PSMA General Chemistry PDF

PHARMACEUTICAL SEMINAR 1 PSMA411 | INORGANIC CHEMISTRY GENERAL CHEMISTRY THERMODYNAMICS Note: Heat does not flow from cold to hot, unless there is...

PHARMACEUTICAL SEMINAR 1 PSMA411 | INORGANIC CHEMISTRY GENERAL CHEMISTRY THERMODYNAMICS Note: Heat does not flow from cold to hot, unless there is pumping of energy into the system. CHEMISTRY – a branch of science that is concerned with the Ex: ball on top of hill will naturally fall down; ball in low position study of matter and the changes it undergoes. will not climb on top of a hill. CHEMISTRY BRANCHES 3. Solid crystalline substance has zero entropy 1. Organic Chemistry – the study of organic compounds Entropy of perfect crystal at absolute zero is containing carbon. exactly equal to zero 2. Inorganic Chemistry – study of elements and All processes cease as temperature approaches compounds considered inorganic. absolute zero 3. Biochemistry – study of chemistry of life 4. Analytical Chemistry – is an area of chemistry that is 4. Zeroth Law used in characterization of matter, both quantitatively If 2 bodies are in equilibrium with the third body and qualitatively. separately, then it follows that the 1st and 2nd body 5. Physical Chemistry – study of macroscopic and are also in thermal equilibrium. particular phenomena in chemical systems in terms of principles, practices and concepts of physics such as motion energy, force, time, thermodynamics, quantum CLASSIFICATION OF MATTER chemistry & chemical equilibrium. MATTER – anything that occupies space & has mass. THERMODYNAMICS Has a composition Study of energy, the conversion of energy between Has structure various forms and ability of energy to do work. Involves change From the word therme (heat) and dynamis (power or Requires energy for changes & interactions energy) Has classifications (SOLID, LIQUID GAS) Can be a pure substance. THERMODYNAMICS LAWS/THEORIES 1. Law of Conservation of Energy PURE SUBSTANCE Total amount of energy in an isolated system kind of matter possessing a definite & unvarying remains constant over time. composition Total energy is conserved over time. can be an element Energy is neither created nor destroyed. if physically combined, they form a mixture If an energy/heat enters a system, it is called internal energy. ELEMENT simplest form of substance that cannot be 2 WAYS OF SYSTEM CAN INCREASE INTERNAL ENERGY decomposed by simple chemical reaction 1. Transfer of energy into the system has 1 material or atom (either positive or negative) 2. The surroundings can perform work into the system Ex: If surroundings perform 300 joules into the system, then COMPOUND the system’s internal energy goes up by 300 joules. ▪ When 2 or more elements chemically combine Change is considered to be positive. Surroundings ▪ Substances composed of 2 or more elements united then decreased in 300 joules of energy. Therefore, chemically in definite proportion energy is not created nor destroyed, just transferred. MIXTURE ENTROPY – thermodynamic property that is a measure of a Made up of 2 or more substances each which system’s thermal energy per unit temperature. retains its own characteristics or properties Not chemically combined 2. Law Entropy Can be homogeneous (mixture is uniform) or The spontaneous natural processes increases heterogeneous (there is insoluble substance & entropy overall. composition is not uniform throughout the mixture; Can lead to higher disorderliness. 2 or more phases) For natural process, the change of entropy is considered positive (higher than 0) ATOMS – consists of a positively charged core (atomic nucleus) Heat can spontaneously be conducted only from a which contains protons & neutrons, & which maintains a higher temperature region to a lower temperature number of electrons to balance the positive charge in the region, but not the other way around. nucleus. Ex: metal rod placed between hot & cold object = heat Basic unit that makes up all matter will flow from hot to that cold object Basic unit of element that can enter into a chemical reaction Protons – positively charged; shown in Ernest Rutherford’s experiment (Gold Foil Experiment) Electrons – negatively charged; shown in Cathode ray tube experiment Karl Ferdinand Braun 1 | PAGE  Neutrons – uncharged particles; discovered by James Considered as ionic compounds Chadwick Metal is always written first followed by non- metal/polyatomic ion MOLECULES – smallest indivisible portion of a pure chemical substance that has its unique set of chemical properties, METALLIC COMPOUNDS EXAMPLES: that is, its potential to undergo a certain set of chemical 1. Calcium carbonate reactions with other substances. 2. Aluminum sulfide Aggregates of at least 2 atoms in a definite arrangement held together by chemical forces. TYPES OF INORGANIC COMPOUNDS Formed when 2 or more atoms of an element chemically join together 1. ACIDS May be homonuclear (consists of atoms of 1 chemical Contain hydrogen which is replaceable by a element. Ex: oxygen) metal Maybe heteronuclear (composed of more than 1 element. Yields hydrogen ions in water solutions Ex: H2O) Donates protons Accepts electrons ION – charged species, an atom or a molecule, that has lost or Turns blue litmus paper into red gained one or more electrons. Most acids contain a hydrogen atom that is Atom or molecule in which the total number of electrons is bonded that can release to yield a cation and not equal to total number of protons. anion in water. Net positive/negative charge Can be a cation or anion 2. BASES Contain metal with hydroxyl group CATION - If a neutral atom loses one or more electrons, it has Accept protons net positive charge. Soapy/slippery in water solutions ANION – If a neutral atom gained an electron it has a net Turns red litmus paper into blue negative charge. 3. SALTS Acid base reaction COMPOUNDS – substances whose molecules are made up of Combination of any positive and negative ions 2 or more kinds of atoms combined in definite proportions. except hydrogen & hydroxyl ion All compounds are molecules Electrolyte that yield neither hydrogen or Not all molecules are compounds hydroxide from dissolved molecule Water – example of compund; atoms that make up water are not 4. OXIDES the same. Consists of oxygen & other elements Binary compounds formed by the reaction of TYPES OF COMPOUNDS oxygen with other elements 1. Ionic compound Oxygen is highly reactive in nature transfer of electrons from an atom to another formed through attraction between positive & negatively TYPES OF ELECTRODES charged ion. 1. Anode If compound is made from metal & nonmetal Positively charged electrode. 