General and Inorganic Chemistry PHCM101 Lecture Notes PDF

General and Inorganic Chemistry PHCM101 Synthetic and Natural Organic Polymers Lecture 12 Dr. Nesrine El-Gohary 1  COMPETENCIES 1-1-1 Define polymers and different types of polymerization reactions 1-1-2 A...

General and Inorganic Chemistry PHCM101 Synthetic and Natural Organic Polymers Lecture 12 Dr. Nesrine El-Gohary 1  COMPETENCIES 1-1-1 Define polymers and different types of polymerization reactions 1-1-2 Articulate knowledge in differentiation between different stereoisomers. 1-1-3 Demonstrate understanding of the different types of polymers (natural and synthetic). 2-2-1 Differentiate between thermoplastic and thermoset polymers. 2  Polymers A polymer is a large molecule (macromolecule) of high molecular weight made up of many repeated chemical units called monomers. Naturally occurring polymers Proteins Dacron Nucleic acids Cellulose Rubber Teflon Synthetic polymers Nylon Dacron (polyester) Kevlar Teflon Kevlar 3  Classification of polymers Polymers are classified according to: 1- Synthesis method (addition polymers or condensation polymers). 2- Origin (natural or synthetic). 3- Physical properties (thermoplastic or thermoset). 4  Revision on some basic organic Chemistry concepts Alkanes are the simplest organic molecules, consisting of only carbon and hydrogen and with only single bonds between carbon atoms. (All bonds in alkanes are sigma bonds). Alkenes are hydrocarbons that contain a carbon-carbon double bond (Molecule contains a pi bond). Ethene (ethylene) Alkynes are hydrocarbons containing at least one carbon—carbon triple bond between two carbon atoms. (Molecule contains 2 pi bonds) Ethyne (acetylene) 5 Revision on some basic Organic Chemistry concepts (cont.) A molecule or an atom having an unpaired valence electron is called a FREE RADICAL. Free radicals are given the symbol R Free radicals are highly reactive, the single unpaired electron is in desperate need to be paired with another electron. FREE RADICALS are a must in free radical addition polymerization THEY ATTACK WEAK BONDS (Pi bonds) present in alkenes or alkynes 6 Revision on some basic Organic Chemistry concepts (cont.) NH2 Carboxylic group Amino group H2O Amide 7 Revision on some basic Organic Chemistry concepts (cont.) HO Carboxylic group Alcohol group H2O Ester 8 Synthesis of polymers 1- Addition Polymerisation (Addition polymers) The monomers simply add together to form the polymer. A carbon – carbon double/ triple bond is needed in the monomer. ethylene polyethylene 9 Mechanism of addition polymerization Initiation light / heat initiator Propagation Repitition of this process thousands of time creates a long-chain polymer. Termination repeating unit (monomer) 10  Addition Polymerisation The polymer is the only product. Reaction proceeds by a free radical mechanism. Involves the opening out of a double / triple bond, where the free radical attacks the pi bond. Conditions are high pressure, temperature and an initiator (to provide the initial free radical). Peroxides(R-O-O-R) are often used as the initiator/catalyst (source of free radical). The conditions of the reaction can alter the properties of the polymer. 11  The simple repeating unit of a polymer is the monomer. Homopolymer is a polymer made up of only one type of monomer Tetrafluoroethylene monomer Ethylene monomer Vinyl chloride monomer ( CF2 CF2 )n ( CH2 CH2 )n ( CH2 CH )n Teflon Polyethylene Cl 12 Polyvinyl chloride (PVC) Copolymer is a polymer made up of two or more monomers n x Styrene monomer Butadiene monomer 25% 75% ( CH CH2 CH2 CH CH CH2 )n Styrene-butadiene rubber 13  Poly(propene) and stereoisomerism Stereoisomers: Two molecules are described as stereoisomers of each other if they are made of the same atoms, connected in the same sequence, but the atoms are positioned differently in space. Isotactic: This is a very regular type of polymer chain. All the methyl groups are on the same side, so the polymer chains pack closely together leading to strongest polymer with highest melting point when compared to other stereoisomers. H CHH CHH CHH CHH CHH CHH CHH CH 3 3 3 3 3 3 3 3 14 Poly(propene) and stereoisomerism Syndiotactic: A slightly less regular but still very ordered polymer. The methyl groups alternate the side of chain they are on. H CH CH H H CH3H H H H 3 3 CH3 CH3 CH3 CH3 15 Poly(propene) and stereoisomerism Atactic: This is a completely random allocation of methyl groups along the carbon skeleton, weakest polymer, lowest melting point. H H H H H H H 16  Poly(propene) and stereoisomerism The varying degree of randomness will affect the strength and melting point of the polymer. The less random, the stronger the polymer and the higher the melting point. This is because in a more ordered polymer the chains can get closer together and hence the intermolecular forces will be greater. By adding certain catalysts, the polymerization could be directed towards the formation of the more ordered polypropene (polypropylene). 17 1- Isotactic, atactic, syndiotactic 2- Syndiotactic, isotactic, atactic 3- Syndiotactic, atactic, isotactic 4- Atactic, isotactic, syndiotactic 5- Atactic, syndiotactic, isotactic  Synthesis of polymers 2- Condensation Polymerisation (Condensation polymers) Involves 2 monomers that have different functional groups. They also involve the elimination of water or another small molecule. Hence the term condensation polymer. Monomer A + Monomer B → Polymer + small molecule (normally water). Common condensation polymers include polyesters (the ester linkage) and polyamides (the amide linkage as in proteins). 19  Polyesters The Dacron is a polyester, a polymer of benzene-1,4- dicarboxylic acid and ethane-1,2-diol (ethylene glycol). The ester linkage is formed between the monomers. O O n HO C C OH + n HO CH2 CH2 OH heat with an acid catalyst O O C C O CH2 CH2 O n poly(ethan-1,2-diyl benzene-1,4-dicarboxylate) 20 Polyamides These involve the linkage of two monomers (amine and an acid) through the amide linkage as in nylon. H H O O N (CH2)6 N C (CH2)4 C H H HO OH 1,6-diaminohexane hexanedioic acid O O N (CH2)6 N C (CH2)4 C H H part of a nylon polymer chain 21 Nylon 6,6 a polyamide H H O O N (CH2)6 N C (CH2)4 C H H HO OH 1,6-diaminohexane hexanedioic acid O O N (CH2)6 N C (CH2)4 C H H part of a nylon polymer chain 22  Kevlar polyamide Benzene 1,4 dicarboxylic acid 1,4 diaminobenzene Uses of polyesters and polyamides The main use of polyesters and polyamides is as fibres in clothing. Most clothing now has a degree of manufactured fibres blended into the natural material (such as cotton). This gives the material more desirable characteristics, such as stretchiness. 