6-chloromethylbenzene? The ring is numbered clockwise in this case because that produces a 2- in the name rather than a 6-. 2 is smaller than 6. Warning! You will find all sorts of variations on this depending on the age of the book you look it up in, and where it was published. What I have described above isn't in strict accordance with the most modern interpretation of the

69 IUPAC recommendations for naming organic compounds.

The names should actually be 1-chloro-2-methylbenzene, 1-chloro-3- methylbenzene, and so on. The substituted groups are named in alphabetical order, and the "1" position is assigned to the first of these - rather than to the more logical methyl group.

This produces some silly inconsistencies. For example, if you had the exactly equivalent compounds containing nitro groups in place of the chlorines, the names would change completely, to 1-methyl-2-nitrobenzene, 1-methyl-3- nitrobenzene, etc. In this case, the normal practice of naming the hydrocarbon first, and then attaching other things to it has been completely wrecked.

Do you need to worry about this? NO! It is extremely unlikely that you would ever be asked to name these in an exam, and it is always easy to write a structure from one of these names - however illogical it may be! There is a simple rule for exam purposes. Unless you are specifically asked for the name of anything remotely complicated, don't give it. As long as you have got the structure right, that's all that matters. 2-hydroxybenzoic acid

This might also be called 2-hydroxybenzenecarboxylic acid. There is a -COOH group attached to the ring and, because the name is based on benzoic acid, that group is assigned the number 1 position. Next door to it in the 2 position is a hydroxy group, -OH.

benzene-1,4-dicarboxylic acid

The di shows that there are two carboxylic acid groups, -COOH, one of them in the 1 position and the other opposite it in the 4 position.

70 2,4,6-trichlorophenol

This is based on phenol - with an -OH group attached in the number 1 position on the ring. There are 3 chlorine atoms substituted onto the ring in the 2, 4 and 6 positions.

methyl 3-nitrobenzoate

This is a name you might come across as a part of a practical exercise in nitrating benzene rings. It's included partly for that reason, and partly because it is a relatively complicated name to finish with!

The structure of the name shows that it is an ester. You can tell that from the oate ending, and the methyl group floating separately from the rest of the name at the beginning. The ester is based on the acid, 3-nitrobenzoic acid - so start with that. There will be a benzene ring with a -COOH group in the number 1 position and a nitro group, NO2, in the 3 position. The -COOH group is modified to make an ester by replacing the hydrogen of the -COOH group by a methyl group. Methyl 3-nitrobenzoate is therefore:

Week No. 17 / ALCOHOL

71 an alcohol is any organic compound in which a hydroxyl group (-OH) is bound to a carbon atom of an alkyl or substituted alkyl group. An important group of alcohols is formed by the simple acyclic alcohols, the general formula for which is CnH2n+1OH. Of those, ethanol (C2H5OH) is the type of alcohol found in alcoholic beverages, and in common speech the word alcohol means, specifically, ethanol.

Other alcohols are usually described with a clarifying adjective, as in isopropyl alcohol (propan-2-ol) or wood alcohol (methyl alcohol, or methanol). The suffix -ol appears in the IUPAC chemical name of all alcohols.

There are three major subsets of alcohols: primary (1°), secondary (2°) and tertiary (3°), based upon the number of carbon atoms the C-OH group's carbon (shown in red) is bonded to. Ethanol is a simple 'primary' alcohol. The simplest secondary alcohol is isopropyl alcohol (propan-2-ol), and a simple tertiary alcohol is tert-butyl alcohol (2-methylpropan-2-ol).

Simple alcohols

The simplest and most commonly used alcohols are methanol and ethanol. Methanol was formerly obtained by the distillation of wood and called "wood alcohol."

Two other alcohols whose uses are relatively widespread (though not so much as those of methanol and ethanol) are propanol and butanol. Like ethanol, they can be produced by fermentation processes. (However, the fermenting agent is a bacterium, Clostridium acetobutylicum, that feeds on cellulose, not sugars like the Saccharomyces yeast that produces ethanol.)

Nomenclature

Systematic names

In the IUPAC system, the name of the alkane chain loses the terminal "e" and adds "ol", e.g. "methanol" and "ethanol". When necessary, the position of the hydroxyl group is indicated by a number between the alkane name and the "ol": propan-1-ol for CH3CH2CH2OH, propan-2-ol for CH3CH(OH)CH3. Sometimes, the position number is written before the IUPAC name: 1-propanol and 2- propanol. If a higher priority group is present (such as an aldehyde, ketone or carboxylic acid), then it is necessary to use the prefix "hydroxy", for example: 1-hydroxy-2-propanone (CH3COCH2OH).

72 Physical and chemical properties

Alcohols have an odor that is often described as “biting” and as “hanging” in the nasal passages.

The hydroxyl group generally makes the alcohol molecule polar. Two opposing solubility trends in alcohols are: the tendency of the polar OH to promote solubility in water, and of the carbon chain to resist it. Thus, methanol, ethanol, and propanol are miscible in water because the hydroxyl group wins out over the short carbon chain. Butanol, with a four-carbon chain, is moderately soluble because of a balance between the two trends. Alcohols of five or more carbons (Pentanol and higher) are effectively insoluble in water because of the hydrocarbon chain's dominance. All simple alcohols are miscible in organic solvents.

Because of hydrogen bonding, alcohols tend to have higher boiling points than comparable hydrocarbons and ethers. The boiling point of the alcohol ethanol is 78.29 °C, compared to 69 °C for the hydrocarbon Hexane (a common constituent of gasoline), and 34.6 °C for Diethyl ether.

Alcohols, like water, can show either acidic or basic properties at the O-H group. With a pKa of around 16-19 they are generally slightly weaker acids than water.

Mean while the oxygen atom has lone pairs of nonbonded electrons that render it weakly basic in the presence of strong acids such as sulfuric acid. For example, with methanol:

73 Alcohols can also undergo oxidation to give aldehydes, ketones or carboxylic acids, or they can be dehydrated to alkenes. They can react to form ester compounds, and they can (if activated first) undergo nucleophilic substitution reactions. The lone pairs of electrons on the oxygen of the hydroxyl group also makes alcohols nucleophiles. For more details see the reactions of alcohols section below

Production

Industrially alcohols are produced in several ways:

 By fermentation using glucose produced from sugar from the hydrolysis of starch, in the presence of yeast and temperature of less than 37°C to produce ethanol. For instance the conversion of invertase to glucose and fructose or the conversion of glucose to zymase and ethanol.  By direct hydration using ethylene (ethylene hydration)[5] or other alkenes from cracking of fractions of distilled crude oil.

Laboratory synthesis

Several methods exist for the preparation of alcohols in the laboratory.

 Primary alkyl halides react with aqueous NaOH or KOH mainly to primary alcohols in nucleophilic aliphatic substitution. (Secondary and especially tertiary alkyl halides will give the elimination (alkene) product instead).  Aldehydes or ketones are reduced with sodium borohydride or lithium aluminium hydride (after an acidic workup). Another reduction by aluminiumisopropylates is the Meerwein-Ponndorf-Verley reduction.  Alkenes engage in an acid catalysed hydration reaction using concentrated sulfuric acid as a catalyst which gives usually secondary or tertiary alcohols. The hydroboration-oxidation and oxymercuration- reduction of alkenes are more reliable in organic synthesis. Alkenes react with NBS and water in halohydrin formation reaction

74  Grignard reagents react with carbonyl groups to secondary and tertiary alcohols. Related reactions are the Barbier reaction and the Nozaki- Hiyama reaction.  Noyori asymmetric hydrogenation is the asymmetric reduction of β-keto- esters  Amines can be converted to diazonium salts which are then hydrolyzed.