2. Cathode Negatively charged electrode. Ex: Na + O = Na2O (product is neutral compound) Mg + O = MgO (product is neutral compound) o Opposite electrical charges attract o Electrolysis can also take place in ionic solutions 2. Covalent compound o The more concentrated the solution the greater ion formed by sharing of electrons between 2 atoms. fluorate. The ion fluorate can be increased by 2 nonmetal ions. increasing the potential difference. PRINCIPLES EX: C + O = CO2 H + O = H2O Law of Definite Proportion AKA “Proust’s Law” BINARY COMPOUNDS Chemical compound always contain exactly the Compounds containing 2 elements same proportion of elements by mass. Can either have ionic or covalent bonding By Joseph Lewis Proust – he found out that copper carbonate is obtained resources is always composed of 5.3 parts of copper & 1 part of carbon & 4 parts of oxygen. Law of Multiple Proportion 3. Metallic compound States that when 2 elements form more than 1 Formed from interactions between elements which compound the ratio of the masses of thr second are metallic but resulting compound became like an element affect mass of the first with the ratio of a ordinary metal small whole number. Contains metals & nonmetals or polyatomic ions 2 | PAGE  Ex: CO2 when decomposed has 32 grams of oxygen 1. J.J THOMPSON (Thompson Model) atom is a for every 12 grams of carbons. spherical mass containing electrons & this spherical mass is positive but is made neutral by the electrons Law of Combining Weights embedded in it. Like a plum pudding. An atom is Proportions by weight when chemical reactions take hollow but filled. AKA “Raisin Bread Model” because place can be expressed in terms of small, integral, negatively charged particles in the positively charged multiple, fixed number called combining weights or particles. Electrons are randomly spread in a cloud of equivalent weights. massless positively charged material. AKA “Law of Reciprocal Proportions” AKA “Law of Equivalence” ATOM – fundamental unit of structure of matter that can enter into a chemical reaction. SUB-ATOMIC PARTICLES OF ATOM 1. PROTON Positively charged 2. ERNEST RUTHERFORD (Rutherford Model) based Discovered by Eugen Goldstein on additional experimental evidence of “alpha scattering elements” The positive charges/protons are 2. NEUTRON concentrated in the nucleus, and the region outside the Uncharged nucleus are the negatively charged/electrons. Ernest Same mass with proton Rutherford used the Geiger Marsden experiment to Discovered by James Chadwick prove most of atoms are in very small space called the atomic nucleus. Rutherford took a photo plate & 3. ELECTRONS surrounded it with gold foil & shot alpha particles, Negatively charged calling it Gold Foil Experiment. ATOMIC MASS o AKA atomic weight o Equal to number of protons + no. of neutrons ATOMIC NUMBER o No. of protons o No. of protons are equal to number of electrons ISOTOPES 3. NEIL BOHR (Bohr Model) also known as Planetary o Atoms of same element having the same atomic Model. In this model, protons are in the nucleus and number, but different atomic mass the electrons are in the orbitals that surround the nucleus. This model showed that the electrons orbit the ISOBAR nucleus in fixed circular orbits. o Atoms of different elements having the same atomic mass ISOTONES o Atoms of different elements having the same number of neutrons ATOMIC THEORY & STRUCTURES Scientist who studied & developed the structure of the atoms: 1. DEMOCRITUS termed the word “atomos” meaning tiny, indivisible particles. (Unable to divide or separate). Considered as 90% right. 2. JOHN DALTON proposed on an atomic theory (John Dalton Theory) based on the facts and experimental evidence that atoms enter into combination with other atoms to form compounds but will remain unchanged during ordinary chemical reactions. Atom is like a billiard ball, it is hard, indestructible sphere. It supports the law of definite proportion & law of multiple 4. RUTHERFORD-BOHR MODEL states that atoms are proportions. in elliptical orbits of increasing number. ATOMIC STRUCTUR 3 | PAGE  Indicates the spin of an electron in its own axis whether clockwise or counterclockwise direction 8. ORBITAL THEORY States that the number of orbital types in a given shell is equal to the shell number. Orbitals have a 3 dimensional region in space where the probability of finding the electron is greatest. HUND’S RULE OF MAXIMUM MULTIPLICITY 5. HEISENBURG UNCERTAINTY PRINCIPLE states Orbitals with same electron of same energy that simultaneous determination of the exact position level must be filled in seemly before the and exact momentum of electron is impossible. baring. 6. WAVE MECHANICAL ATOM AKA Electron Cloud 9. ELECTRONIC CONFIGURATION THEORY Model the nucleus is a single cluster of particles at Distribution of electrons in the different shells the center of the atom while the electrons are & subshells of orbitals within the atom. everywhere in rotating motion. 1st MAIN ENERGY LEVEL have the maximum number of 2 electrons Maximum number: 2 (1s2) 2ND MAIN ENERGY LEVEL Maximum number: 8 (2s2, 2p6) 3Rd MAIN ENERGY LEVEL Maximum number: 18 (3s2 3s2 3p6 3d10) PAULI’S EXCLUSION PRINCIPLE is formulated by 4th MAIN ENERGY LEVEL Wolfgang Pauli (1925), stating that no 2 electrons in Maximum number: 32 (s2, p6, d10, f14) an atom can have the same set of quantum numbers. S orbital – can only have 2 7. SCHROEDINGER “QUANTUM MODEL” formulated P orbital – can only have 6 by Erwin Schroedinger. This theory makes the D orbital – can have 10 assertion that electromagnetic ration like X-rays. Gamma rays, radio waves and light rays are made up F orbital – can have 14 of small bits of energy. It does not define the exact path of an electron but rather it predicts the odds AUFBAU PRINCIPLE of the location of the electron. This model can be portrayed as a nucleus surrounded by an electron cloud, where the cloud is most dense. The electron is likely in a less dense area of the cloud. This introduced the concept of sub-energy levels. QUANTUM NUMBERS – electron wave dimensions indicated by numbers quantum numbers 1. Principal Quantum Number Represented by small letter “n” Corresponds to main energy level Determines the size of particle Relates average distance of electrons from nucleus 2. Azithmuthal Quantum Number Represented by small letter “l” Atoms built by progressively filling of the main energy Gives and measures the angular momentum of an electron in its motion about the nucleus levels, sublevels & orbitals with