Chemistry Lecture Notes - PDF

Prof. Dr. Hossieny Ibrahim Badr University in Assiut School of Biotechnology [email protected] Office number: Bio-326 1  Ethics:  Mobiles should be silent (Please Check)  Questions only with permission  Side discussions are not allowed  Entrance &...

Prof. Dr. Hossieny Ibrahim Badr University in Assiut School of Biotechnology [email protected] Office number: Bio-326 1  Ethics:  Mobiles should be silent (Please Check)  Questions only with permission  Side discussions are not allowed  Entrance & leave should be quite and organized 2 Lecture Monday: 1:0 – 3:0 Room No. V-301 Assessment : 10 Marks: Quizzes 10 Marks: Midterm Exam 15 Marks: Oral Exam 15 Marks: Practical Exam 50 Marks: End Term Exam Textbooks: 1- Atkins P.W., 2006, Physical Chemistry, 8th ed., Oxford University Press. 2- McMurry & Thompson, Fundamentals of Organic Chemistry, Brooks-Cole 2002. 3  What is the main purpose for this course?  To know the basis of organic chemistry and its importance.  To know relationship between the molecular structure and properties of organic compounds.  To recognize the functional groups in organic compounds and classification of organic compounds.  To know the nomenclature of organic compounds by using IUPAC rules and common names.  To know relationship between the functional groups and reactivity of organic compounds.  To recognize the fundamental reactions in organic chemistry (substitution – Addition – elimination ).  To study the different sections of the aliphatic organic compounds in terms of the nomencluture , structure , physical properties , methods of preparation , chemical reactions and common uses.  Students learn the importance of :  Structure of atom.  Chemical bonding.  Chemical equilibrium. 4   Course Content 1. Introduction 1.1. Types of bonds 1.5. Resonance 1.2. Electronegativity 1.6. Hybridization 1.3. Lewis’s structure & Formal charge 1.7. Isomer 1.4. Lewis’s acid & base 1.8. Functional groups 2. Hydrocarbons 2.1. Aliphatic compounds: Alkanes & cycloalkanes – Alkenes – Alkynes 2.2. Aromatic compounds 3. Alcohols and Alkyl halides 4. Carbonyl compounds: Aldehydes & Ketones 5. Carboxylic acids and derivatives 6. Amines 7. Structure of atom and chemical bonding 8. Chemical Equilibrium 5  Introduction to Organic Chemistry  The early definition related to compounds obtained only from living things.  Today, it is a major branch of chemistry that deals with compounds of Carbon. It is such a complex branch of chemistry because.........  Carbon has a bonding capacity of 4.  The carbon-carbon bonds can be single, double or triple. C C C C C C  Carbon atoms can be arranged in: ❶ Straight chains ❷ Branched chains ❸ Rings 6   Introduction to Organic Chemistry  Other atoms/groups of atoms can be placed on the carbon atoms.  The basic atom is Hydrogen but groups containing oxygen, nitrogen, halogens and sulphur are very common. C C C C C C C C C C Carbon Functional Carbon Functional skeleton group skeleton group  Groups can be placed in different positions on a carbon Skeleton  The c=c double bond is in a different position.  The chlorine atom is in a different position. Cl Cl 7   The Significance of Carbon  Nearly 16 million carbon-containing organic compounds are known  Types of Carbon compounds in organisms include Carbohydrates, Lipids, Proteins, and Nucleic acids.  Carbohydrates: Contain only carbon, hydrogen, and oxygen.  All consist of one or more smaller units called monosaccharides.  Lipids: Contain carbon, hydrogen, and oxygen.  Lipid molecules consist of glycerol & 3 fatty acids.  Proteins: Contain carbon, hydrogen, oxygen, nitrogen  Made of smaller units called amino acids.  Nucleic acids: Contain carbon, hydrogen, oxygen, nitrogen, and phosphorus. 