The formation of a secondary alcohol via reduction and hydration is shown:

Reactions

Deprotonation

Alcohols can behave as weak acids, undergoing deprotonation. The deprotonation reaction to produce an alkoxide salt is either performed with a strong base such as sodium hydride or n-butyllithium, or with sodium or potassium metal.

- + 2 R-OH + 2 NaH → 2 R-O Na + 2H2↑

− 2 R-OH + 2Na → 2R-O Na + H2

− E.g. 2 CH3CH2-OH + 2 Na → 2 CH3-CH2-O Na + H2

Water is similar in pKa to many alcohols, so with sodium hydroxide there is an equilibrium set up which usually lies to the left:

- + R-OH + NaOH <=> R-O Na + H2O (equilibrium to the left) alcohols may likewise be converted to alkyl bromides using hydrobromic acid or phosphorus tribromide, for example:

3 R-OH + PBr3 → 3 RBr + H3PO3

Dehydration

+ Alcohols are themselves nucleophilic, so R−OH2 can react with ROH to produce ethers and water in a dehydration reaction, although this reaction is rarely used except in the manufacture of diethyl ether.

75 More useful is the E1 elimination reaction of alcohols to produce alkenes. The reaction generally obeys Zaitsev's Rule, which states that the most stable (usually the most substituted) alkene is formed. Tertiary alcohols eliminate easily at just above room temperature, but primary alcohols require a higher temperature.

This is a diagram of acid catalysed dehydration of ethanol to produce ethene:

A more controlled elimination reaction is the Chugaev elimination with carbon disulfide and iodomethane.

Esterification

To form an ester from an alcohol and a carboxylic acid the reaction, known as Fischer esterification, is usually performed at reflux with a catalyst of concentrated sulfuric acid:

R-OH + R'-COOH → R'-COOR + H2O

In order to drive the equilibrium to the right and produce a good yield of ester, water is usually removed, either by an excess of H2SO4 or by using a Dean- Stark apparatus. Esters may also be prepared by reaction of the alcohol with an acid chloride in the presence of a base such as pyridine.

Oxidation Main article: Oxidation of primary alcohols to carboxylic acids

Primary alcohols (R-CH2-OH) can be oxidized either to aldehydes (R-CHO) or to carboxylic acids (R-CO2H), while the oxidation of secondary alcohols (R1R2CH-OH) normally terminates at the ketone (R1R2C=O) stage. Tertiary alcohols (R1R2R3C-OH) are resistant to oxidation.

The direct oxidation of primary alcohols to carboxylic acids normally proceeds via the corresponding aldehyde, which is transformed via an aldehyde hydrate (R-CH(OH)2) by reaction with water before it can be further oxidized to the carboxylic acid.

76 Mechanism of oxidation of primary alcohols to carboxylic acids via aldehydes and aldehyde hydrates

 Catalytic TEMPO in presence of excess bleach (NaOCl) (Anelli's oxidation).

Week No. 18 / ETHER

The general structure for an ether

Ether is a class of organic compounds that contain an ether group — an oxygen atom connected to two alkyl or aryl groups — of general formula R–O–R.[1] A typical example is the solvent and anesthetic diethyl ether, commonly referred to simply as "ether" (CH3-CH2-O-CH2-CH3). Ethers are common in organic chemistry and pervasive in biochemistry, as they are common linkages in carbohydrates and lignin.

77 Physical properties

Ether molecules cannot form hydrogen bonds amongst each other, resulting in a relatively low boiling point compared to that of the analogous alcohols. The difference, however, in the boiling points of the ethers and their isometric alcohols become smaller as the carbon chains become longer, as the van der waals interactions of the extended carbon chain dominate over the presence of hydrogen bonding.

Ethers are slightly polar, as the COC bond angle in the functional group is about 110 degrees, and the C-O dipoles do not cancel out. Ethers are more polar than alkenes but not as polar as alcohols, esters, or amides of comparable structure. However, the presence of two lone pairs of electrons on the oxygen atoms makes hydrogen bonding with water molecules possible, causing the solubility of alcohols (for instance, butan-1-ol) and ethers (ethoxyethane) to be quite dissimilar.

Cyclic ethers such as tetrahydrofuran and 1,4-dioxane are miscible in water because of the more exposed oxygen atom for hydrogen bonding as compared to aliphatic ethers.

Reactions

Structure of the polymeric diethyl ether peroxide

Ethers in general are of low chemical reactivity, but they are more reactive than alkanes (epoxides, ketals, and acetals are unrepresentative classes of ethers and are discussed in separate articles). Important reactions are listed below.[2]

Ether cleavage

Although ethers resist hydrolysis, they are cleaved by mineral acids such as hydrobromic acid and hydroiodic acid. Hydrogen chloride cleaves ethers only slowly. Methyl ethers typically afford methyl halides:

ROCH3 + HBr → CH3Br + ROH

+ - These reactions proceed via onium intermediates, i.e. [RO(H)CH3] Br .

78 Peroxide formation

Primary and secondary ethers with a CH group next to the ether oxygen form peroxides, e.g. diethyl ether peroxide. The reaction requires oxygen (or air) and is accelerated by light, metal catalysts, and aldehydes. The resulting peroxides can be explosive. For this reason, diisopropyl ether and THF are often avoided as solvents.

Ethers can be prepared in the laboratory in several different ways.

Dehydration of alcohols

The Dehydration of alcohols affords ethers:

2 R-OH → R-O-R + H2O

This direct reaction requires elevated temperatures (about 125 °C). The reaction is catalyzed by acids, usually sulfuric acid. The method is effective for generating symmetrical ethers, but not unsymmetrical ethers. Diethyl ether is produced from ethanol by this method. Cyclic ethers are readily generated by this approach. Such reactions must compete with dehydration of the alcohol:

R-CH2-CH2(OH) → R-CH=CH2 + H2O

The dehydration route often requires conditions incompatible with delicate molecules. Several milder methods exist to produce ethers.

Week No.19 / Phenols

Phenol

79 Hydroxybenzene

Properties

Molecular C6H6O formula

Molar mass 94.11 g mol−1

Appearance White Crystalline Solid

Density 1.07 g/cm³

Melting point 40.5 °C, 314 K, 105 °F Boiling point 181.7 °C, 455 K, 359 °F Solubility in 8.3 g/100 ml (20 °C) water

Acidity (pKa) 9.95 Dipole 1.7 D moment

Phenol, also known as carbolic acid, is an organic compound with the chemical formula C6H5OH. It is a white, crystalline solid. This functional group consists of a phenyl, bonded to a hydroxyl (-OH). It is produced on a large scale (about 7 billion kg/year) as a precursor to many materials and useful compounds.[1] It is a mildly acidic compound that requires careful handling.



Properties

Phenol is appreciably soluble in water, with about 8.3 g dissolving in 100 ml (0.88 M). The sodium salt of phenol, sodium phenoxide, is far more water soluble. It is a reactive molecule.

80 Acidity

It is slightly acidic: the phenol molecule has weak tendencies to lose the H+ ion from the hydroxyl group, resulting in the highly water-soluble phenolate anion − [2] C6H5O , called phenoxide anion. Compared to aliphatic alcohols, phenol shows much higher acidity (about 1 million times more acidic). It reacts completely with aqueous NaOH to lose H+, whereas most alcohols react only partially. Phenols are less acidic than carboxylic acids.