electrons according to increasing sequence 3. Magnetic Quantum Number Levels of lower energy are occupied first (1s->2s->3p->) Represented by small letter “m” Ex: Indicates the behavior of electrons in the magnetic Aluminum (13 atomic number) = 1s2 2s2 2p6 3s2 film 3p1 Chlorine (Cl-1) (17 atom9ic number 🡪 18) = 1s2 Range is from –1 to +1 2s2 2p6 3s2 3p6 4. Spin Quantum Number TYPES OF CHEMICAL BONDS Represented by small letter “s” 4 | PAGE  Chemical Bond – connection between electrons and atoms. At constant temp, volume is inversely proportional to These are forces that hold atoms together pressure 1. IONIC BOND AKA Electrovalent bond Transfer of 1 or more electrons from one atom to another CHARLES L AW Held by strong electrostatic forces which is more stable At constant pressure, volume is directly proportional to 2. COVALENT BOND absolute temp Sharing of electrons between two atoms 3. METALLIC BOND Electrostatic attraction of positive and negative ions 4. HYDROGEN BOND GAY-LUSSAC’S LAW OF COMB9INING VOLUMES At constant volume and mass, when temperature and Hydrogen ion in combination with other atoms pressure are measured, the ratio of the volumes of the 5. VAN DER WAALS reacting gases are small whole numbers Attractive forces of polar substances INTERMOLECULAR FORCES OF ATTRACTION VAN DER WAALS 1. London/Dispersion Forces IDEAL GAS LAW/ GENERAL GAS LAW o Temporary attractive force that results when Pressure, temp and volume of gas is are related to each electrons of two adjacent atoms occupy positions other that make atoms form temporary dipoles o Weakest IMF o AKA Induced dipole-induced dipole attraction 2. Keesom Forces o AKA Dipole-dipole attraction COMBINED GAS LAW o Attraction between two polar molecules Volume is inversely proportional to pressure and directly 3. Keesom Forces proportional to temp o AKA Dipole-induced dipole attraction o Between polar and nonpolar molecule HYDROGEN BONDS Dipole-dipole attraction Strong electronegative elements ION-INDUCED DIPOLE DALTON’S LAW OF PARTIAL PRESSURE Weak attraction that result when the approach of an ion Total pressure is equal to the sum of partial pressure of all induces a dipole in an atom or in a nonpolar molecule by components disturbing the arrangement of electrons in the nonpolar species CHEMICAL F ORMULAS AVOGADRO’S L AW STRUCTURAL FORMULA Equal volumes of different gases having the same temp and Shows how atoms are bonded to one another in a pressure will contain equal number of molecules molecule (drawing ng structure niya) MOLECULAR FORMULA Shows the exact number of atoms of each element in the smallest unit of substances Based on the actual molecule GRAHAM ’S LAW OF DIFFUSION EMPERICAL FORMULA The ratio of the diffusion of gas is inversely proportional to Tells which element are present and expressed in the the square root of their molecular weight simplest whole number ratio GAS LAWS BOYLE’S LAW 5 | PAGE  Evolution of gas Formation of precipitate Emission of light Generation of electricity Production of mechanical energy Absorption/liberation of heat Chemical Reaction – occurs when valence electrons around ACID AND BASE THEORY the nucleus interacts. This means removal of electrons and addition of electrons to a partly filled valence shell or ARRHENIUS T HEORY sharing a pair of electrons Based on aq soln Reactants – substances that enter a chemical reaction Acids dissociate to form H ions (H+) Products – substances formed Bases dissociate to form hydroxide ions (OH-) - Chemical equation – representation of chemical reaction BRONSTED & LOWRY T HEORY Acids donate protons (H+) to bases TYPES OF CHEMICAL REACTION Bases accept donated protons Direct union A + B 🡪 AB LEWIS T HEORY Reaction of two or more substances react to form Bases donate electrons to base one compound Acids accept electrons Ex. 2Ca + O2 🡪 2CaO Decomposition DEBYE-HUKEL T HEORY AB 🡪 A + B One compound decomposes to form two or more Simplified model of electrolyte soln that gave accurate substances prediction of mean activity coefficient for ions in dilute soln Provides a starting point for modern treatment of non- Types: ideality of electrolyte soln o Hydrates to water and anhydrous salt o Chlorates to chlorides plus oxygen ADDITIONAL LECTURE o Metal oxides to metal and oxygen MATTER o Carbonates to oxides and CO2 o Bicarbonates to oxide, water and CO2 Anything that occupies space and has mass Single replacement CHARACTERISTICS OF MATTER AX + B 🡪 A + BX) Double displacement 1. It has composition AX + BY 🡪 AY + BX 2. It has structure Ex. 2HNO3 + Ca(OH)2 🡪 Ca(NO3)2 + 2H2O 3. It involves change 4. It requires energy PROCESSES INVOLVED IN CHEMICAL CHANGE Mass – amount of matter present in a material 1. Oxidation – union of oxygen Weight – mass with pull of gravity 2. Reduction – removal of oxygen or hydrogen is added 3. Neutralization – acid + base = salt + water PROPERTIES OF MATTER 4. Hydrolysis – water + salt = acid + base 5. Saponification – alkali + fat/oil = soap + glycerol Intrinsic/Intensive 6. Fermentation – action of microorganism, produces Independent of amount of matter alcohol Ex. Density, specific gravity, melting point Extrinsic/Extensive COLLISION THEORY Dependent on mass or amount of matter Collision of particles provide energy required to break Ex. Weight, volume, pressure, heat content bonds CHANGES UNDERGONE BY MATTER Activation energy – minimum amount of energy required for successful collision Physical – change in physical property or phase without A chemical change takes place as a result of collision of changing its chemical composition. Ex evaporation, molecules Chemical – change in both intrinsic and extrinsic properties. The greater the number of collisions per unit time, the Change in chemical composition. Ex oxidation greater the conversion of initial substances into products – Nuclear change – change in structure, properties, greater the speed of reaction composition of the nucleus of an atom resulting in transmutation of the element into another element Note: Nuclear fission – splitting of an atom (heavy atom, 1. Matter consists of moving particles mass number > 200) 2. In a chemical reaction, bonds must be broken in order Nuclear fusion – union of two light atoms to form new ones 3. Energy for this comes from particle collision 4. Collisions have A variety of energy EVIDENCES OF CHEMICAL CHANGE 6 | PAGE  5. Reaction cannot occur unless the molecule possesses 🡪 ↓ increases sufficient energy to get over the activation energy barrier REQUIREMENT OF SUCCESSFUL COLLISION Sufficient energy to break chemical bonds Favorable geometry Favorable