24  Natural and synthetic polymers 1- Natural polymers (Proteins) Condensation polymer, the repeating units are amino acids. All amino acids contain at lease one carboxylic group and one amino group. Proteins are polyamides specifically called polypeptides. H O H O Play a key role in nearly all biological processes H2N C C OH + H2N C C OH − Enzymes, the catalysts of R1 R2 biochemical reactions. Peptide (amide) bond −Storage and transport of materials. H O H O − Hormones. H2N C C N C C OH + H2O − Mechanical support. R1 H R2 − Protection against diseases. 25 26  Natural and synthetic polymers (cont.) 2- Natural polymers (Nucleic acids :DNA and RNA) Condensation polymers, the repeating units are nucleotides. Nucleic acids are polynucleotides. Nucleotide These molecules contain all the information required to build a functioning organism, like a human body. DNA and RNA contain the information required to build cells, assemble organs, grow, and reproduce. 27  Natural and synthetic polymers (cont.) 3- Natural polymers (Cellulose) Condensation polymer, made of a long chain of repeating glucose units. Cellulose is a polysaccharide. It is the major component of woody plants and cotton. It has remarkable strength, so used in textile industry and in manufacturing of paper. Natural cellulose could be chemically modified by treatment with different reagents. Example: Treatment with nitric acid gives cellulose nitrate (guncotton), which is used as gun 28 powder to make explosives. Natural and synthetic polymers (cont.) 4- Natural polymers (Rubber) Addition polymer, made of repeating units of a monomer called isoprene. Isoprene Natural rubber is obtained from the latex of rubber tree. Latex is an aqueous suspension of rubber particles. https://www.youtube.com/watch?v=5iTz9yN4v4k Natural rubber is Tough Elastic (it will stretch up to 10 times its length, and when released will return to its original size) so rubber is called an Elastomer. Soft and sticky when hot, which makes it unsuitable to be used in the manufacture of tires. 29  Natural and synthetic polymers (cont.) In 1839, Goodyear, an American chemist, accidentally found that if sulfur is added to rubber and the resulting mixture is heated, the rubber became much more stronger but still elastic so could be used in tires. A process called vulcanization. In Vulcanization, S atoms become bonded between carbon atoms on different chains, so they act as bridges between the polymer chains, thus linking the chains together. Rubber molecules ordinarily are bent and convoluted. Parts (a) and (b) represent the long chains before and after vulcanization, respectively; (c) shows the alignment of molecules when stretched. Without vulcanization these molecules would slip past one another, and rubber's 30 elastic properties would be gone. Natural and synthetic polymers (cont.) Synthetic polymers All examples that were previously discussed: Polyethene Polypropylene Dacron Nylon Kevlar Styrene-butadiene rubber (synthetic rubber) which is more resistant to abrasion and oxidation than natural rubber.  Thermoplastic and Thermoset polymers THERMOPLASTIC THERMOSET Polymer that can be Polymer that cannot be softened by heating softened by heating and and then formed into so cannot be remolded desired shapes by applying pressure Recyclable. Non recyclable. 32  Exercise: 1- Which one of the following is an addition polymer a- Protein b- Cellulose c- Rubber d- Nylon 2- Nylon threads are made of a- Polyester polymer b- Polyethylene polymer c- Polyamide polymer d- Polypropylene polymer O O n HO 3-C Addition polymerization C OH + n HO results in2 elimination CH2 CH OH of a small molecule during polymer formation heat with an acid a- True catalyst b- False O O 4- C C O CH2 CH2 O n a- Polyester polymerbenzene-1,4-dicarboxylate) poly(ethan-1,2-diyl b- Polyethylene polymer c- Polyamide polymer d- Polypropylene polymer 33  References and resources Raymond Chang “Chemistry”, 11th Ed. The McGraw Hill Companies, ISBN 978-0-07-017264-7. Chapter 25 34 General and Inorganic Chemistry (PHCM101) Lecture 7 Bonding: General Concepts II (Molecular Geometry and Polarity) Dr. Nesrine El Gohary Office: B5.104 [email protected] Office hours: Monday 4th slot Tuesday 4th slot  COMPETANCIES 1-1-1 Demonstrate understanding of the Valence-Shell Electron Pair Repulsion Model (VSEPR) 2-2-1 Predict the effect of unshared pairs on geometry 2-2-2 Predict the effect of multiple bonds on geometry 2-2-3 Predict the polarity of molecules 2 Recap: Drawing Lewis structure 1. Count the number of valence electrons. 2. Draw a skeleton structure for the species, joining atoms by single bonds. 3. Determine the number of valence electrons still available for distribution. 4. Determine the number of valence electrons still required to fill out an octet for each atom (except H) in the skeleton. 5. Determine the difference between available and required valence electrons. 3 Recap: Drawing Lewis structure Available – Required = Zero -2 -4 > zero Draw the Draw One Draw one Put extra e’s on structure without Double bond triple bond or the central atom, any multiple two Double which will exceed bonds bonds the octet rule NCl3 HCO2- & NO3- H2C2 & OCS ClF2- & IF4- Molecular Geometry It is the general shape of a molecule, as determined by the relative positions of the atomic nuclei. i.e it is the three dimensional arrangement of atoms in a molecule. The geometry of a diatomic molecule such as Cl2 or HCl can be described very simply. Because two points can define a straight line, the molecule must be linear. H Cl With molecules containing three or more atoms, the geometry is not obvious. Here the angles between bonds, called bond angles, must be considered. Molecular Geometry It is the general shape of a molecule, as determined by the relative positions of the atomic nuclei. i.e it is the three dimensional arrangement of atoms in a molecule. ⮚For example, the geometry of a molecule of the type YX2, where Y represents the central atom and X is an atom bonded to it, could be either : Linear : Bent: with a bond angle of 180° with a bond angle < 180° Y X Y X X X 6 Valence-Shell Electron Pair Repulsion Model (VSEPR) This model predicts the shapes of molecules and ions by assuming that valence shell electron pairs are arranged about each atom so that they are kept as far away from one another as possible, thus minimizing the electron-pair repulsions. VSEPR model is used in predicting the geometry of some simple molecules and polyatomic ions in which a central atom is surrounded by 2 - 6 pairs of electrons. 