Phenoxide anion

Phenol can be deprotonated with moderate base such as triethylamine, forming the nucleophilic phenoxide anion or phenolate anion, which is highly water- soluble.

Reactions

Phenol is highly reactive toward electrophilic aromatic substitution as the oxygen atom's pi electrons donate electron density into the ring. By this general approach, many groups can be appended to the ring, via halogenation, acylation, sulfonation, and other processes.

Production

Phenol can be made from the partial oxidation of benzene, by the cumene process, or by the Raschig-Hooker process. It can also be found as a product of coal oxidation. The dominant method starts from cumene (isopropylbenzene):[1]

C6H5CH(CH3)2 + O2 → C6H5OH + (CH3)2CO Uses

The major uses of phenol involve its conversion to plastics or related materials. Condensation with acetone gives bisphenol-A, a key building block for polycarbonates. Condensation with formaldehyde gives phenolic resins, the most famous of which is Bakelite. Hydrogenation of phenol gives cyclohexanone, an intermediate en route to nylon. Nonionic detergents are

81 produced by alkylation of phenol to give the alkylphenols, which are then subjected to ethoxylation.[1]

Phenol is also a versatile precursor to a large collection of drugs, most notably aspirin but also many herbicides and pharmaceuticals.

Niche uses

Phenol is the preferred chemical for embalming bodies for study because of its ability to preserve tissues for extended periods of time. However, formaldehyde is usually preferred over phenol for embalming with intent of public viewing because of phenol's tendency to turn tissues an unpleasant bleach-white color.

Phenol is also used in the preparation of cosmetics including sunscreens,[8] hair dyes, and skin lightening preparations.[9]

In cosmetic surgery, phenol serves as an exfoliant. It is also used in phenolization, a surgical procedure used to treat an ingrown nail, in which it is applied to the nail bed to prevent regrowth of nails. 5% Phenol is sometimes injected near a sensory nerve in order to temporarily (up to a year) stop it from transmitting impulses in some intractable cases of chronic neuropathic pain.

In molecular biology, phenol is used with chloroform to dissolve proteins and cellular debris when purifying nucleic acids, in particular plasmid preparation.

Toxicity

Phenol and its vapors are corrosive to the eyes, the skin, and the respiratory tract. Repeated or prolonged skin contact with phenol may cause dermatitis, or even second and third-degree burns due to phenol's caustic and defatting properties. Inhalation of phenol vapor may cause lung edema.[14] The substance may cause harmful effects on the central nervous system and heart, resulting in dysrhythmia, seizures, and coma. The kidneys may be affected as well. Exposure may result in death and the effects may be delayed. Long-term or repeated exposure of the substance may have harmful effects on the liver and kidneys."[17] There is no evidence to believe that phenol causes cancer in humans.[18] Besides its hydrophobic effects, another mechanism for the toxicity of phenol may be the formation of phenoxyl radicals.[19]

Chemical burns from skin exposures can be decontaminated by washing with polyethylene glycol,[20] isopropyl alcohol,[21] or perhaps even copious amounts of water.[22] Removal of contaminated clothing is required, as well as immediate

82 hospital treatment for large splashes. This is particularly important if the phenol is mixed with chloroform (a commonly-used mixture in molecular biology for DNA & RNA purification from proteins).

Week No. 20, 21 / KETONE & ALDEHYDE

Ketone

In organic chemistry, a ketone (pronounced /ˈkiːtoʊn/ KEE-toan) is a type of compound that features one carbonyl group (C=O) bonded to two other carbon atoms, i.e., R3CCO-CR3 where R can be a variety of atoms and groups of atoms.[1] With carbonyl carbon bonded to two carbon atoms, ketones are distinct from many other functional groups, such as carboxylic acids, aldehydes, esters, amides, and other oxygen-containing compounds. The double-bond of the carbonyl group distinguishes ketones from alcohols and ethers.

A carbon atom adjacent to a carbonyl group is called an α-carbon. Hydrogen atoms attached to these α-carbon centers are called α-hydrogens. Ketones with α-hydrogen centers participate in a so-called keto-enol tautomerism. The reaction with a strong base gives the corresponding enolate, often by deprotonation of the enol.

Classes of ketones

Ketones are classified on the basis of their substituents. One broad classification subdivides ketones into symmetrical and unsymmetrical derivatives, depending on the equivalency of the two organic substituents attached to the carbonyl center. Acetone and benzophenone are symmetrical ketones. Acetophenone (C6H5C(O)CH3) is an unsymmetrical ketone. In the area of stereochemistry, unsymmetrical ketones are known for being prochiral.

83 Diketones

Many kinds of diketones are known, some with unusual properties. The simplest is biacetyl (CH3C(O)C(O)CH3), once used as butter-flavoring in popcorn. Acetylacetone (pentane-2,4-dione) is virtually a misnomer (inappropriate name) because this species exists mainly as the monoenol CH3C(O)CH=C(OH)CH3. Its enolate is a common ligand in coordination chemistry.

Unsaturated ketones

Ketones containing alkene and alkyne units are often called unsaturated ketones. The most widely used member of this class of compounds is methyl vinyl ketone, CH3C(O)CH=CH2, which is useful in Robinson annulation reaction.

Cyclic ketones

Many ketones are cyclic. The simplest class have the formula (CH2)nCO where n varies from 3 for cyclopropanone to the teens. Larger derivatives exist. Cyclohexanone, a symmetrical cyclic ketone, is an important intermediate in the production of nylon.

Many methods exist for the preparation of ketones in industrial scale, biology, and in academic laboratories. In industry, the most important method probably involves oxidation of hydrocarbons. For example, billion kilograms of cyclohexanone are produced annually by aerobic oxidation of cyclohexane. The oxidation of secondary alcohols is also common:

H3C-CH(OH)-CH3 + O → H3C-CO-CH3 + H2O

Typically, such oxidations employ air or oxygen for industrial processes. For specialized applications, such reactions rely on a strong oxidant such as potassium permanganate or a Cr(VI) compound. Milder conditions such as use of the Dess-Martin periodinane or the Moffatt-Swern oxidation are commonly employed in organic synthesis.

Many other methods have been developed including:

 Ketones can be prepared by geminal halide hydrolysis.  Alkynes can be converted to enols through a hydration reaction in the presence of an acid and HgSO4, and subsequent enol-keto tautomerization gives a ketone. This reaction always produces a ketone, even with a terminal alkyne.

84  Aromatic ketones can be prepared in the Friedel-Crafts acylation, the related Houben-Hoesch reaction and the Fries rearrangement.

Reactions

Ketones engage in many organic reactions but the most important reactions follow from the susceptibility of the carbonyl carbon toward nucleophilic addition and the tendency for the enolates to add to electrophiles. Nucleophilic additions include in approximate order of their generality:

 reaction with water (hydration) gives geminal diols  the reaction with the anion of a terminal alkyne gives a hydroxyalkyne  the reaction with ammonia or a primary amine gives an imine + water  the reaction with secondary amine gives an enamine + water  the reaction with a Grignard reagent or organolithium reagent to give, after aqueous workup, a tertiary alcohol  the reaction with an alcohol, an acid or base gives a hemiketal and water and further reaction with an alcohol gives the ketal + water. This is a carbonyl-protecting reaction.  reaction of RCOR' with sodium amide results in cleavage with formation of the amide RCONH2 and the alkane R'H, a reaction called the Haller- Bauer reaction.[3]  Electrophilic addition, reaction with an electrophile gives a resonance stabilized cation.  the reaction with phosphonium ylides in the Wittig reaction gives alkenes  reaction with thiols gives a thioacetal  reaction with hydrazine or derivatives of hydrazine gives hydrazones  reaction with a metal hydride gives a metal alkoxide salt and then with water an alcohol  reaction of an enol with halogens to form α-haloketone  a reaction at an α-carbon is the reaction of a ketone with heavy water to give a deuterated ketone-d.  fragmentation in photochemical Norrish reaction  reaction with halogens and base of methyl ketones in the Haloform reaction  reaction of 1,4-aminodiketones to oxazoles by dehydration in the Robinson-Gabriel synthesis  reaction of aryl alkyl ketones with sulfur and an amine to amides in the Willgerodt reaction

85 ALDEHYDE

An aldehyde is an organic compound containing a terminal carbonyl group. This functional group, which consists of a carbon atom bonded to a hydrogen atom and double-bonded to an oxygen atom (chemical formula O=CH-), is called the aldehyde group. The aldehyde group is also called the formyl or methanoyl group.