orientation Possess minimum energy (activation energy) REACTION RATES CAN INCREASE DUE TO: 1. More collisions 2. Harder collisions – greater collision energy 3. Lower activation energy which allows low energy collisions to be more effective PERIODIC TABLE GROUPINGS IN PERIODIC T ABLE A. Representative Elements Group 1A-7A Wide range of properties Some are metals. Nonmetals or metalloids. Solid, liquid or gas Outer s and p configuration are not filled B. Noble Gases Inert gases S and p sublevels completely full C. Transition Metals Outer s sublevel full and is now filling d sublevel Transition between metal and nonmetal Brightly colored HISTORY D. Inner Transition Antoine Lavoiser – 1789: list of 33 elements Filling the f sublevel Johann Wolfgang Dobereiner – 1829: triads Rare earth metals Leopold Gmelin – 1843: 10 triads, 3 by 4, 1 by 5 Lanthanides – rare earth metals Jean-Baptiste Dumas – 1857: relationship between Actinides – heavy rare earth metals, radioactive various groups of metals All elements after uranium are artificial August Kekule – 1858: carbon bonds to atoms with ratio of 1:4 John Newlands – 1864-1865: law of octaves KINEMATIC MOLECULAR THEORY Dmitri Mendeleev and Julius Lothar Meyer - 1869: Periodic law (order of increasing atomic mass) Explains phases of matter based on movement of Henry Mosely – modern periodic table (based on atomic molecules number) PERIODIC T RENDS Solid Liquid Gas Atomic Radius Holds shape Shape of Shape of Half the distance between two nuclei container container 🡪 ↓ increases Electronegativity Fixed volume Free surface Volume of Ability to attract an electron to itself container 🡪 ↑ increases Ionization potential Fixed volume Energy required to remove an electron. How strongly atoms hold electrons 🡪 ↑ increases Electron affinity SOLUTIONS Ability to accept electrons. How strongly atoms Solute + Solvent 🡪 Solution attract electrons 🡪 ↑ increases Nonmetallic character 🡪 ↑ increases Metallic character 7 | PAGE  Equal forward and backward reaction TYPES OF EQUILIBRIA Homogenous equilibria Reacting substances are on the same phase Heterogenous equilibria Reacting substances are in two or more phases LE CHATELIER PRINCIPLE When stress is applied to a system in equilibrium, the equilibrium will shift in such a manner as to relieve TYPES OF SOLUTIONS ACCORDING TO THE SOLUBILITY OF stress THE SOLUTE FACTORS AFFECTING EQUILIBRIUM CONSTANT Saturated solution Unsaturated 1. Concentration Supersaturated – causes precipitation Increase in reactant favors forward reaction (🡪) FACTORS AFFECTING SOLUBILITY Increase in product favors reverse reaction (🡪) 2. Temperature Nature of the solute and solvent Increase in temp favors endothermic reaction Solubility – amount of solute in grams that can be (absorbs heat). ΔH = + dissolved in 100 grams of solvent Miscibility – ability of one substance to mix with Decrease in temp favors exothermic (release heat). ΔH = - another 3. Pressure Temperature Decrease in pressure faovors die with more Solubility of gas decreases with an increase in temp moles of gas Exothermic – solubility decreases with increasing 4. Catalyst temp Equal effect on both forward and backward Endothermic – solubility increases with decreasing reaction temp Pressure LAW OF MASS ACTION Henry’s law – solubility of gas increases with increasing pressure The rate of chemical reaction is directly proportional to the Particle size/ surface area concentration of the reactant at a given temp ↓particle size = ↑surface area LAW OF CHEMICAL EQUILIBRIUM Presence of salt Salting out – salt decreases solubility Concentration of reactant and product are raised to the Salting in – salt increases solubility power corresponding to the coefficients in the balanced equation is equal to a constant COLLIGATIVE PROPERTIES OF SOLUTION Colligative property – depends on concentration of solute but not on nature of the chemical species Vapor pressure lowering Equilibrium Constant Solute will take up spaces at the surface of liquid, Large Keq = mostly products are present at equilibrium limiting solvent molecules at the surface Small Keq = mostly reactants are present at equilibrium Solute lowers number of solvent molecules coming and going, lowering equilibrium vapor pressure Between (close to 1) = contains both reactants and products Boiling point elevation Other Equilibnum Constants Boiling point – temp at which vapor pressure is equal A. Ion Product Constant for water (Kw) to external pressure (1 atm) Need to exceed boiling point temp in order to Product of H+ and OH- ion concentration in mole/liter evaporate liquid Kw = [H+] [OH-] Kw = 1.00x10E-14 ↑Vapor pressure = add heat to reach vapor pressure B. Ionization Constant (Ki) Adding solute increases boiling point Freezing point depression ↑Ki = stronger acid or base Ka is the equilibrium constant for the ionization of weak acid Freezing point – temp at which solid and liquid in water coexist in equilibrium Kb is the equilibrium, constant for the ionization of weak Lowering temp = creates order base in water Adding solute adds entropy (mas nagkakagulo. C. Solubility Product Constant (Ksp) Kailangan babaan pa ang temp para magkaorder) Common in heterogenous system Osmotic pressure ↑Ksp = more soluble the salt Semi-permeable membrane Equilibrium constant of a slightly soluble salt\ CHEMICAL EQUILIBRIUM Ksp > Ki = no ppt (unsaturated soln) Ksp < Ki = ppt forms (supersaturated) State in which both reactant and products are present in Ksp = Ki – in equilibrium concentrations which have no tendency to change 8 | PAGE  T HERMOCHEMISTRY Organic Compounds Inorganic Compounds Study of heat and energy associated with a chemical Flammable Non-flammable reaction or a physical transformation Energy changes in a chemical reaction is primarily due to Low melting point High melting point breaking up of existing bonds between atoms Energy given out during chemical change will appear in the Low boiling point High boiling point form of heat Soluble in nonpolar Insoluble in non-polar All thermochemical reactions are governed by two laws solvent solvents T HERMOCHEMICAL EQUATION Insoluble in water Soluble in water An equation which indicates the evolution or absorption of heat in the reaction or process is called a thermochemical Contain covalent bonds Contains ionic bonds equation -organic compounds are Ex. C + O2 🡪 CO2 + 393.5kJ generally covalent LAWS OF T HERMODYNAMICS Organic compounds do Reactions take place Laplace Law not contain ions between ions A.L. Lavoisier and P.S. Laplace gave the law in 1780 Contains relatively many Contains relatively