7 Molecular Geometry Molecules with the general formula AX2 Lewis structure shows a central atom surrounded by two bonding pairs of electrons. i.e. 1 atom is a central atom (A) and 2 atoms are terminal atoms (X). Molecular geometry is Linear Bond angle = 180o Example: BeH2 8 Molecular Geometry Molecules with the general formula AX3 Lewis structure shows a central atom surrounded by 3 bonding pairs of electrons. i.e. 1 atom is a central atom and 3 atoms are terminal. Molecular geometry is Triangular (Trigonal) Planar Bond angle = 120o Example: BF3 9 Molecular Geometry Molecules with the general formula AX4 Lewis structure shows a central atom surrounded by 4 bonding pairs of electrons. i.e. 1 atom is a central atom and 4 atoms are terminal. Molecular geometry is Tetrahedron Bond angle = 109.5o Example: CH4 10 Molecular Geometry Molecules with the general formula AX5 Lewis structure shows a central atom surrounded by five bonding pairs of electrons. i.e. 1 atom is a central atom and 5 atoms are terminal. Molecular geometry is Trigonal bipyramid Bond angle = 90o , 120o , 180o Example: PF5 11 Molecular Geometry Molecules with the general formula AX5 Lewis structure shows a central atom surrounded by five bonding pairs of electrons. i.e. 1 atom is a central atom and 5 atoms are terminal. Molecular geometry is Trigonal bipyramid Bond angle = 90o , 120o , 180o Example: PF5 12 Molecular Geometry Molecules with the general formula AX6 Lewis structure shows a central atom surrounded by six bonding pairs of electrons. i.e. 1 atom is a central atom and 6 atoms are terminal. Molecular geometry is Octahedron Bond angle = 90o, 180o Example: SF6 13 Molecular Geometry Molecules with the general formula AX6 Lewis structure shows a central atom surrounded by six bonding pairs of electrons. i.e. 1 atom is a central atom and 6 atoms are terminal. Molecular geometry is Octahedron Bond angle = 90o, 180o Example: SF6 14 Molecular Geometry Summary Species type Orientation of Predicted bond Ball and stick model electron pairs angles AX2 Linear 180o Trigonal AX3 120o planar AX4 Tetrahedron 109.5o Trigonal AX5 90o, 120o, 180o bipyramid AX6 Octahedron 90o, 180o 15 Molecular Geometry Multiple Bonds ⮚The VSEPR model can be extended to species in which double or triple bonds are present. ⮚Electron pairs in a multiple bond must occupy the same region of space as those in a single bond. ⮚Hence, the extra electron pairs in a multiple bond has no effect on the geometry. The four electrons in a double bond, or the six electrons in a triple bond, must be located between the two atoms, as are the two electrons in a single bond. 16 Molecular Geometry Multiple Bonds Consider the CO2 molecule: The central carbon has two double bonds and no lone pairs ⮚ The double bond is counted as a single bond, ignoring extra bonding pairs. The bonds are directed to be as far apart as possible, giving a 180°O-C-O. CO2 and BeF2 are both linear molecules. 17 Molecular Geometry Effect of lone pairs ▪ In many molecules and polyatomic ions, one or more of the electron pairs around the central atom are unshared. The VSEPR model is readily extended to predict the geometries of these species. ▪ In describing molecular geometry, we refer only to the positions of the bonded atoms. These positions can be determined experimentally. ▪ Positions of unshared pairs cannot be established by experiment. Hence the locations of unshared pairs are not specified in describing molecular geometry. ▪ Electron-pair geometry: specifies the positions of all electron pairs (bonding and non bonding) 18 Molecular Geometry Effect of lone pairs ▪ Molecular geometry differs when one or more unshared (lone) pairs of electrons are present around the central atom. ▪ Electron-pair geometry is the same as that observed when only bonding pairs are involved, and bond angles are ordinarily a little smaller than ideal values listed. 19 Molecular Geometry Effect of lone pairs Lewis structure shows a central atom surrounded by two bonding pairs of electrons and one lone pair (AX2E). Electron pair geometry Molecular geometry Triangular planar Bent Example: SO2 20 Molecular Geometry Effect of lone pairs Lewis structure shows a central atom surrounded by three bonding pairs of electrons and one lone pair (AX3E). Electron pair geometry Molecular geometry Tetrahedron Triangular pyramid Example: NH3 21 Molecular Geometry Effect of lone pairs Lewis structure shows a central atom surrounded by two bonding pairs of electrons and two lone pairs (AX2E2). Electron pair geometry Molecular geometry Tetrahedron Bent Example: H2O 22 Molecular Geometry Effect of lone pairs Consider the NH3 and H2O molecules: Triangular pyramid 107° Bent 105° Why are the bond angles less than the ideal (109.5°)? 23 Molecular Geometry Effect of lone pairs Explanation: ⮚ An unshared pair is attracted by one nucleus, that of the atom to which it belongs. ⮚In contrast, a bonding pair is attracted by two nuclei, those of the two atoms it joins. ⮚Hence the electron cloud of an unshared pair is expected to spread out over a larger volume than that of a bonding pair. In NH3, this tends to force the bonding pairs closer to one another thus reducing the bond angle. In H2O, this effect is more pronounced due to the presence of 24 two unshared pairs. Molecular Geometry Effect of lone pairs Bonding electron pair Nonbonding electron pair Repulsion: Lone-pair vs lone-pair > lone-pair vs bonding > bonding-pair vs bonding-pair 25 Molecular Geometry Effect of lone pairs Summary Geometries of 2, 3 or 4 electron pairs around the central atom Species Structure Molecular Ideal Species Structure Molecular Ideal type geometry bond type geometry bond angles angles Tetrahedro AX2 Linear 180o AX4 109.5o n Triangular Triangular AX3 120o AX3E 109.5o * planar pyramid AX2E Bent 120o * AX2E2 Bent 109.5o * * In these species, the observed bond angle is less than the ideal value. Molecular Geometry Effect of lone pairs Lewis structure shows a central atom surrounded by four bonding pairs of electrons and 1 lone pair (AX4E). Electron pair geometry Molecular geometry Triangular bipyramid See-saw Example: SF4 27 Molecular Geometry Effect of lone pairs Lewis structure shows a central atom surrounded by three bonding pairs of electrons and two lone pairs (AX3E2). Electron pair geometry Molecular geometry Triangular bipyramid T-Shape Example: ClF3 28 Molecular Geometry Effect of lone pairs Lewis structure shows a central atom surrounded by two bonding pairs of electrons and three lone pairs (AX2E3). Electron pair geometry Molecular geometry Triangular bipyramid Linear Example: XeF2 29 Molecular Geometry Effect of lone pairs Lewis structure shows a central atom surrounded by five bonding pairs of electrons and one lone pair (AX5E). Electron pair geometry Molecular geometry Octahedron Square pyramid Example: ClF5 30 Molecular Geometry Effect of lone pairs Lewis structure shows a central atom surrounded by four bonding pairs of electrons and two lone pairs (AX4E2). Electron pair geometry Molecular geometry Octahedron Square planar Example: XeF4 31 Molecular Geometry Effect of lone pairs Summary 5 electron pairs 6 electron pairs Species Structure Molecular Bond Species Structure Molecular Bond type geometry angles type geometry angles 90o, Triangular 90o, AX5 120o, AX6 Octahedron bipyramid 180o 180o 90o, AX4E See-saw 120o, Square 90o, AX5E 180o pyramid 180o 90o, AX3E2 T-shape 180o Square 90o, AX4E2 planar 180o AX2E3 Linear 180o 32 Molecular Geometry Predict the molecular geometry of: NH4+, GeF2, PF3 it is tetrahedral with this is an AX4 species bond angle of 109.5° Lewis structure it is bent shape with bond this is an AX2E species angle less than 120° Lewis structure it is triangular pyramid with this is an AX3E species bond angle less than 109.5° Lewis structure 33 Molecular Geometry Predict the geometry of: ClO3- and NO3– ions, which have the below Lewis structures: a) The central atom, chlorine, is bonded to 3 oxygen atoms, it has one unshared pair. The ClO3- ion is of the type AX3E. This a triangular pyramid. b) The central atom, nitrogen, is bonded to 3 oxygen atoms, it has no unshared pairs. The NO3– ion is of the type AX3. It has a triangular planar structure. 