The aldehyde group is polar. Oxygen, more electronegative than carbon, pulls the electrons in the carbon-oxygen bond towards itself, creating an electron deficiency at the carbon atom.

There are several methods for preparing aldehydes:

 Reacting a primary alcohol with a chromium (VI) oxidizing agent. In the laboratory, this may be achieved by heating the alcohol with an acidified solution of potassium dichromate. In this case, excess dichromate will further oxidize the aldehyde to a carboxylic acid, so either the aldehyde is distilled out as it forms (if volatile) or milder reagents such as PCC are used. The reaction is illustrated below with propan-1-ol being oxidised to form propionaldehyde, and again with pentan-1-ol being oxidized to form pentanal.

CH3CH2CH2OH —→ CH3CH2CHO

 Oxidation of primary alcohols can form aldehydes and can be achieved under milder, chromium-free conditions by employing methods or reagents such as IBX acid, Dess-Martin periodinane, Swern oxidation, TEMPO, or the Oppenauer oxidation.  Ozonolysis of non-fully-substituted alkenes yield aldehydes upon reductive work-up.  Reduction of an ester with diisobutylaluminium hydride (DIBAL-H) or sodium aluminium hydride yields an aldehyde.  Reduction of an acid chloride via the Rosenmund reduction or using lithium tri-t-butoxyaluminium hydride (LiAlH(O-t-C4H9)3).

86  Reaction of ketones with methoxymethylenetriphenylphosphine in a modified Wittig reaction.  Various formylation reactions, such as the Vilsmeier-Haack reaction, can be used to introduce an aldehyde group.  In the Nef reaction, aldehydes form by hydrolysis of salts of primary nitro compounds.  Zincke aldehydes form by reaction of pyridinium salts with secondary amines followed by hydrolysis.  In the Stephen aldehyde synthesis aldehydes form from nitriles, tin(II) chloride, and hydrochloric acid.  In the Meyers synthesis they form by hydrolysis of an oxazine.  The McFadyen-Stevens reaction is a base-catalyzed thermal decomposition of acylsulfonylhydrazides

] Examples of aldehydes

 Methanal (Formaldehyde)  Ethanal (Acetaldehyde)  Propanal (Propionaldehyde)  Butanal (butyraldehyde)  Benzaldehyde  Cinnamaldehyde

Week No. 22 /Carboxylic acid

Carboxylic acids are organic acids characterized by the presence of a carboxyl group, which has the formula -C(=O)OH, usually written -COOH or -CO2H. Carboxylic acids are Brønsted-Lowry acids — they are proton donors. Salts and anions of carboxylic acids are called carboxylates.

87 Physical properties

Solubility

Carboxylic acid dimers

Carboxylic acids are polar. Because they are both hydrogen-bond acceptors (the carbonyl) and hydrogen-bond donors (the hydroxyl), they also participate in hydrogen bonding. Together the hydroxyl and carbonyl group forms the functional group carboxyl. Carboxylic acids usually exist as dimeric pairs in nonpolar media due to their tendency to “self-associate.” Smaller carboxylic acids (1 to 5 carbons) are soluble with water, whereas higher carboxylic acids are less soluble due to the increasing hydrophobic nature of the alkyl chain. These longer chain acids tend to be rather soluble in less-polar solvents such as ethers and alcohols.[2]

Boiling points

Carboxylic acids tend to have higher boiling points than water, not only because of their increased surface area, but because of their tendency to form stabilised dimers. Carboxylic acids tend to evaporate or boil as these dimers. For boiling to occur, either the dimer bonds must be broken, or the entire dimer arrangement must be vaporised, both of which increase enthalpy of vaporisation requirements significantly.

Acidity

Carboxylic acids are typically weak acids, meaning that they only partially dissociate into H+ cations and RCOO- anions in neutral aqueous solution. For example, at room temperature, only 0.02 % of all acetic acid molecules are dissociated. Electronegative substituents give stronger acids.

Ionization of a carboxylic acid gives a carboxylate anion, which is stabilized because the negative charge is shared (delocalized) between the two oxygen atoms. Each of the carbon-oxygen bonds in a carboxylate anion has partial double-bond character.

88 Odor

Carboxylic acids often have strong odors, especially the volatile derivatives. Most common are acetic acid (vinegar) and butyric acid (rancid butter). On the other hand, esters of carboxylic acids tend to have pleasant odors and many are used in perfumes.

Laboratory methods

Preparative methods for small scale reactions for research, instruction, or for production of small amounts of fine chemicals often employ expensive consumable reagents.

 oxidation of primary alcohols or aldehydes with strong oxidants such as potassium dichromate, Jones reagent, potassium permanganate, or sodium chlorite. The method is amenable to laboratory conditions compared to the industrial use of air, which is “greener” since it yields less inorganic side products such as chromium or manganese oxides.  Oxidative cleavage of olefins by ozonolysis, potassium permanganate, or potassium dichromate.  Carboxylic acids can also be obtained by the hydrolysis of nitriles, esters, or amides, generally with acid- or base-catalysis.  Carbonation of a Grignard and organolithium reagents:

RLi + CO2 RCO2Li

RCO2Li + HCl RCO2H + LiCl

 Halogenation followed by hydrolysis of methyl ketones in the haloform reaction  The Kolbe-Schmitt reaction provides a route to salicylic acid, precursor to aspirin.

The most widely practiced reactions convert carboxylic acids into esters, amides, carboxylate salts, acid chlorides, and alcohols. Carboxylic acids react with bases to form carboxylate salts, in which the hydrogen of the hydroxyl (- OH) group is replaced with a metal cation. Thus, acetic acid found in vinegar reacts with sodium bicarbonate (baking soda) to form sodium acetate, carbon dioxide, and water:

− + CH3COOH + NaHCO3 → CH3COO Na + CO2 + H2O

Carboxylic acids also react with alcohols to give esters. This process is heavily used in the production of polyesters. Similarly carboxylic acids are converted

89 into amides, but this conversion typically does not occur by direct reaction of the carboxylic acid and the amine. Instead esters are typical precursors to amides. The conversion of amino acids into peptides is a major biochemical process that requires ATP.

The hydroxyl group on carboxylic acids may be replaced with a chlorine atom using thionyl chloride to give acyl chlorides. In nature, carboxylic acids are converted to thioesters.

The carboxylic acid can be reduced to the alcohol by hydrogenation or using stoichiometric hydride reducing agents such as lithium aluminum hydride.