few which states that the enthalpy of a reaction is atoms atoms exactly equal but opposite in sign for the reverse reaction Most are complex Simpler structure Heat change is exactly equal but opposite in sign for the reverse reaction ΔH forward reaction = ΔH backward reaction Hess’s Law UREA G.H. Hess proposed a law regarding the heat or enthalpies of reaction in 1840. This law states that 1st organic compound synthesized heat change in a particular reaction is the same Most abundant organic compound whether intakes place in one step or several steps CHARACTERISTICS Amount of heat absorbed in a chemical reaction Catenation – ability to form long chains depends only upon the energy of the initial reactants Isomerism – compounds with same molecular formula exist and final reactants. Heat change is independent of in different forms owing to their different organization of the path or manner in which the change has taken atoms place Friedrich Wöhler produced urea In 1828 AKA carbamide T ERMINOLOGIES Constituent of urine that came from the reaction ammonium cyanate Heat (q) – an energy transfers due to temp difference Work (w) – A form of energy transfers between a system FUNCTIONAL GROUPS and its surroundings in the fdorm of compression and expansion of the gas C=C Alkenes Internal energy (U) – total energy attributed to the particles of matter and their interactions within the system, C≡c Alkynes composed of thermal energy and chemical energy OH Hydroxyl Enthalpy (H0 – energy of reaction Entropy (s) – degree of disorderliness C=O Carbonyl Heat capacity (c) – amount of heat required to raise the temp of an object or substance by one degree C=O-OH Carboxyl CHEMICAL REACTIONS NH2 Amine Endothermic – heat is absorbed by the system indicated C=N Nitriles by (+) change in enthalpy Exothermic - Heat is released by the system indicated by OR Alkoxy (–) change in enthalpy X Halide ORGANIC CHEMISTRY C=O-OR Ester Study of carbon containing compounds except carbonate, bicarbonates, cyanides and oxides SH Sulfhydryl O-C=O-R Carboxylate Catenation Ability to bond other carbon atoms to itself to form very large NO2 Nitro and complex molecules 9 | PAGE  ISOMERS Tautomers – two readily inconvertible structures that differ only in electron distribution and position in hydrogen atoms. These are compounds having the same molecular formula Proton transfer in an intramolecular fashion and same molecular weight but different structural formula, thus differ physical and chemical properties CLASSIFICATION 1. Constitutional; isomers o Same molecular formula but different connectivity 2. Stereoisomers/ Configurational Isomers o Atoms are connected in the same way but differ in how it is arranged in space TYPES OF ISOMERS Structural isomers/ Constitutional isomer – differs in Stereoisomers – isomers which differ from each other in order of bonding of atoms the arrangement of atoms A. Geometric isomers – stereoisomers that cannot be inconvertible without breaking a chemical bond Cis, E – same side Trans, Z – opposite side Butane (C4H10) ‘ A. Chain isomers – alteration in the way carbons are joined with no change in hybridization. Same molecular formula but differ in the Arrangement of carbon skeleton 2-methylpropane (C4H10) B. Positional isomers – a change in position of the double bonds or functional or substituent groups. Based on the movement of double bond or functional group B. Optical Isomers - Functional groups – part of a molecule that gives it Has asymmetric carbon/ chiral carbon reactivity Chiral carbon – carbon atom attached to 4 different types of atoms. 1-propanol 2-propanol C. Functional isomer – difference in type of formula. Same molecular formula but differs in functional group Optical activity – ability to rotate the plane polarized light due to chiral carbon Dextro (d) - right Levo (l) – left 1. Enantiomer – non superimposable mirror image Ethanol methoxymethane/ dimethyl ether 10 | PAGE  Compounds which contains carbon and hydrogen only 1. Aromatic hydrocarbon It is a hydrocarbon with alternating double and single bonds between carbon atoms Contains a six carbon atoms in aromatic compounds (benzene ring) Some non-benzene-based compounds called heteroarenes which follow Huckel’s rule are 2. Diastereomers – non superimposable non-mirror also aromatic compounds images Epimer CONDITIONS FOR AROMATICITY a. Should be cyclic b. Should be planar c. Should be conjugated d. It follows Huckel’s rule of aromaticity HUCKEL’S RULE Estimates whether a planar ring molecule will have an aromatic property. In these compounds, at least one carbon atom is replaced by one of the heteroatoms (O, N, S) Electrons = 4n + 2 Anomers – carbon bonded to 2 oxygen atom, usually carbon 1. Produced through the mutarotation when the sugar is suspended in aq soln producing alpha and beta. Detectable only if sugar is in cyclic form (Haworth projection) - Whole number = more stable - Fraction = unstable SULFUR-CONTAINING COMPOUNDS Compound Example Formula Mercaptans/ CH3SH R-SH Thiols Thioethers CH3CH2SCH3 R-S-R - OH up = beta, OH down = alpha (BUDA) Disulfides CH3CH2CH2SSCH2CH3 R-SS-R Racemic Mixtures Equimolar mixture of 2 enantiomers. Equal amount of Thiocyanates/ CH3CH2CH2SCN R-SCN dextro and levo rhodanide Optically inactive Thiophenol Ph-SH C. Conformational Isomer – a rotation that allows interconversion Envelope Chair Boat - Thiophenol – organosulfur compound NITROGEN-CONTAINING COMPOUNDS Compound Example Formula Hydrazine/ N2H4 diazine -used as foaming agent in HYDROCARBONS preparing polymer 11 | PAGE  foams, o Note: In naming, start the numbering with the precursor to nearest location of the double bond pharmaceut icals ALKYNES Having a triple bond between two carbons Nitro- Nitroglycerin, RNO2 CnH2n-2 compounds nitrobenzene sp hybrid IUPAC Naming ends with –yne Diazo R2C=N2 compounds 1-propyne Acid nitriles or R-C≡N cyanides o Note: In naming, start the numbering with the nearest location of the triple bond HYDROCARBON DERIVATIVES 2. Aliphatic Hydrocarbons Is a hydrocarbon containing carbon and Compounds made primarily of hydrogen and carbon in hydrogen joined together in straight chains, which it has a specific group of atoms that are attached. branched chains or non-aromatic rings This specific groups are called functional groups Contains at least 1 element that is considered to be 3 TYPES OF ALIPHATIC HYDROCARBON electronegative - Alkane/ Saturated HC One or more hydrogen atoms in the molecules is replaced o Characterized by single covalent bonds by certain group of atoms between carbon atoms o Cnh2n + 2 (R-H) HYDROXY