34 Polarity of molecules Polar molecule Non-polar molecule Is a molecule which has a positive Is a molecule which has no charge and negative poles (partial positive separation (no positive and negative & partial negative charges). poles). i.e. has a net dipole. i.e. has no net dipole. If the molecule consists of only two atoms (diatomic molecule) bonded together, then by determining bond polarity, we can predict if the molecule is polar or non-polar. If the molecule consists of more than two atoms bonded together, not only bond polarity but also molecular geometry determines the polarity of the molecule. 35 Polarity of molecules Examples Diatomic molecules Non-polar Polar if there is no difference in if there is a significant electronegativity. difference in electronegativity. Molecules consisting of more than two atoms Non-polar Polar if there is symmetrical distribution of if there is an asymmetrical distribution electrons, which leads to a molecule of electrons, the molecule contains a with no positive or negative poles. positive & a negative pole (dipole). 36 Polarity of molecules The orientation of polar molecules in an electric field Molecules become oriented so that their negative ends are directed towards the positively charged plate and their positive ends towards the negatively charged plate. Polarity of molecules Consider the following molecules: All bonds in these molecules are polar as shown by the symbol: in which the arrow points to the more negative end of the bond and the + indicates the more positive end. BeF2 H2O CCl4 CHCl3 38 Polarity of molecules Example: BeF2 F Be F There are two polar Be—F bonds. In both bonds, the electron density is concentrated around the more negative fluorine atom. However, because BeF2 molecule is linear, the two dipoles are in opposite directions and cancel one another. Therefore, BeF2 has no net dipole moment and hence is nonpolar. 39 Polarity of molecules Example: H2O In the bent H2O molecule, the two H→O dipoles do not cancel each other. Instead, they add to give the H2O molecule a net dipole. The center of negative charge is located at the O atom, this is the negative pole of the molecule. The center of positive charge is located midway between the two H atoms. Therefore, H2O is a polar molecule. 40 Polarity of molecules Example: CCl4 There are four polar C-Cl bonds. However, because the four bonds are arranged symmetrically (tetrahedral) around the central carbon atom, their dipole moments cancel out each other. Therefore, CCl4 has no net dipole moment and hence is nonpolar. 41 Polarity of molecules Example: CHCl3 If one of the Cl atoms in CCl4 is replaced by a hydrogen atom, the situation changes. In the CHCl3 molecule, the H→C dipole does not cancel with the three C→Cl dipoles. Therefore, CHCl3 has a net dipole moment and hence is polar. 42  AX2 AX4 AX5 AX6 Linear, non-polar Triangular bipyramid, non-polar Octahedron, AX3 Tetrahedron, non-polar non-polar AX4E AX5E AX3E Triangular planar, See-saw, polar non-polar Square pyramid, Triangular pyramid, polar AX3E2 polar AX4E2 AX2E AX2E2 T-shaped, polar AX2E3 Square planar, Bent, polar Bent, polar Linear, non-polar Non-polar 43 Note: The non-polar molecules mentioned in the previous slide are non-polar only if the terminal atoms are identical. If the terminal atoms are different, the molecule becomes asymmetric and polar. 44  Geometry Project (8%) Each group of 4 students will construct a ball and stick model for a certain pharmaceutical compound. You will present the models in your last tutorial. You will also prepare a presentation about the drug assigned to your group, the presentation should cover the following points: 1- Chemical name. 2-Molecular formula and molar mass 3-Class of pharmaceuticals 4- Uses or indications 5- Route of administration 6-The hybridization of each central atom. 7- The geometry around each central atom 8- The polarity of the molecule 9- Formal charge of each atom. 45  Geometry Project (8%) These are the criteria upon which we can judge the presentation: 1- Not reading from the slides nor from note cards. 2- Good understanding of the aspects of the presented topic. 3- Capability to transfer the knowledge in a simple manner to audience. 4- Quality of the slides (not too much text, no plagiarism, good referencing, good organization). 5- Commitment to the time limit of the presentation (10 min.) 46 References Chemistry, 10th ed., Chang, Chapter 10. 47  General and Inorganic Chemistry (PHCM101) Lecture 8 Covalent bonding: Orbitals hybridization and intermolecular forces Dr. Nesrine El Gohary  COMPETANCIES 1-1-1 Identify different types of intermolecular forces. 1-1-2 Demonstrate understanding of the valence bond theory. 1-1-3 Demonstrate understanding of hybridization and hybrid orbitals. 2-2-1 Apply different types of orbital hybridization. 2-2-2 Determine the effect of multiple bonds. 2 Introduction to Intermolecular Forces Intermolecular forces are attraction forces between the molecules. Intramolecular forces are the chemical bonds holding the atoms together within a molecule. Intermolecular forces are much weaker than intramolecular forces. Intermolecular forces are responsible for many of the physical properties exhibited by substances. 3 Introduction to Intermolecular Forces Types Van der Ion-dipole Hydrogen Ionic Waals forces bonding bonding forces 4 Introduction to Intermolecular Forces Dipole-dipole forces Dipole- I- Van der Waals Ion-induced induced forces dipole forces dipole forces London dispersion forces 5 Introduction to Intermolecular Forces I- Van der Waals: 1- Dipole-dipole forces: ⮚ Present in polar molecules which are molecules that have a net permanent dipole (+ve end and -ve end). ⮚ Molecules with dipole moments can attract each other electrostatically by lining up so that the +ve and -ve ends are close to each other. This is called dipole-dipole attraction. ⮚ Dipole forces are 1% as strong as covalent or ionic bonds and they become weaker especially in the gas phase where the distance between dipoles increases. 