Week No. 23 / ESTER, AMINE & AMIDE ESTER

Esters are chemical compounds derived by reacting an oxoacid (one containing an oxo group, X=O) with a hydroxyl compound such as an alcohol or phenol.[1] Esters are usually derived from an inorganic acid or organic acid in which at least one -OH (hydroxyl) group is replaced by an -O-alkyl (alkoxy) group, and most commonly from carboxylic acids and alcohols.

Esters are ubiquitous. Many naturally occurring fats and oils are the fatty acid esters of glycerol. Esters with low molecular weight are commonly used as fragrances and found in essential oils and pheromones. Phosphoesters form the backbone of DNA molecules. Nitrate esters, such as nitroglycerin, are known for their explosive properties, while polyesters are important plastics, with monomers linked by ester moieties.

Nomenclature Main article: IUPAC nomenclature of organic chemistry#Esters

The names for esters are derived from the names of the parent alcohol and the carboxylic acid. For esters derived from the simplest carboxylic acids, the traditional names are used, such as formate, acetate, propionate, and butyrate.

90 For esters from more complex carboxylic acids, the systematic name for the acid is used, followed by the suffix -oate. For example, hexyl octanoate, also called hexyl caprylate, has the formula CH3(CH2)6CO2(CH2)5CH3.

The chemical formulas of esters are typically written in the format of RCO2R', where R and R' are the organic parts of the carboxylic acid and alcohol respectively. For example butyl acetate, derived from butanol and acetic acid would be written CH3CO2C4H9. Alternative presentations are common including BuOAc and CH3COOC4H9.

Cyclic esters are called lactones. One example of many is valerolactone

Appendix of ester odorants

Many esters have distinctive fruit-like odors, which has led to their commonplace use in artificial flavorings and fragrances.

Ester Name Structure Odor or occurrence

Allyl hexanoate pineapple

Benzyl acetate pear, strawberry, jasmine

Bornyl acetate pine tree flavor

91 Butyl butyrate pineapple

Ethyl acetate nail polish remover, model paint, model airplane glue

Ethyl butyrate banana, pineapple, strawberry

Ethyl hexanoate pineapple,waxy-green banana

Ethyl cinnamate cinnamon

Ethyl formate lemon, rum, strawberry

Ethyl heptanoate apricot, cherry, grape, raspberry

Ethyl isovalerate apple

Ethyl lactate butter, cream

Ethyl nonanoate grape

Ethyl pentanoate apple

Geranyl acetate geranium

Geranyl butyrate cherry

Geranyl pentanoate apple

92 Isobutyl acetate cherry, raspberry, strawberry

Isobutyl formate raspberry

Isoamyl acetate pear, banana (flavoring in Pear drops)

Isopropyl acetate fruity

Linalyl acetate lavender, sage

Linalyl butyrate peach

Linalyl formate apple, peach

Methyl acetate glue

Methyl anthranilate grape, jasmine

Methyl benzoate fruity, ylang ylang, feijoa

Methyl butyrate (methyl pineapple, apple, strawberry butanoate)

Methyl cinnamate strawberry

93 Methyl pentanoate (methyl flowery valerate)

Methyl phenylacetate honey

Methyl salicylate (oil of Modern root beer, wintergreen, Germolene and wintergreen) Ralgex ointments (UK)

Nonyl caprylate orange

Octyl acetate fruity-orange

Octyl butyrate parsnip

Amyl acetate (pentyl apple, banana acetate)

Pentyl butyrate (amyl apricot, pear, pineapple butyrate)

Pentyl hexanoate (amyl apple, pineapple caproate)

Pentyl pentanoate (amyl apple valerate)

Propyl ethanoate pear

Propyl isobutyrate rum

Terpenyl butyrate cherry

94 Amide

In chemistry, an amide is usually an organic compound that contains the functional group consisting of an acyl group (R-C=O) linked to a nitrogen atom (N). The term refers both to a class of compounds and a functional group within those compounds. The term amide also refers to deprotonated form of ammonia - (NH3) or an amine, often represented as anions R2N . The remainder of this article is about the carbonyl-nitrogen sense of amide. For discussion of these "anionic amides," see the articles sodium amide and lithium diisopropylamide.

Amides are commonly formed via reactions of a carboxylic acid with an amine. Many methods are known for driving the unfavorable equilibrium to the right:

RCO2H + R'R"NH RC(O)NR'R" + H2O

For the most part, these reactions involve "activating" the carboxylic acid and the best known method, the Schotten-Baumann reaction, which involves conversion of the acid to the acid chlorides:

Amine Primary amine Secondary amine Tertiary amine

Amines are organic compounds and functional groups that contain a basic nitrogen atom with a lone pair. Amines are derivatives of ammonia, wherein

95 one or more hydrogen atoms have been replaced by a substituent such as an alkyl or aryl group. Important amines include amino acids, biogenic amines, trimethylamine, and aniline; see Category:Amines for a list of amines. Inorganic derivatives of ammonia are also called amines, such as chloramine (NClH2).

Compounds with the nitrogen atom attached to a carbonyl of the structure R-

C(=O)NR2 are called amides and have different chemical properties from amines

Application of amines

Dyes

Primary aromatic amines are used as a starting material for the manufacture of azo dyes. It reacts with nitric(III) acid to form diazonium salt, which can undergo coupling reaction to form azo compound. As azo-compounds are highly coloured, they are widely used in dyeing industries, such as:

 Methyl orange  Direct brown 138  Sunset yellow FCF  Ponceau

Drugs

Many drugs are designed to mimic or to interfere with the action of natural amine neurotransmitters, exemplified by the amine drugs:

 Chlorpheniramine is an antihistamine that helps to relieve allergic disorders due to cold, hay fever, itchy skin, insect bites and stings.  Chlorpromazine is a tranquillizer that sedates without inducing sleep. It is used to relieve anxiety, excitement, restlessness or even mental disorder.  Ephedrine and Phenylephrine, as amine hydrochlorides, are used as decongestants.  Amphetamine, Methamphetamine, and Methcathinone are amines that are listed as controlled substances by the DEA.  Amitriptyline, Imipramine, Lofepramine and Clomipramine are tricyclic antidepressants and tertiary amines  Nortriptyline, Desipramine, and Amoxapine are tricyclic antidepressants and secondary amines  (The tricyclics are grouped by the nature of the final amine group on the side chain.)

96 Safety

Low molecular weight amines are toxic and some are easily absorbed through the skin. Many higher molecular weight amines are highly active biologically.

Week No. 24/ Acids and Bases

Acids and Bases, two classes of chemical compounds that display generally opposite characteristics. Acids taste sour, turn litmus (a pink dye derived from lichens) red, and often react with some metals to produce hydrogen gas. Bases taste bitter, turn litmus blue, and feel slippery. When aqueous (water) solutions of an acid and a base are combined, a neutralization reaction occurs. This reaction is characteristically very rapid and generally produces water and a salt. For example, sulfuric acid and sodium hydroxide, NaOH, yield water and sodium sulfate:

H2SO4 + 2NaOH⇄2H2O + Na2SO4

Names of acids :The names of acids without oxygen in the molecule have the prefix hydro- (sometimes shortened to hydr-) and the suffix -ic attached to the stem based on the names of the constituent elements (other than hydrogen). For example, HCl (made of hydrogen and chlorine) is hydrochloric acid; HBr (made of hydrogen and bromine) is hydrobromic acid; HI (made of hydrogen and iodine) is hydroiodic acid; HCN (made of hydrogen, carbon, and nitrogen) is hydrocyanic acid; and H2S (made of hydrogen and sulfur) is hydrosulfuric acid.