DERIVATIVES o Sp3 hybrid o IUPAC naming ends with –ane ALCOHOLS o With –OH functional group o R-OH 5 4 3 o IUPAC Naming ends with –ol 2 1 Classification of Alcohols 2,3-dimethylpentane 1. According to the number of alkyl groups attached Alkane Series to the carbon bearing the hydroxyl groups A. Primary (1°) 1C Methane 10C Decane o -OH group is attached to 1 carbon that has 1 or no 2C Ethane 11C Undecane carbon atoms attached to it 3C Propane 12C Dodecane o o Ethanol 4C Butane 13C Tridecane B. Secondary (2°) 5C Pentane 14C Tetradecane o –OH group is attached to a carbon atom having 2 6C Hexane 15C Pentadecane carbon atoms attached to it 7C Heptane 20C Eicosane 8C Octane 30C Triacontane o o Isopropyl/2-propanol 9C nonane C. Tertiary (3°) TYPES OF ALIPHATIC HYDROCARBONS o –OH group is attached to a carbon atom having 3 ALKENES carbon atoms attached to it Characterized by having double bonds between carbon atoms CnH2n o o 2-methyl-2-propanol sp2 hybrid IUPAC Naming ends with –ene 2. According to the number of hydroxyl groups 3-methyl-1-butene A. Monohydric 12 | PAGE  o Presence of 1 -OH group ▪ Exist when one compound or the two attached is not the same o ▪ CH30CH3CH3 : Methoxy ethane B. Dihydric 1. Cyclic Ethers Ring o Presence of 2 -OH group o C. Polyhydric o Presence of 2 or more -OH group o o With 2 Carbons. Thus called cyclic ether o Name: Ethylene oxide AKA Epoxy ethane o CARBONYL COMPOUNDS PHENOLS o The –OH group is bonded to a conjugated cyclic A. ALDEHYDES planar system Formed by the oxidation of primary alcohol o Chemical behavior is different in some respect RCH=O from that of the alcohol. It is considered to be IUPAC Naming ends with –al sensible to treat them as similar but characteristically distinct group o More acidic but less basic than alcohols o Ar-OH Name: butanal/butyraldehyde ▪ Phenols are considered to be aromatic NUMBER OF CARBON ATOMS 1C- formaldehyde 6C- caproaldehyde 2C- acetaldehyde 7C- enantaldehyde 3C- propionaldehyde 8C- caprylaldehyde 4C- butyraldehyde 9C- pelargonaldehyde 5C- valeraldehyde 10C- caproldehyde o B. KETONES o –OH is connected to benzene ring Formed by the oxidation of secondary alcohol R-CO-R ETHERS IUPAC Naming ends with –one Formed by the reaction of alcohol with sulfuric acid that removes water from two molecules of alcohol Class of organic compounds characterized by an oxygen atom bonded to two alkyl or aryl groups o 2-pentanone R-O-R Note: In naming, start the numbering with the nearest IUPAC Naming ends with –oxy + Alkane series location of the double bond Carbonyl should be assigned in the lowest number of C atoms CARBOXYLIC ACIDS Name: ALPHABETICALLY Results from the oxidation of aldehyde o Common Name: Ethyl methyl ether R-COOH o IUPAC Name: Methoxy ethane IUPAC Naming ends with –oic acid CLASSIFICATION OF ETHERS 1. Open Chain o Butanoic acid Linear CH3CH2COOH o Symmetrical o Propanoic acid ▪ Made up of exactly similar parts facing each other or around its axis ▪ CH3OCH3 (Both are methyl) : Methoxy SATURATED ACIDS methane 1C Formic Acid o Asymmetrical 13 | PAGE  2C Acetic Acid R-CO-NH2 IUPAC Naming ends with –amide 3C Propanoic/Propioni c Acid 4C Butanoic Acid o Ethanamide 5C Valeric Acid ESTERS 6C Caproic Acid Produced by the reaction of organic acid with an alcohol R-COO-R 7C Enantic/Heptanoic IUPAC Naming the alkyl group R’ is named 1st followed by Acid the name of RCOO ending with –oate 8C Caprylic/Octanoic Acid 9C Pelargonic/Nonanoi o Propyl ethanoate c Acid ▪ Note: In naming, the –oate will be added in 10C Caprylic/Decanoic the group where COO is located Acid NITROGEN CONTAINING COMPOUNDS 11C Undecyclic/Undeca noic Acid AMINES Classified according to the number of C atoms that are 12C Dodecanoic/Loric bonded directly to a nitrogen atom Acid Considered basic compounds because it is Nitrogen containg compound 13C Tridecanoic Acid Resembles ammonia structurally where nitrogen can bind to 3 nitrogen atoms 14C Tetradecanoic Acid Organic compound derived from ammonia 15C Pentadecanoic R-NH2 Acid IUPAC Naming ends with –amine 16C Hexadecanoic Acid CLASSIFICATION OF AMINES Primary amine (1R) 17C Heptadecanoic Acid o 1 hydrogen attached to the amine is replaced by 1 alkyl or aryl group 18C Octadecanoic Acid o 20C Eicosanoic Acid ▪ Methylamine Secondary amine (2R) DICARBOXYLIC ACIDS o 2 hydrogen attached to the amine are replaced by o Organic compound that contains 2 carboxyl 2 organic substituents which can either be 2 alkyl functional groups or aryl group o Can be aliphatic or aromatic o COOH-R-COOH o DICARBOXYLIC ACIDS Tertiary amine (1R) 0CH2 Oxalic Acid/Ethanedioic acid 1CH2 Malonic Acid/Propanedioic acid o 3 hydrogen attached to the amine are replaced 2 2CH2 Succinic acid/Butanedioic acid organic substituents. It can be as an aromatic 3CH2 Pentanedioc Acid compound or aryl compound 4CH2 Hexanedioic Acid 5CH2 Heptanedioic Acid o 6CH2 Octanedioic Acid 7CH2 Nonanedioic Acid Quaternary amine (1R) 8CH2 Decanedioic Acid o 4 hydrogen attached to the amine are replaced 2 9CH2 Undecanoic Acid organic substituents. AMIDES Formed by the reactions of organic acids with ammonia or with amines o 14 | PAGE NITRILES Organic compound in which the carbon is triple bonded with a nitrogen Nitriles are usually known as the CYANIDES Smallest organic nitrile: Ethane nitrile Alkyl derivatives of hydrogen cyanide, HCN R-CN Named by giving the name of the alkyl group and ends with –nitrile o Propane nitrile ALKYL HALIDES AKA Halogenoalkanes or simply Haloalkanes Simplest alkyl haliude: Chloromethane Derivatives of alkanes in which one hydrogen atom in an alkane has been replaced by a halogen atom R-X o R: Alkyl radical o X- Halogenic compound Where R is any alkyl radical and X is any halogenide (F, Cl, I, Br) TYPES OF ALKYL HALIDE Primary Alkyl Halide (R-X) o Carbon which carries the halogen atom is only attached to 1 alkyl group o Secondary Alkyl Halide (R-CHX) o Carbon which carries the halogen atom is attached to 2 alkyl groups o Secondary Alkyl Halide (R-CX) o Carbon which carries the halogen atom is attached to 3 alkyl groups o 15 | PAGE PHARMACEUTICAL SEMINAR 1 PSMA411 | ORGANIC CHEMISTRY ORGANIC CHEMISTRY  Branch of chemistry that deals with carbon-containing compounds  except carbonates, bicarbonates, cyanides and oxides HISTORY  In 1828, Friedrich Wohler, a German chemist, disproved the “Vitalism” theory which states that all organic compounds come from living things. He was able to isolate DISTRIBUTION OF ORBITALS WITHIN SHELLS urea from an inorganic compound, ammonium cyanate  Each shell contains subshells known as atomic orbitals.  