6 Introduction to Intermolecular Forces I- Van der Waals: 2- London dispersion forces (induced fluctuating dipole): ⮚ Non-polar molecules and noble gases exhibit temporary dipoles, where it happens that for a very short time period the electrons are not distributed symmetrically around the molecule (the electron cloud is more dense on one side of the molecule in comparison to the other) leading to temporary δ+ and δ- charges (Instantaneous dipole). ⮚ This instantaneous dipole distorts the electron distribution in neighboring molecules thus inducing dipoles in these neighboring molecules (Induced dipole). ⮚ This leads to interatomic or intermolecular attraction that is relatively weak and short lived. 7 Introduction to Intermolecular Forces I- Van der Waals: 2- London dispersion forces (induced fluctuating dipole): H‫ــ‬H H‫ــ‬H No polarization Molecule A Molecule B δ- δ+ H‫ــ‬H H‫ــ‬H Instantaneous dipole on molecule A δ- δ+ δ- δ+ H‫ــ‬H H‫ــ‬H Induced dipole on molecule B 8 Introduction to Intermolecular Forces I- Van der Waals: 3- Ion-induced dipole interactions Are attractive forces that arise between ions and temporary dipoles induced in atoms or molecules. 9 Introduction to Intermolecular Forces I- Van der Waals: 4- Dipole-induced dipole interactions Are attractive forces that arise between dipoles and temporary dipoles induced in atoms or molecules. 10 Introduction to Intermolecular Forces II- Hydrogen Bonding: ⮚ It is a STRONG dipole-dipole attraction force. ⮚ Present among polar molecules in which Hydrogen is bound to a highly electronegative atom, such as Nitrogen, Oxygen or Fluorine. ⮚ Two factors account for the strengths of these interactions: 1. The great difference in the electronegativity between H and (N, O or F) leading to the great polarity of the molecule. 2. The close approach of the dipoles allowed by the small size of the H atom. 11 Introduction to Intermolecular Forces II- Hydrogen Bonding: 12 Introduction to Intermolecular Forces III- Ion-dipole forces: ⮚Are forces which attract an ion and a polar molecule to each other. ⮚Strength of interaction depends on the charge and size of the ion and on the magnitude of the dipole moment and size of the molecule. Interaction of a water molecule with a Na+ ion and a Mg2+ ion 13 Introduction to Intermolecular Forces IV- Ionic bonding: ⮚It is always found in compounds that are composed of both metallic and nonmetallic elements (example: NaCl). ⮚Each ion is attracted electrostatically to several neighboring ions of opposite charges. ⮚These attraction forces are called ionic bonds or interionic attractions. ⮚The interionic attractions increase as the charge on the ions increase and as the ionic radii decrease. Introduction to Intermolecular Forces IV- Ionic bonding: In solid ionic compounds, the ions are fixed in the crystalline lattice, therefore will not conduct electricity. However, when the ionic solid is dissolved in water the ions are free to move, so solutions of ionic compounds act as good conductors of electricity. Introduction to Intermolecular Forces Comparison of Strengths of Intermolecular Forces Understanding Chemical Bonding Two theories based on quantum mechanics explain the nature of the bonds formed between atoms. Valence Bond Theory Molecular orbital Theory Uses atomic orbitals to Uses molecular orbitals explain bonding to explain bonding 17 Understanding Chemical Bonding Valence Bond Theory According to this theory, a bond forms between two atoms when the following conditions are met: 1. An orbital on one atom comes to occupy a portion of the same region of space as an orbital on the other atom. This is called overlapping of atomic orbitals. 2. The total number of electrons in both orbitals is no more than two. 3. Overlap results in energy release, therefore the energy of the system reaches a minimum and is most stable (bond formation: exothermic reaction!) 18 Understanding Chemical Bonding Valence Bond Theory Let’s have a look at C atom: ↑ ↑ ↑↓ The electronic configuration of C is: 1s2 2s2 2p 2s 2p2 You might expect C to bind only with 2 H atoms and form CH2, since it has 2 single electrons in two of the p orbitals. Promotion But this doesn’t happen, C is known to bind ↑ ↑ with 4 H and forms CH4 (methane). ↑↓ 2p So we might explain this as follows: 2s An electron from the 2s got excited (promoted) and entered the empty p orbital. ↑ ↑ ↑ ↑ Now we have 4 single electrons in 4 orbitals 2p 2s ready for overlapping with the 4 H (1s). 19 Understanding Chemical Bonding Valence Bond Theory But this means that one bond would be ↑ ↑ ↑ formed from the overlap of the C (2s) orbital ↑ 2p and the H (1s) orbital, and three bonds 2s would be formed from the overlap of the C p orbitals with the 1s of the other three H atoms. This is not possible since it is known that the four C-H bonds in methane are identical. This means that the carbon orbitals involved in bonding are also equivalent. 20 Understanding Chemical Bonding Valence Bond Theory For this reason the valence bond theory assumes that the four valence orbitals of the carbon atom combine during the bonding process to form four new but equivalent hybrid orbitals, and these new hybrid orbitals are the ones involved in bonding. Hybridization is the mixing of native atomic orbitals to form special identical orbitals for bonding. The four new orbitals are called sp3 orbitals because they are formed from one 2s and three 2p orbitals. That is sp3 hybridization. The four orbitals are identical in shape, each one having a large lobe and a small lobe. 21 Zonkey is a real life example of hybridization! 22  Valence Bond Theory Sp3 Hybridization of Carbon in CH4 ↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑ 2p sp3 sp3 ↑ C-H bonds 2s ↑↓ ↑↓ ↑↓ 1s 1s 1s Promotion Hybridization Bond formation 23  Valence Bond Theory Sp3 Hybridization (CH4) sp3 ⮚ The four hybrid orbitals are identical in shape, each one having a large lobe and a small lobe. ⮚ Four hybrid orbitals are directed toward the four corners of a regular tetrahedral. ⮚ CH4 has a tetrahedral shape, and all the HCH angles are 109.5°. 24  Valence Bond Theory Sp3 Hybridization (CH4) Shape and arrangement of the hybrid orbitals 25  Valence Bond Theory Different types of hybridization sp Hybridization sp2 Hybridization sp3 Hybridization s p s p s p sp sp sp2 sp2 sp2 sp3 sp3 sp3 sp3 dsp3 Hybridization d2sp3 Hybridization s p d s p d dsp3 dsp3 dsp3 dsp3 dsp3 d2sp3 d2sp3 d2sp3 d2sp3 d2sp3 d2sp3 26  Valence Bond Theory How to determine the type of hybridization? ⮚ Count the number of effective electron pairs. What is considered as an effective electron pair? ⮚ Lone pair ⮚ Single bond ⮚ Double bond ⮚ Triple bond Each one of the above is counted as 1 effective electron pair. 27  Valence Bond Theory How to determine the type of hybridization? ⮚ The relationship between the number of effective pairs and their hybrid orbitals are given in this table: No of effective Type of Arrangement of electron pairs hybridization hybrid orbitals 2 sp Linear 3 sp2 Triangular planar 4 sp3 Tetrahedral 5 dsp3 Triangular bipyramid 6 d2sp3 Octahedral 28  Valence Bond Theory How to determine the type of hybridization? ⮚ The relationship between the number of effective pairs and their hybrid orbitals are given in this table: No of effective Arrangement of No of effective Arrangement of electron pairs hybrid orbitals electron pairs hybrid orbitals 2 5 3 6 4 29  Valence Bond Theory Sp3 Hybridization of Oxygen in H2O ↑↓ ↑ ↑ ↑↓ ↑↓ ↑ ↑ ↑↓ ↑↓ ↑ ↑ 2p sp3 sp3 Lone pairs O-H bonds ↑↓ 2s ↑↓ ↑↓ ↑↓ 1s 1s 1s Ground state Hybridization Bond formation 30  Valence Bond Theory Sp3 Hybridization of Oxygen in H2O ⮚ Two of the four hybrid orbitals of oxygen form covalent O-H bonds. The other two hybrid orbitals accommodate the lone pair on oxygen. ⮚ Repulsion between the lone-pair electrons and bonding-electron pairs decreases the HOH bond angles from 109.5° to 104.5 ° ⮚ Hybridization and VSEPR model are RELATED! 