Names of acids containing oxygen (known as oxoacids) are derived from the number of oxygen atoms in the molecules of a series, or class, of acids. An example of an oxoacid series is as follows: HClO, HClO2, HClO3, HClO4. If a class of acids contains only one member, its name is given the suffix –ic. For example, H2CO3 is carbonic acid. If an acid series contains two acids, such as H2SO4 and H2SO3, the acid containing more oxygen atoms is given the suffix – ic, while the acid with fewer oxygen atoms is given the suffix –ous. For example, H2SO4 is sulfuric acid, and H2SO3 is sulfurous acid. Similarly, HNO3 is nitric acid, and HNO2 is nitrous acid. In the case of an extensive acid series (such as HClO, HClO2, HClO3, HClO4), the acid with the fewest oxygen atoms is given the prefix hypo- and the suffix –ous, and the acid with the most oxygen atoms is given the prefix per-. In the above example, HClO is perchloric acid. In the classical naming system, acids are named according to their anions. That ionic suffix is dropped and replaced with a new suffix (and sometimes prefix), according to the table below. For example, HCl has chloride as its anion, so the -ide suffix makes it take the form hydrochloric acid. In the IUPAC naming

97 system, "aqueous" is simply added to the name of the ionic compound. Thus, for hydrogen chloride, the IUPAC name would be aqueous hydrogen chloride. The prefix "hydro-" is added only if the acid is made up of just hydrogen and one other element.

Classical naming system:

Anion prefix Anion suffix Acid prefix Acid suffix Example per ate per ic acid perchloric acid (HClO4)

ate ic acid chloric acid (HClO3)

ite ous acid chlorous acid (HClO2) hypo ite hypo ous acid hypochlorous acid (HClO) ide hydro ic acid hydrochloric acid (HCl)

Classification of acids

Mineral acids

 Hydrogen halides and their solutions: hydrochloric acid (HCl), hydrobromic acid (HBr), hydroiodic acid (HI)  Halogen oxoacids: hypochloric acid, chloric acid, perchloric acid, periodic acid and corresponding compounds for bromine and iodine  Sulfuric acid (H2SO4)  Fluorosulfuric acid  Nitric acid (HNO3)  Phosphoric acid (H3PO4)  Fluoroantimonic acid  Fluoroboric acid  Hexafluorophosphoric acid  Chromic acid (H2CrO4)  Boric acid (H3BO3)

Sulfonic acids

 Methanesulfonic acid (or mesylic acid, CH3SO3H)  Ethanesulfonic acid (or esylic acid, CH3CH2SO3H)  Benzenesulfonic acid (or besylic acid, C6H5SO3H)  p-Toluenesulfonic acid (or tosylic acid, CH3C6H4SO3H)  Trifluoromethanesulfonic acid (or triflic acid, CF3SO3H)  Polystyrene sulfonic acid (sulfonated polystyrene, [CH2CH(C6H4)SO3H]n)

Carboxylic acids

 Acetic acid

98  Citric acid  Formic acid  Gluconic acid  Lactic acid  Oxalic acid  Tartaric acid

Vinylogous carboxylic acids

 Ascorbic acid  Meldrum's acid

Nucleic acids

 Deoxyribonucleic acid (DNA)  Ribonucleic acid (RNA)

A base:

in chemistry is generally an aqueous substance that can accept hydronium ions. Bases are also the oxides or hydroxides of metals. A soluble base is also often referred to as an alkali if hydroxide ions (OH−) are involved. Such a base in solution has a surplus of OH- ions. This refers to the Brønsted-Lowry theory of acids and bases. Alternative definitions of bases include electron pair donors (Lewis)[1], and as sources of hydroxide anions (Arrhenius).[2] In addition to this, bases can commonly be thought of as any chemical compound that, when dissolved in water, gives a solution with a hydrogen ion activity lower than that of pure water, i.e. a pH higher than 7.0 at standard conditions. Examples of common bases are sodium hydroxide and ammonia.

 Bases can be thought of as the chemical opposite of acids. A reaction between an acid and base is called neutralization. Bases and acids are seen as opposites because the effect of an acid is to increase the + hydronium ion (H3O ) concentration in water, whereas bases reduce this concentration. Bases and acids are typically found in aqueous solution forms. Aqueous solutions of bases react with aqueous solutions of acids to produce water and salts in aqueous solutions in which the salts separate into their component ions. If the aqueous solution is a saturated solution with respect to a given salt solute any additional such salt present in the solution will result in formation of a precipitate of the salt.

99 Properties

Some general properties of bases include:

 Slimy or soapy feel on fingers, due to saponification of the lipids in human skin  Concentrated or strong bases are caustic on organic matter and react violently with acidic substances  Aqueous solutions or molten bases dissociate in ions and conduct electricity.  Reactions with indicators: bases turn red litmus paper blue, phenolphthalein pink, keep bromthymol blue in its natural colour of blue, and turn methyl orange yellow.  the pH is above 7

Neutralization of acids

When dissolved in water, the strong base sodium hydroxide decomposes into hydroxide and sodium ions:

NaOH → Na+ + OH− and similarly, in water hydrogen chloride forms hydronium and chloride ions:

+ − HCl + H2O → H3O + Cl

+ − When the two solutions are mixed, the H3O and OH ions combine to form water molecules:

+ − H3O + OH → 2 H2O

If equal quantities of NaOH and HCl are dissolved, the base and the acid exactly neutralize, leaving only NaCl, effectively table salt, in solution.

Weak bases, such as soda or egg white, should be used to neutralize any acid spills. Neutralizing acid spills with strong bases, such as sodium hydroxide or potassium hydroxide can cause a violent exothermic reaction, and the base itself can cause just as much damage as the original acid spill.

Alkalinity of non-hydroxides

Bases are generally compounds that can neutralize an amount of acids. Both sodium carbonate and ammonia are bases, although neither of these substances contains OH− groups. Both compounds accept H+ when dissolved in water:

100 + - - Na2CO3 + H2O → 2 Na + HCO3 + OH

+ - NH3 + H2O → NH4 + OH

From this, a pH, or acidity, can be calculated for aqueous solutions of bases. Bases also directly act as electron-pair donors themselves:

2- + - CO3 + H → HCO3

+ + NH3 + H → NH4

Carbon can act as a base as well as nitrogen and oxygen. This occurs typically in compounds such as butyl lithium, alkoxides, and metal amides such as sodium amide. Bases of carbon, nitrogen and oxygen without resonance stabilization are usually very strong, or superbases, which cannot exist in a water solution due to the acidity of water. Resonance stabilization, however, enables weaker bases such as carboxylates; for example, sodium acetate is a weak base.

Strong bases

A strong base is a basic chemical compound that is able to deprotonate very weak acids in an acid-base reaction. Common examples of strong bases are the hydroxides of alkali metals and alkaline earth metals like NaOH and Ca(OH)2. Very strong bases are even able to deprotonate very weakly acidic C–H groups in the absence of water. Here is a list of several strong bases:

 Potassium hydroxide (KOH)  Barium hydroxide (Ba(OH)2)  Caesium hydroxide (CsOH)  Sodium hydroxide (NaOH)  Strontium hydroxide (Sr(OH)2)  Calcium hydroxide (Ca(OH)2)  Lithium hydroxide (LiOH)  Rubidium hydroxide (RbOH)

The cations of these strong bases appear in the first and second groups of the periodic table (alkali and earth alkali metals).

Acids with a pKa of more than about 13 are considered very weak, and their conjugate bases are strong bases.