Uniqueness of Carbon  Electrons are said to occupy orbitals in an atom.  Carbon is able to form 4 covalent bonds (4 valence electrons) with other carbon or other elements.  Carbon atoms have the ability to bond to each other to form long chains or rings.  Ability to Catenate - Carbon atoms link together to form chains of varying length, branched chains and rings of different sizes ATOMIC STRUCTURE Shell Orbitals contained Maximum Relative  Elements: Fundamental building blocks of all substances In each shell number of energies of  Atoms: Smallest particle of an element electrons electrons in  Neutron: Neutral subatomic particle shell can each shell hold  Proton: Positively charged subatomic particle (+1 charge) 4 one 4s three 4p five 4d 2 + 6 + 10 +  Electron: Negatively charged subatomic particle (-1 seven 4f orbitals 14 = 32 charge) 3 one 3s three 3p five 3d 2 + 6 + 10 =  Nucleus: Center of an atom; contains protons and neutrons orbitals 18  An atom consists of a nucleus surrounded by electrons that 2 one 2s three 2p 2+6=8 are equal in number to the protons of the nucleus orbitals 1 one 1s 2 ATOMIC NUMBER AND ATOMIC MASS  The atomic number (Z) - number of protons in nucleus  The mass number (A) - number of protons plus neutrons ISOTOPES  atoms of the same element with different numbers of neutrons and thus different mass number (A). ELECTRON PRINCIPLES  AUFBAU PRINCIPLE - states that electrons fill lower- energy atomic orbitals before filling higher-energy ones  PAULI’S EXCLUSION PRINCIPLE - maximum of 2 electrons can occupy the same orbital only if they have opposite spins  HUND’S RULE - for degenerate orbitals, electrons fill the orbitals singly before they pair up  HEISENBERG’S UNCCERTAINTY PRINCIPLE - No 2 electrons can have the same set of 4 quantum numbers  Quantum Numbers SYMBOL VALUES FUNCTION ORBITALS 1. PRINCIPAL N 1,2,3 Determine the QN size of the  region of space where there is a certain probability of finding particle an electron Can hold 2 electrons Also known as WAVE FUNCTION 1|PAGE  2. l 0 to (n-1) Subshell or  O has two bonds and two unshared pair of electrons F, Cl, AZIMUTHAL sublevel, Br, and I have one bond and three unshared pairs of or ANGULAR determines electrons. the shape 3. MAGNETIC m or ml ml -1 to +1 orbitals, determines orientation 4. SPIN s or ms -1/2 to +1/2 direction of spin or orientation ELECTRON CONFIGURATION  symbolic notation of the manner in which the electrons of its atoms are distributed over different atomic orbitals  summary of where the electrons are around a nucleus RULES  Lowest-energy orbitals fill first: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 (Aufbau “build-up” principle)  Electrons act as if they were spinning around an axis. Electron spin can have only two orientations, up ↑ and down ↓. Only two electrons can occupy an orbital, and they must be of opposite spin (Pauli exclusion principle) to have unique wave equations.  If two or more empty orbitals of equal energy are available, KEKULE STRUCTURE electrons occupy each with spins parallel until all orbitals  Line bond structure have one electron (Hund’s rule)  Each shared electron is represented by line between the atom symbols CHEMICAL BONDING  Octet rule – atoms react in a way that achieve valence shell of eight valence electrons.  IONIC BOND  bond between anion and cation  atom may lose or gain enough electrons to acquire a completely filled valence shell COVALENT BONDING  anions (-) gain electrons; cations (+) lose electrons  joining of two atoms in a stable arrangement  Lone-pair electrons or non-bonding electrons  may occur between atoms of the same or different  Pair of valence electrons that are not used for elements. bonding  favorable process because it always leads to lowered energy and increased stability  Atoms form bonds because the resulting compound is more stable than the separate atoms  Ionic bonds in salts form by electron transfers  Organic compounds have covalent bonds from sharing electrons MULTIPLE BOND  When 2 atoms share more than 1 pair of electron FORMATION OF IONS  DOUBLE BOND  Octet Rule – The tendency among atoms of group 1A-7A  TRIPLE BOND elements to react in ways that achieve an outer shell of eight valence electrons. IDENTIFYING FORMAL CHARGES  Anion – an atom or group of atoms bearing a negative  Formal Charge - Associated with any atom that does not charge. exhibit the appropriate number of valence electrons.  Cation – an atom or group of atoms bearing a positive  1 st: Determine the number of valence electrons charge.  2nd: Determine whether the atom exhibits LEWIS STRUCTURE appropriate number of electrons  FC = [# valence e-] – [non-bonded e- + number of  Electron dot structure bonds]  Valence shell electrons of an atom are represented as dot  H has one bond ELECTRONEGATIVITY AND BOND POLARITY  C has four bonds  ELECTRONEGATIVITY - Measure of the ability of an atom  N has three bonds and one unshared pair of electrons to attract electrons INDUCTION AND POLAR COVALENT BONDS 2|PAGE  Difference in electronegativity < 0.5 HYDROGEN BONDING  Non-polar covalent bond  Hydrogen bonding is a complex interaction that includes  Equally shared electrons between the 2 atoms dipole-dipole, as well as orbital interactions and the transfer  Difference in electronegativity 0.5-1.7 of electron density between molecules.  polar covalent bond  These are the strongest of the IMFs and range from  not equally shared electrons between atoms 5 – 25 kJ/mol  INDUCTION withdrawal of electrons towards a highly  Occur primarily between OH, NH and FH. The more EN the electronegative atom which causes the formation of partial atom the stronger the interaction. (The atom H is attached charges to usually has a lone pair of e- ) DRAWING CHEMICAL STRUCTURES Shorthand ways of writing structures  Condensed Structure – C-H and C-C and single bonds are not shown but understood  If C has 3 H bonded to it, write CH3  If C has 2 H bonded to it, write CH2 and so on. Sometimes bonds between carbons are not shown in condensed structure- here the CH3, CH2 and CH units are simply drawn next to one another, but some bonds are added for clarity. The compound called 2- DIPOLE-DIPOLE methylbutane for example is written as follows  Dipole-dipole forces arise from the attraction of oppositely charged atoms (other than H) in molecules. These molecules may have a permanent dipole moment. Generally in organic molecules they results from the presence of C-X bonds where X is more electronegative that C.  