31  Valence Bond Theory Sp2 Hybridization of Boron in BF3 ↑ ↑ ↑ ↑ ↑ ↑ 2p 2p p sp2 ↑↓ ↑ Empty p orbital 2s 2s Ground State: Promotion: Hybridization: Ground state electron Energy promote an Since all B-F bonds are configuration can form electron from the 2s to an identical, therefore only one bond! empty 2p orbital, thus can sp2 hybridization must form three bonds! have occurred! 32  Valence Bond Theory Sp2 Hybridization of Boron in BF3 ⮚ Three hybrid orbitals are directed toward the 3 corners of an equatorial, planar triangle. ⮚ BF3 has a triangular planar shape, and all the FBF angles are 120°. 33  Valence Bond Theory Sp2 Hybridization of Carbon in Ethylene (C2H4) ⮚ Each C is surrounded by 3 effective electron pairs in Lewis structure, so we conclude that the hybridization must be sp2. ↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑ 2p 2p p sp2 ↑↓ ↑ Unhybridized p orbital 2s 2s Promotion: Hybridization: Ground State an electron from the 2s is Hybridization involves configuration of promoted to the empty 2p mixing of the S orbital valence electrons orbital and two p orbitals to give 3 identical sp2 orbitals 34  Valence Bond Theory Sp2 Hybridization of Carbon in Ethylene (C2H4) ⮚ In forming the sp2 orbitals, one 2p orbital on carbon has not been used. This remaining p orbital is oriented perpendicular to the plane sp2 orbitals. ⮚ The arrangement of the hybrid orbitals is triangular planar arrangement. 35  Valence Bond Theory Sp2 Hybridization of Carbon in Ethylene (C2H4) ⮚ The three sp2 orbitals on each carbon atom can be used to share electrons. In each of these bonds, the electron pair is shared in an area centered on a line running between the atoms. This type of bonds is called sigma(σ) bonds. In the σ bond, the electron pair occupies the space between the two atoms. ⮚ What about the second bond between the C atoms???? Remember each C atom has an unhybridized 2p orbital. Thus the second bond between the C atoms can be formed using the 2p orbital perpendicular to the sp2 hybrid orbitals on each carbon atom. 36  Valence Bond Theory Sp2 Hybridization of Carbon in Ethylene (C2H4) ⮚ The 2p orbitals are parallel to each other and undergo a side to side overlap (parallel overlap). ⮚ This results in sharing an electron pair in the space above and below the σ bond forming a pi (π) bond. 37  Valence Bond Theory Sp2 Hybridization of Carbon in Ethylene (C2H4) * Note that, σ bonds are formed from orbitals whose lobes point towards each other, but π bonds result from parallel orbitals. 38  Types of Covalent bonds Sigma Bond (σ) Pi Bond (π) ⮚ End-to-end (head to head) ⮚ Sideways (parallel) orbital orbital overlap. overlap. ⮚ Electron density concentrated ⮚ Electron density concentrated between the nuclei of the above and below the plane of bonding atoms. the nuclei of the bonding atoms. There is Stronger overlapping in sigma bonds, as a result sigma bonds are stronger than pi bonds. 39  Valence Bond Theory Sp Hybridization of Be in BeH2 ⮚ Number of effective electron pairs around Be is 2. Hybridization must be sp. ↑ ↑ ↑ 2p 2p p sp Empty p orbitals ↑↓ ↑ 2s 2s Promotion: Hybridization: Ground State energy promoted an Mixing of 2s and 2p electron electron from the 2s to orbitals gives two configuration cannot the 2p orbital, thus can identical sp hybrid form any bonds! form two bonds! orbitals. 40  Valence Bond Theory Sp Hybridization of Boron in BeH2 Be atom: BeH2 molecule: 1s atomic.... 1s atomic orbital of H orbital of H Overlap region 41  Valence Bond Theory Sp Hybridization in CO2 ⮚ Number of effective electron pairs around C is 2 so hybridization of C is sp. ⮚ Number of effective electron pairs around O is 3 so hybridization of O is sp2. Carbon Oxygen 2s 2p 2s 2p Ground State ↑↓ ↑ ↑ ↑↓ ↑↓ ↑ ↑ Ground State Promotion ↑ ↑ ↑ ↑ Hybridization ↑↓ ↑↓ ↑ ↑ Hybridization sp2 p ↑ ↑ ↑ ↑ sp p 42  Valence Bond Theory Sp Hybridization in CO2 The hybrid orbitals in CO2 molecule can be represented as: ⮚ Note the two 2p orbitals remain unchanged on the sp hybridized carbon. These are used to form the two π bonds with the oxygen atom. 43  Valence Bond Theory Sp Hybridization in CO2 Complete picture of CO2 molecule ⮚ Note the two 2p orbitals remain unchanged on the sp hybridized carbon. These are used to form the two π bonds with the oxygen atom. 44  Valence Bond Theory Sp Hybridization in N2 ⮚ Number of effective electron pairs around each N is 2 so hybridization of each N is sp. 2s 2p Ground State ↑↓ ↑ ↑ ↑ Hybridization ↑↓ ↑ ↑ ↑ unhybridized p orbitals are used to sp p form the pi bonds 45  Valence Bond Theory Some important terms The bond order is the number of electron pairs being shared by a given pair of atoms. ⮚ A single bond consists of one bonding pair and has a bond order of 1. ⮚ A double bond consists of two bonding pairs and has a bond order of 2. ⮚ A triple bond consists of three bonding pairs and has a bond order of 3. The bond energy (BE) or bond enthalpy is the energy needed to overcome the attraction between the nuclei and the shared electrons , in other words it is the energy needed to break the bond between the atoms in a molecule. The stronger the bond the higher the bond energy. 46  Valence Bond Theory dsp3 Hybridization Consider SF4 and PCl5, a set of five effective pairs around the S or P atom is present. In general, a set of five effective pairs around a given atom always requires a triangular bipyramid arrangement, which in turn requires dsp3 hybridization of that atom d2sp3 Hybridization Consider SF6 and XeF4 (2 lone pairs around Xe). In general, six electron pairs around an atom are always arranged octahedrally and require d2sp3 hybridization of the atom. 47  Valence Bond Theory Hybridization with d-orbitals, 3rd row: Extension of Octet rule (SF4) d p Promotion then dsp3 s ↑↓ ↑ ↑ Hybridization ↑↓ ↑ ↑ ↑ ↑ ↑↓ Trigonal bi- pyramidal Bonding F F F F ↑↓ ↑ ↑ ↑ ↑ dsp3 48  Valence Bond Theory Hybridization with d-orbitals, 3rd row: Extension of Octet rule (SF6) d p Promotion then d2sp3 s ↑↓ ↑ ↑ Hybridization ↑ ↑ ↑ ↑ ↑ ↑ ↑↓ Octahedral Bonding F F F F F F ↑ ↑ ↑ ↑ ↑ ↑ d2sp3 49  Practice Exercises For each of the following molecules or ions, predict the hybridization of each atom: a. BF4- Draw Lewis structure first. B is surrounded by 4 effective electron pairs so the hybridization is sp3. The hybrid orbitals have a tetrahedral arrangement. Each fluorine atom has four electron pairs so it is sp3 hybridized. 