Salts are named for the ions that compose them. The cation (positively charged ion) within the compound is named first. Examples are NaCl (sodium chloride),

101 BaO (barium oxide), Fe(NO3)2 [iron (II) nitrate], and Fe(NO3)3 [iron (III) nitrate].

Week No. 25, 26, 27, 28 /Polymer

Appearance of real linear polymer chains as recorded using an atomic force microscope on surface under liquid medium. Chain contour length for this polymer is ~204 nm; thickness is ~0.4 nm.

A polymer is a large molecule (macromolecule) composed of repeating structural units typically connected by covalent chemical bonds. While polymer in popular usage suggests plastic, the term actually refers to a large class of natural and synthetic materials with a variety of properties.

Due to the extraordinary range of properties accessible in polymeric materials they have come to play an essential and ubiquitous role in everyday life—from plastics and elastomers on the one hand to natural biopolymers such as DNA and proteins that are essential for life on the other. A simple example is polyethylene, whose repeating unit is based on ethylene (IUPAC name ethene) monomer. Most commonly, as in this example, the continuously linked backbone of a polymer used for the preparation of plastics consists mainly of carbon atoms. However, other structures do exist; for example, elements such as silicon form familiar materials such as silicones, examples being silly putty and waterproof plumbing sealant. The backbone of DNA is in fact based on a phosphodiester bond, and repeating units of polysaccharides (e.g. cellulose) are joined together by glycosidic bonds via oxygen atoms.

Natural polymeric materials such as shellac, amber, and natural rubber have been in use for centuries. Biopolymers such as proteins and nucleic acids play crucial roles in biological processes. A variety of other natural polymers exist, such as cellulose, which is the main constituent of wood and paper.

The list of synthetic polymers includes synthetic rubber, Bakelite, neoprene, nylon, PVC, polystyrene, polyethylene, polypropylene, polyacrylonitrile, PVB, silicone, and many more.

102 Polymers are studied in the fields of polymer chemistry, polymer physics, and polymer science.

Polymer synthesis Main article: Polymerization

The repeating unit of the polymer polypropylene

Polymerization is the process of combining many small molecules known as monomers into a covalently bonded chain. During the polymerization process, some chemical groups may be lost from each monomer. This is the case, for example, in the polymerization of PET polyester. The monomers are terephthalic acid (HOOC-C6H4-COOH) and ethylene glycol (HO-CH2-CH2-OH) but the repeating unit is -OC-C6H4-COO-CH2-CH2-O-, which corresponds to the combination of the two monomers with the loss of two water molecules. The distinct piece of each monomer that is incorporated into the polymer is known as a repeat unit or monomer residue.

Laboratory synthesis

Laboratory synthetic methods are generally divided into two categories, step- growth polymerization and chain-growth polymerization[4]. The essential difference between the two is that in chain growth polymerization, monomers are added to the chain one at a time only, whereas in step-growth polymerization chains of monomers may combine with one another directly. However, some newer methods such as plasma polymerization do not fit neatly into either category. Synthetic polymerization reactions may be carried out with or without a catalyst. Efforts towards rational synthesis of biopolymers via laboratory synthetic methods, especially artificial synthesis of proteins, is an area of intense research.

Biological synthesis Main article: Biopolymer

There are three main classes of biopolymers: polysaccharides, polypeptides, and polynucleotides. In living cells, they may be synthesized by enzyme-mediated processes, such as the formation of DNA catalyzed by DNA polymerase. The synthesis of proteins involves multiple enzyme-mediated processes to transcribe genetic information from the DNA to RNA and subsequently translate that

103 information to synthesize the specified protein from amino acids. The protein may be modified further following translation in order to provide appropriate structure and functioning.

Week No. 29, 30 /Chromatography

Chromatography (from Greek χρώμα:chroma, color and γραφειν:graphein to write) is the collective term for a set of laboratory techniques for the separation of mixtures. It involves passing a mixture dissolved in a "mobile phase" through a stationary phase, which separates the analyte to be measured from other molecules in the mixture based on differential partitioning between the mobile and stationary phases. Subtle differences in compounds partition coefficient results in differential retention on the stationary phase and thus changing the separation.

Chromatography may be preparative or analytical. The purpose of preparative chromatography is to separate the components of a mixture for further use (and is thus a form of purification). Analytical chromatography is done normally with smaller amounts of material and is for measuring the relative proportions of analytes in a mixture. The two are not mutually exclusive.

Chromatography terms

 The analyte is the substance that is to be separated during chromatography.  Analytical chromatography is used to determine the existence and possibly also the concentration of analyte(s) in a sample.  A bonded phase is a stationary phase that is covalently bonded to the support particles or to the inside wall of the column tubing.  A chromatogram is the visual output of the chromatograph. In the case of an optimal separation, different peaks or patterns on the chromatogram correspond to different components of the separated mixture.

104 Plotted on the x-axis is the retention time and plotted on the y-axis a signal (for example obtained by a spectrophotometer, mass spectrometer or a variety of other detectors) corresponding to the response created by the analytes exiting the system. In the case of an optimal system the signal is proportional to the concentration of the specific analyte separated.

 A chromatograph is equipment that enables a sophisticated separation e.g. gas chromatographic or liquid chromatographic separation.  Chromatography is a physical method of separation in which the components to be separated are distributed between two phases, one of which is stationary (stationary phase) while the other (the mobile phase) moves in a definite direction.  The effluent is the mobile phase leaving the column.  An immobilized phase is a stationary phase which is immobilized on the support particles, or on the inner wall of the column tubing.  The mobile phase is the phase which moves in a definite direction. It may be a liquid (LC and CEC), a gas (GC), or a supercritical fluid (supercritical-fluid chromatography, SFC). A better definition: The mobile phase consists of the sample being separated/analyzed and the solvent that moves the sample through the column. In the case of HPLC the mobile phase consists of a non-polar solvent(s) such as hexane in normal phase or polar solvents in reverse phase chromotagraphy and the sample being separated. The mobile phase moves through the

105 chromatography column (the stationary phase) where the sample interacts with the stationary phase and is separated.  Preparative chromatography is used to purify sufficient quantities of a substance for further use, rather than analysis.  The retention time is the characteristic time it takes for a particular analyte to pass through the system (from the column inlet to the detector) under set conditions. See also: Kovat's retention index  The sample is the matter analyzed in chromatography. It may consist of a single component or it may be a mixture of components. When the sample is treated in the course of an analysis, the phase or the phases containing the analytes of interest is/are referred to as the sample whereas everything out of interest separated from the sample before or in the course of the analysis is referred to as waste.  The solute refers to the sample components in partition chromatography.  The solvent refers to any substance capable of solubilizing other substance, and especially the liquid mobile phase in LC.  The stationary phase is the substance which is fixed in place for the chromatography procedure. Examples include the silica layer in Chromatography#Thin layer chromatography Submitted by James Nguyen

Techniques by chromatographic bed shape

Column chromatography For more details on this topic, see Column chromatography.

Column chromatography is a separation technique in which the stationary bed is within a tube. The particles of the solid stationary phase or the support coated with a liquid stationary phase may fill the whole inside volume of the tube (packed column) or be concentrated on or along the inside tube wall leaving an open, unrestricted path for the mobile phase in the middle part of the tube (open tubular column). Differences in rates of movement through the medium are calculated to different retention times of the sample.

In 1978, W. C. Still introduced a modified version of column chromatography called flash column chromatography (flash). The technique is very similar to the traditional column chromatography, except for that the solvent is driven through the column by applying positive pressure. This allowed most separations to be performed in less than 20 minutes, with improved separations compared to the old method. Modern flash chromatography systems are sold as pre-packed plastic cartridges, and the solvent is pumped through the cartridge. Systems may also be linked with detectors and fraction collectors providing

106 automation. The introduction of gradient pumps resulted in quicker separations and less solvent usage.