8 NOTE Carbon compounds that are exceptions and considered Inorganic are compounds like: carbon monoxide [CO(g)] and carbon dioxide [CO2(g)], carbonate (CO32- ), cyanide (CN-), and carbides [CaC2 (calcium carbide)] Problem Classify the following compounds as organic or inorganic. CH2O , CO2 , CO , CH4 , NaHCO3 , C2H6 , MgCO3 , C6H12O6 , KCN , CCl4 , CH3COOH , CaC2 9 Atomic Structure and Electronic Configuration Electronic Configuration of Elements 10 Atomic Structure and Electronic Configuration Nucleus Electron Orbit Energy Levels 11 12   Valence Electrons & Lewis structure  A valence electron is an outer shell electron that is associated with an atom, 5 B and that can participate in the formation of a chemical bond if the outer shell is not closed. Boron Atom  A Lewis structure shows the symbol of the element, surrounded by a number of dots equal to the number of electrons in the outer shell (Valence electrons) of an Atom of that element. Note: Group No. = No. of valence electron 13   Formation of Ions & Lewis Model of Bonding  The tendency of atoms to react in ways that achieve an outer shell of eight valence electrons is particularly common among elements of Groups 1A–7A (the main-group elements). We give this tendency the special name, the Octet Rule.  In gaining electrons, the atom becomes a negatively charged ion called an Anion.  An atom with only one or two valence electrons tends to lose the number of electrons required to have the same electron configuration as the noble gas nearest it in atomic number. In losing one or more electrons, the atom becomes a positively charged ion called a Cation. 14   Formation of Chemical Bonds  According to the Lewis model of bonding, atoms acquire completed valence shells in two ways: 1- Ionic bond: A chemical bond 2- Covalent bond : A chemical resulting from the electrostatic bond resulting from the sharing attraction of an anion and a cation. of one or more pairs of electrons. 15  Ionic Bond or a Covalent Bond  One way to answer this question is to consider the relative positions of the two atoms in the Periodic Table.  Ionic Bonds usually form between a metal and a nonmetal.  An example of an ionic bond is that formed between the metal sodium and the nonmetal chlorine in the compound sodium chloride (NaCl).  By contrast, when two nonmetals or a metalloid and a nonmetal combine, the bond between them is usually covalent.  Examples of compounds containing covalent bonds between nonmetals include Cl2, H2O, CH4, and NH3.  Examples of compounds containing covalent bonds between a metalloid and a nonmetal include BF3, SiCl4, and AsH4. 16   Electronegativity and Chemical Bonds  Another way to identify the type of bond is to compare the electronegativities of the atoms involved.  Electronegativity (EN) is an atom’s attraction for the electrons in a covalent bond. The more electronegative an atom, the more strongly it pulls shared electrons toward itself. 17   Polar Covalent Bonds Polar covalent bond: is a bond formed when a shared pair of electrons are not shared equally. We can show the presence of a bond dipole by an arrow. 18 Example: Classify each bond as nonpolar covalent, polar covalent, or ionic: (a) O—H (b) N—H (c) Na—F (d) C—Mg 19  Problem Using a bond dipole arrow ( ) and the symbols δ- and δ+, indicate the direction of polarity in these polar covalent bonds: (a) C—O (b) N—H (c) C—F (d) C—N (e) N—O (f) C—Cl Problem Arrange the single covalent bonds Covalent bond (CB) within each set in order of increasing Polar Covalent Bond (PCB) polarity: (a) C—H, O—H, N—H (b) C—H, C—Cl, C— F (c) C—C, C—O, C—N 20   Lewis Structures of Molecules and Ions Rules: 1- Total Number of Valence Electrons (TNVE)  TNVE is the sum of group number for each atom + (negative charge). – (positive charge). So, for H2O, the TNVE = 2 x 1 + 1 x 6 = 8 electrons CO2 : TNVE = 1 x 4 + 2 x 6 = 16 electrons CO32- : TNVE = 1 x 4 + 3 x 6 + 2 (-ve charge) = 24 electrons 2- Octet Rule : every atom should have an octet (8) electrons associated with it. Hydrogen should only have 2 (a duet).  