50  Practice Exercises For each of the following molecules or ions, predict the hybridization of each atom: b. XeF2 The Lewis structure shows five electron pairs on the xenon atom, which requires a triangular bipyramidal arrangement. To accommodate five pairs at the vertices of a trigonal bipyramid requires that the Xenon atom adopt a set of five dsp3 orbitals. Each fluorine atom has four electron pairs and can be assumed to be sp3 hybridized. The XeF2 molecule has a linear arrangement of atoms (molecular geometry). *Note: It is not a must that all the atoms have the same type of hybridization 51  Practice Exercises For each of the following molecules or ions, predict the hybridization of each atom: c. 52 References Chemistry, 10th ed., Chang, Chapter 10 and 11. 53 General and Inorganic Chemistry PHCM101 Lecture 9 Covalent bonding: Molecular orbital theory Dr. Nesrine El Gohary  COMPETANCIES 1-1-1 Identify the different types of molecular orbitals. 1-1-2 Articulate knowledge in designing molecular orbital diagrams for diatomic molecules. 2-2-1 Determine the bond order for diatomic molecules. 2-2-2 Determine the magnetic properties of diatomic molecules. 2  Molecular orbital theory The theory describes covalent bonds in terms of molecular orbitals. Molecular orbitals result from interaction of the atomic orbitals of the bonding atoms. Molecular orbitals are orbitals that belong to the whole molecule rather than to a single atom. The electrons are found in the molecular orbitals.  Molecular orbital theory Atomic orbital Molecular orbital Describes the volume of space around the nucleus of an atom Describes the volume of space around where an electron is likely to be a molecule where an electron is likely found to be found Atomic orbitals are associated with Molecular orbitals are associated with only one atom. the whole molecule Atomic orbitals have specific sizes, Molecular orbitals have specific sizes, shapes and energies shapes and energies An Atomic orbital carries a maximum A Molecular orbital carries a maximum of two electrons with opposite spin of two electrons with opposite spin  Molecular orbital theory Molecular orbital theory helped in the explanation of some properties of molecules that could not be explained by Lewis structure and the Valence bond theory. Example: Oxygen molecule O2 This is Lewis structure for O2 According to this structure all the electrons are paired, so O2 is a diamagnetic molecule. The reality is O2 is a paramagnetic molecule. This means that O2 should have unpaired electrons to be paramagnetic. The presence of unpaired electrons and its paramagnetism could be explained by the molecular orbital theory. https://www.youtube.com/watch?v=Lt4P6ctf06Q  Molecular orbital theory How are molecular orbitals formed? Atomic orbitals of the bonding atoms combine together to form molecular orbitals. In this theory, the wave nature of electrons explains the different types of molecular orbitals obtained. Remember! Electrons in orbitals have wave properties, they behave like waves. It is known that waves can interact either constructively or destructively. If waves (atomic orbitals) interact constructively, the resulting molecular orbital is lower in energy and more stable than the two atomic orbitals. It is thus called a bonding molecular orbital. If waves (atomic orbitals) interact destructively, the resulting molecular orbital is higher in energy and is less stable than the two atomic orbitals. It is thus called an antibonding molecular orbital. constructive interference; destructive interference  Molecular orbital theory [H2 molecule] Each hydrogen atom has its 1s atomic orbital. As the hydrogen atoms approach each other the two 1s atomic orbitals interact to form two molecular orbitals: Bonding molecular orbital which is more Antibonding molecular orbital which is less stable and of lower energy than the stable and of higher energy than the isolated isolated 1s atomic orbitals. 1s atomic orbitals. In the bonding molecular orbital, the In the antibonding molecular orbital, the electron density in this molecular orbital is electron density in this molecular orbital is concentrated in the region between the also concentrated along the internuclear axis two nuclei along the internuclear axis, so it but it is on either side of the nuclei. Between is called a sigma () bonding molecular the two nuclei we have a node orbital. In this case it is called  1s. ( region where there is no probability of finding electrons). This antibonding molecular orbital is called * 1s.  Molecular orbitals: s orbitals [H2 molecule] The bonding  1s molecular orbital is lower in energy than the antibonding * 1s, consequently it is filled with electrons before the antibonding * 1s. Any molecular orbital can be filled with a maximum of 2 electrons and they have to have opposite spin. Thus, the two electrons of the hydrogen molecule are present in the bonding  1s. This favors the bonding in hydrogen molecule since this molecular orbital is more stable and of lower energy than the isolated 1s atomic orbital of each hydrogen atom. So with the presence of the two electrons in this more stable molecular orbital, it favors bonding and the formation of the hydrogen molecule. Imagine if the two electrons were present in the antibonding * 1s which is less stable and of higher energy than the isolated 1s atomic orbitals. This condition will not favor bonding at all (ANTIBONDING) and it will be more stable for the hydrogen to remain as separated atoms and not combine into a molecule. Molecular orbital energy level diagram of H2 molecule We conclude that presence of electrons in bonding molecular orbitals favors bonding, while the presence of electrons in antibonding molecular orbitals is against bonding.  Molecular orbitals: p orbitals What are the molecular orbitals formed from interaction of p atomic orbitals?? Remember that there are three p atomic orbitals, that are oriented along, x, y and z axes Interaction of two p orbitals, one from each atom results in the formation of two molecular orbitals Bonding and antibonding molecular orbitals. end-to-end overlap Sigma () molecular orbitals Side by side overlap Pi (π) molecular orbitals Electron density is concentrated above and below a line joining the two nuclei. Note that: The two bonding π molecular Note that: The two antibonding π* molecular orbitals are equal in energy (degenerate). orbitals are equal in energy (degenerate).  Molecular orbital energy level diagram Rules Governing Molecular Electron Configuration and Stability: 1. The number of molecular orbitals formed is always equal to the number of atomic orbitals combined. 