In expanded bed adsorption, a fluidized bed is used, rather than a solid phase made by a packed bed. This allows omission of initial clearing steps such as centrifugation and filtration, for culture broths or slurries of broken cells.

Planar chromatography

Thin layer chromatography is used to separate components of chlorophyll

Planar chromatography is a separation technique in which the stationary phase is present as or on a plane. The plane can be a paper, serving as such or impregnated by a substance as the stationary bed (paper chromatography) or a layer of solid particles spread on a support such as a glass plate (thin layer chromatography). Different compounds in the sample mixture travel different distances according to how strongly they interact with the stationary phase as compared to the mobile phase. The specific Retardation factor (Rf) of each chemical can be used to aid in the identification of an unknown substance.

Paper chromatography For more details on this topic, see Paper chromatography.

Paper chromatography is a technique that involves placing a small dot or line of sample solution onto a strip of chromatography paper. The paper is placed in a jar containing a shallow layer of solvent and sealed. As the solvent rises through the paper, it meets the sample mixture which starts to travel up the paper with the solvent. This paper is made of cellulose, a polar substance, and the

107 compounds within the mixture travel farther if they are non-polar. More polar substances bond with the cellulose paper more quickly, and therefore do not travel as far.

Thin layer chromatography For more details on this topic, see Thin layer chromatography.

Thin layer chromatography (TLC) is a widely employed laboratory technique and is similar to paper chromatography. However, instead of using a stationary phase of paper, it involves a stationary phase of a thin layer of adsorbent like silica gel, alumina, or cellulose on a flat, inert substrate. Compared to paper, it has the advantage of faster runs, better separations, and the choice between different adsorbents. For even better resolution and to allow for quantification, high-performance TLC can be used.

Displacement Chromatography

The basic principle of displacement chromatography is: A molecule with a high affinity for the chromatography matrix (the displacer) will compete effectively for binding sites, and thus displace all molecules with lesser affinities. There are distinct differences between displacement and elution chromatography. In elution mode, substances typically emerge from a column in narrow, Gaussian peaks. Wide separation of peaks, preferably to baseline, is desired in order to achieve maximum purification. The speed at which any component of a mixture travels down the column in elution mode depends on many factors. But for two substances to travel at different speeds, and thereby be resolved, there must be substantial differences in some interaction between the biomolecules and the chromatography matrix. Operating parameters are adjusted to maximize the effect of this difference. In many cases, baseline separation of the peaks can be achieved only with gradient elution and low column loadings. Thus, two drawbacks to elution mode chromatography, especially at the preparative scale, are operational complexity, due to gradient solvent pumping, and low throughput, due to low column loadings. Displacement chromatography has advantages over elution chromatography in that components are resolved into consecutive zones of pure substances rather than “peaks”. Because the process takes advantage of the nonlinearity of the isotherms, a larger column feed can be separated on a given column with the purified components recovered at significantly higher concentrations.

Techniques by physical state of mobile phase

Gas chromatography For more details on this topic, see Gas chromatography.

108 Gas chromatography (GC), also sometimes known as Gas-Liquid chromatography, (GLC), is a separation technique in which the mobile phase is a gas. Gas chromatography is always carried out in a column, which is typically "packed" or "capillary" (see below) .

Gas chromatography (GC) is based on a partition equilibrium of analyte between a solid stationary phase (often a liquid silicone-based material) and a mobile gas (most often Helium). The stationary phase is adhered to the inside of a small-diameter glass tube (a capillary column) or a solid matrix inside a larger metal tube (a packed column). It is widely used in analytical chemistry; though the high temperatures used in GC make it unsuitable for high molecular weight biopolymers or proteins (heat will denature them), frequently encountered in biochemistry, it is well suited for use in the petrochemical, environmental monitoring, and industrial chemical fields. It is also used extensively in chemistry research.

Liquid chromatography

Liquid chromatography (LC) is a separation technique in which the mobile phase is a liquid. Liquid chromatography can be carried out either in a column or a plane. Present day liquid chromatography that generally utilizes very small packing particles and a relatively high pressure is referred to as high performance liquid chromatography (HPLC).

In the HPLC technique, the sample is forced through a column that is packed with irregularly or spherically shaped particles or a porous monolithic layer (stationary phase) by a liquid (mobile phase) at high pressure. HPLC is historically divided into two different sub-classes based on the polarity of the mobile and stationary phases. Technique in which the stationary phase is more polar than the mobile phase (e.g. toluene as the mobile phase, silica as the stationary phase) is called normal phase liquid chromatography (NPLC) and the opposite (e.g. water-methanol mixture as the mobile phase and C18 = octadecylsilyl as the stationary phase) is called reversed phase liquid chromatography (RPLC). Ironically the "normal phase" has fewer applications and RPLC is therefore used considerably more.

Specific techniques which come under this broad heading are listed below. It should also be noted that the following techniques can also be considered fast protein liquid chromatography if no pressure is used to drive the mobile phase through the stationary phase. See also Aqueous Normal Phase Chromatography.

Affinity chromatography For more details on this topic, see Affinity chromatography.

109 Affinity chromatography is based on selective non-covalent interaction between an analyte and specific molecules. It is very specific, but not very robust. It is often used in biochemistry in the purification of proteins bound to tags. These fusion proteins are labeled with compounds such as His-tags, biotin or antigens, which bind to the stationary phase specifically. After purification, some of these tags are usually removed and the pure protein is obtained.

Supercritical fluid chromatography For more details on this topic, see Supercritical fluid chromatography.

Supercritical fluid chromatography is a separation technique in which the mobile phase is a fluid above and relatively close to its critical temperature and pressure.

Techniques by separation mechanism

Ion exchange chromatography For more details on this topic, see Ion exchange chromatography.

Ion exchange chromatography uses ion exchange mechanism to separate analytes. It is usually performed in columns but can also be useful in planar mode. Ion exchange chromatography uses a charged stationary phase to separate charged compounds including amino acids, peptides, and proteins. In conventional methods the stationary phase is an ion exchange resin that carries charged functional groups which interact with oppositely charged groups of the compound to be retained. Ion exchange chromatography is commonly used to purify proteins using FPLC.

Size exclusion chromatography For more details on this topic, see Size exclusion chromatography.

Size exclusion chromatography (SEC) is also known as gel permeation chromatography (GPC) or gel filtration chromatography and separates molecules according to their size (or more accurately according to their hydrodynamic diameter or hydrodynamic volume). Smaller molecules are able to enter the pores of the media and, therefore, take longer to elute, whereas larger molecules are excluded from the pores and elute faster. It is generally a low-resolution chromatography technique and thus it is often reserved for the final, "polishing" step of a purification. It is also useful for determining the tertiary structure and quaternary structure of purified proteins, especially since it can be carried out under native solution conditions.

110 Special techniques

Reversed-phase chromatography For more details on this topic, see Reversed-phase chromatography.

Reversed-phase chromatography is an elution procedure used in liquid chromatography in which the mobile phase is significantly more polar than the stationary phase.

Two-dimensional chromatography

In some cases, the chemistry within a given column can be insufficient to separate some analytes. It is possible to direct a series of unresolved peaks onto a second column with different physico-chemical (Chemical classification) properties. Since the mechanism of retention on this new solid support is different from the first dimensional separation, it can be possible to separate compounds that are indistinguishable by one-dimensional chromatography.

111