These are generally weaker than H-bonding, ranging from about 5-10 kJ/mol. PHYSICAL PROPERTY  A property that does not affect the chemical identity of a VAN DER WAALS compound  Can be observed and measured without changing a  Van der Waals or (London) dispersion forces arise from compound’s composition of matter the movement of electrons within a molecule. This natural motion can produce an uneven distribution of the electrons  Any substance that has mass and can occupy space (polarization of the distribution) resulting in a temporary INTERMOLECULAR FORCES dipole moment in the molecule. This will induce the movement of electrons in adjacent molecules producing a  The physical properties of molecules are in part dependent dipole moment in them. on the type's of intermolecular forces (IMF) present.  These “induced” dipole moments are very brief as  Boiling points (BP) are also dependent on the mass of the they disappear when the electrons move to new molecule. \ locations within the molecule, so they forces are very  Solubility, the ability to dissolve into a solvent is dependent brief and weak, only 2-5 kJ/mol.\ on IMFs.  The strength of the interaction between molecules is also STRUCTURAL EFFECTS ON IMFS dependent on the overall shape of the molecule. There are  The strength of the IMFs depend on the amount of contact 3 types of IMFs, by decreasing strength they are: between the molecules, especially for dispersion forces. 1. Hydrogen bonding Hence the shape of the molecule can affect the surface 2. Dipole-dipole area of contact, long thin molecules have more surface in 3. Van der Waals or London Dispersion contact than spherical molecules. FACTORS AFFECTING THE PHYSICAL PROPERTIES OF ORGANIC COMPOUNDS 3|PAGE  Structure of Functional Group  Anything with a charged group (eg. ammonium,  Molecules having a polar functional group have a higher carboxylate, phosphate) is almost certainly water soluble, b.p. than others with a non-polar functional group of similar unless it has large nonpolar group, in which case it will most molecular masses. likely be soluble in the form of micelles, like a soap or detergent.  Any functional group that can donate a hydrogen bond to water (eg. alcohols, amines) will significantly contribute to water solubility.  Any functional group that can only accept a hydrogen bond from water (eg. ketones, aldehydes, ethers) will have a somewhat smaller but still significant effect on water solubility.  Other groups that contribute to polarity (eg. alkyl halides, thiols sulfides) will make a small contribution to water solubility. BOILING POINT AND MELTING POINT  Melting and boiling are processes in which noncovalent interactions between identical molecules in a pure sample are disrupted. The stronger the noncovalent interactions, the more energy that is required, in the form of heat, to break them apart. CHEMICAL PROPERTIES  A chemical reaction occurs when one substance is converted into another substance(s). A chemical reaction is accompanied by breaking of some bonds and by making of some others. REACTION MECHANISM  Define as the detailed knowledge of the steps involved in a process in which the reactant molecules change into products.  Chemical reactions involve breaking of one or more of the existing chemical bonds in reactant molecule(s) and LENGTH OF CARBON CHAINS formation of new bonds leading to products.  Molecules with higher molecular masses have higher m.p.,  The breaking of a covalent bond is known as bond fission. b.p. and density  During bond breaking or bond fission, the two shared  Higher molecular masses electrons can be distributed equally or unequally between the two bonded atoms.  Large molecular sizes  Stronger London dispersion forces among HOMOLYTIC FISSION molecules  Molecules with branched chains  The fission of a covalent bond with equal sharing of bonding  b.p. and density lower than its straight-chain isomer electrons.  Free radicals are neutral but reactive species having an  Straight-chain isomers have greater surface area in contact unpaired electron and these can also initiate a chemical with each other reaction.  Greater attractive force among the molecules  As a rule, larger molecules have higher boiling (and melting) points HETEROLYTIC FISSION  The fission of a covalent bond involving unequal sharing of bonding electrons. SOLUBILITY  This type of bond fission results in the formation of ions. The  If the solvent is polar, like water, then a smaller hydrocarbon ion which has a positive charge on the carbon atom, is component and/or more charged, hydrogen bonding, and known as the carbonium ion or a carbocation. On the other other polar groups will tend to increase the solubility. hand, an ion with a negative charge on the carbon atom is  The number of Carbons. More carbons means more of a known as the carbanion. non-polar/hydrophobic character, and thus lower solubility in water. 4|PAGE   An elimination reaction is characterized by the removal of a small molecule from adjacent carbon atoms and the formation of a double bond.  The charged species obtained by the heterolytic fission initiate chemical reactions and they are classified as electrophiles and nucleophiles. ADDITION  Electrophiles: An electrophile is an electron deficient species and it may be positively charged or neutral.  Unsaturated hydrocarbons such as alkenes and alkynes  Examples are H+ , AlCl3 , Br2 , Cl2 , Ag+ , CH3 +, BF3 etc. are extremely reactive towards a wide variety of reagents.  Nucleophiles : A nucleophile is negatively charged or The carbon-carbon double bond (–C=C–) of an alkene electron rich neutral species. contains two types of bonds. In alkynes, three carbon- carbon bonds.  Examples of nucleophiles are OH– , –NO2+ , H2O, :NH3 etc. MOLECULAR REARRANGEMENTS TYPES OF REACTIONS IN ORGANIC COMPOUNDS  proceeds with a fundamental change in the hydrocarbon skeleton of the molecule. During this reaction, an atom or SUBSTITUTION group migrates from one position to another.  A substitution reaction involves the displacement of one atom or group in a molecule by another atom or group. Aliphatic compounds undergo nucleophilic substitution reactions.  For example, a haloalkane can be converted to a wide variety of compounds by replacing halogen atom (X) with ISOMERS different nucleophiles as shown below.  These are compounds with the same molecular formula and same molecular weight but different structural formula, this differ in physical and chemical properties STRUCTURAL ISOMERS: CHAIN ISOMERS  Same molecular formula, but different arrangements of the carbon ‘skeleton’.  The positions of the carbon atoms can be rearranged to give ‘branched’ carbon chains coming off the main chain.