2 x no. of H-atoms + 8 x any atom Ex. H2O : 2 x 2 + 8 x 1 = 12 electrons CO2 : 2 x zero + 8 x 3 = 24 electrons CO32- : 2 x zero + 8 x 4 = 32 electrons 21 3- Number of bonding electrons (NBE) = Step (2) – Step (1) 4- Number of bonds = Step (3) ÷ 2 5- Determine which atom is the “central” atom. The central atom is the least electronegative element that isn’t hydrogen. 6- Stick everything to the central atom using a single bond. 7- Fill the octet of every atom by adding dots. 8- Check the “Formal charge” (FC) of each atom. Formal charge (FC) = Gp.No. ̶ (No. of bonds + No. of unshared electrons)  H has one bond.  C has four bonds.  N has three bonds and one unshared pair of electrons.  O has two bonds and two unshared pairs of electrons.  F, Cl, Br, and I have one bond and three unshared pairs of electrons. 22 Note: Bonding electrons: Valence electrons shared in a covalent bond. Nonbonding electrons: Valence electrons not involved in forming covalent bonds, that is, unshared electrons.  Nonbonding electrons  Unshared electrons  Lone pairs Bonding electrons Bonding electrons 23  Writing Lewis Structures: Ex. PCl3 1- Total Number of Valence Electrons (TNVE) 1 x 5 + 3 x 7 = 26 e 2- Octet rule: 2 x zero + 8 x 4 = 32 e 3- NBE = 32 – 26 = 6 e 4- No. of bonds = 3 bonds  P is the central atom FC (P) = 5 – (3 + 2) = zero Fill the Octet of every atom by adding dots. Calculate the Formal charge (FC) of each atom FC (Cl) = 7 – (1 + 6) = zero 24  Writing Lewis Structures: Ex. O3 1- Total Number of Valence Electrons (TNVE) 3 x 6 = 18 e 2- Octet rule: 2 x zero + 8 x 3 = 24 e 3- NBE = 24 – 18 = 6 e FC (O) = 6 – (3 + 2) = +1 4- No. of bonds = 3 bonds Fill the Octet of every atom by adding dots. FC (O) = 6 – (1 + 6) = -1 Calculate the Formal charge (FC) of each atom Note: Total Formal charge = Zero FC (O) = 6 – (2 + 4) = 0 25  A Summary of Common Formal Charges 26  Problem  Calculate the formal charges on the atoms shown in red: (a) (c) (b) (d) (e) Problem  Assign formal charges to the atoms in each of the following molecules: (a) (b) (c) 27 Problem Following the rule that each atom of carbon, oxygen, and nitrogen reacts to achieve a complete outer shell of eight valence electrons, add unshared pairs of electrons as necessary to complete the valence shell of each atom in the following ions. Then, assign formal charges as appropriate 28   Lewis Acids (LA) and Lewis Bases (LB)  In 1923 G. N. Lewis proposed a theory that significantly broadened the understanding of acids and bases.  A Lewis acid is a substance that accepts an electron pair.  A Lewis base is a substance that donates an electron pair. LB LA Write an equation that shows the Lewis acid and Lewis base in the reaction of bromine (Br2) with ferric bromide (FeBr3). LA-LB adduct LB LA 29   Lewis Acids (LA) and Lewis Bases (LB)  Some further examples of Lewis acids follow: 30   Lewis Acids (LA) and Lewis Bases (LB)  In a more general sense, most oxygen- and nitrogen-containing organic compounds can act as Lewis bases because they have lone pairs of electro 31  Using curved arrows, show how acetaldehyde, CH3CHO, can act as a Lewis base.  Strategy A Lewis base donates an electron pair to a Lewis acid. We therefore need to locate the electron lone pairs on acetaldehyde and use a curved arrow to show the movement of a pair toward the H atom of the acid. curved arrow LA LB 32 Problem Write equations showing the Lewis acid–base reaction that takes place when: (a) Methanol (CH3OH) reacts with BF3. (b) Chloromethane (CH3Cl) reacts with AlCl3. (c) Dimethyl ether (CH3OCH3) reacts with BF3. (d) Trimethyl amine [ (CH3)3N ] reacts with B(CH3)3 (e) Ammonia (NH3) reacts with AlCl3.  Add curved arrows to the following reactions to indicate the flow of electrons for all of the bond-forming and bond-breaking steps. 33   Resonance Structures  In cases in which more than one reasonable Lewis structure can be drawn for a species, these structures are called resonance structures or resonance contributors.  Let’s consider the example of the carbonate anion, CO32- 34  Draw all the equivalent resonance structures for following species: Problem (a) O3 molecule (b) Nitrate anion NO3– (c) SO2 35