2. The more stable the bonding molecular orbital, the less stable the corresponding antibonding molecular orbital. 3. The filling of molecular orbitals proceeds from low to high energies. In a stable molecule, the number of electrons in bonding molecular orbitals is always greater than that in antibonding molecular orbitals. 4. Like an atomic orbital, each molecular orbital can accommodate up to two electrons with opposite spins in accordance with the Pauli exclusion principle. 5. When electrons are added to molecular orbitals of the same energy, we follow Hund’s rule; that is, electrons enter singly with parallel spin before we start pairing. 6. The number of electrons in the molecular orbitals is equal to the sum of all the electrons on the bonding atoms.  Molecular orbital energy level diagram Rules Governing Molecular Electron Configuration and Stability: 1. The number of molecular orbitals formed is always equal to the number of atomic orbitals combined. 2. The more stable the bonding molecular orbital, the less stable the corresponding antibonding molecular orbital. 3. The filling of molecular orbitals proceeds from low to high energies. In a stable molecule, the number of electrons in bonding molecular orbitals is always greater than that in antibonding molecular orbitals. 4. Like an atomic orbital, each molecular orbital can accommodate up to two electrons with opposite spins in accordance with the Pauli exclusion principle. 5. When electrons are added to molecular orbitals of the same energy, we follow Hund’s rule; that is, electrons enter singly with parallel spin before we start pairing. 6. The number of electrons in the molecular orbitals is equal to the sum of all the electrons on the bonding atoms.  Molecular orbital energy level diagram Example Draw the molecular orbital energy level diagram for the oxygen molecule. Calculate the total electrons present in the molecule. Oxygen atom has 8 electrons, so we σ* 2px have a total of 16 electrons. Draw the molecular orbitals п* 2pz represented as boxes or lines , п* 2py starting from the bottom which п 2pz п 2py represents the molecular orbital of lowest energy. σ 2px Arrange the rest of the molecular orbitals moving upwards according to σ* 2s increasing energy. If two orbitals have the same energy σ 2s such as π 2pz and π 2py draw them at the same level. σ* 1s Start putting electrons in the bottom (orbital of lowest energy) and σ 1s continue filling moving upwards. If all electrons are paired so diamagnetic O2 molecule has unpaired electrons so it is molecule. PARAMAGNETIC. If there is unpaired electrons so paramagnetic  Molecular orbital energy level diagram Example Draw the molecular orbital energy level diagram for the oxygen molecule. Calculate the total electrons present in the molecule. Oxygen atom has 8 electrons, so we σ* 2px have a total of 16 electrons. Draw the molecular orbitals п* 2pz represented as boxes or lines , п* 2py starting from the bottom which п 2pz п 2py represents the molecular orbital of lowest energy. σ 2px Arrange the rest of the molecular orbitals moving upwards according to σ* 2s increasing energy. If two orbitals have the same energy σ 2s such as π 2pz and π 2py draw them at the same level. σ* 1s Start putting electrons in the bottom (orbital of lowest energy) and σ 1s continue filling moving upwards. If all electrons are paired so diamagnetic O2 molecule has unpaired electrons so it is molecule. PARAMAGNETIC. If there is unpaired electrons so paramagnetic  Molecular orbital energy level diagram Example Draw the molecular orbital energy level diagram for the oxygen molecule. Calculate the total electrons present in the molecule. Oxygen atom has 8 electrons, so we σ* 2px have a total of 16 electrons. Draw the molecular orbitals п* 2pz represented as boxes or lines , п* 2py starting from the bottom which п 2pz п 2py represents the molecular orbital of lowest energy. σ 2px Arrange the rest of the molecular orbitals moving upwards according to σ* 2s increasing energy. If two orbitals have the same energy σ 2s such as π 2pz and π 2py draw them at the same level. σ* 1s Start putting electrons in the bottom (orbital of lowest energy) and σ 1s continue filling moving upwards. If all electrons are paired so diamagnetic O2 molecule has unpaired electrons so it is molecule. PARAMAGNETIC. If there is unpaired electrons so paramagnetic  Bond order in molecular orbital theory It is the difference between the number of bonding electrons and the number of anti bonding electrons divided by two. Bond order indicates the approximate strength of a bond, the bigger the bond order, the stronger is the bond. If Bond order > zero Bond order =zero the molecule is the molecule is unstable and relatively stable and cannot exist. exists. Bond order containing fractions are possible.  Molecular orbital energy level diagram Example Draw the molecular orbital energy level diagram for the oxygen molecule. Calculate the total electrons present in the molecule. Oxygen atom has 8 electrons, so we σ* 2px have a total of 16 electrons. Draw the molecular orbitals п* 2pz represented as boxes or lines , п* 2py starting from the bottom which п 2pz п 2py represents the molecular orbital of lowest energy. σ 2px Arrange the rest of the molecular orbitals moving upwards according to σ* 2s increasing energy. If two orbitals have the same energy σ 2s such as π 2pz and π 2py draw them at the same level. σ* 1s Start putting electrons in the bottom (orbital of lowest energy) and σ 1s continue filling moving upwards. If all electrons are paired so diamagnetic O2 molecule has unpaired electrons so it is molecule. PARAMAGNETIC. If there is unpaired electrons so paramagnetic  Bond order Compare the bond order of H2, H2+ and H2- H2 H2+ H2-   1s 1s 1s 1s AO of H AO of H AO of AO of H  H-  MO of H2- bond order = 1/2(1-0) = 1/2 bond order = 1 H2+ does exist 1/2(2-1) B.O = (2 − 0) = 1 = 1/2 2 H2- does exist  Bond order Compare the bond order of He2, He2+ He2+ He2 1 1 B.O = (2 − 1) = 1/2 B.O = (2 − 2) = 0 2 2 He2+ Exists He2 Cannot Exist  Homonuclear Diatomic Molecules of Second-Period Elements They are diatomic molecules containing atoms of the same elements. E.g.Li2 Electronic configuration of Li atoms: 1s2 2s1 Electron configuration of molecular orbitals in Li2: (1s)2 (*1s)2 (2s)2. Bond Order = 1 Molecule is diamaganetic  Homonuclear Diatomic Molecules of Second-Period p-block elements 2px, 2py and 2pz atomic orbitals 2px and 2px M.O 2py ,2pz , *2py and * 2pz (overlap of 2px orbitals) (overlap of 2py & 2pz orbitals) Normally, overlap of the two p orbitals is greater in a  molecular orbital than in a  molecular orbital, so we would expect the former to be lower in energy. However, the order 2p and 2p molecular orbitals is flipped for light molecules such as B2, C2, N2.  Homonuclear Diatomic Molecules of Second-Period p-block elements Light molecules with  half filled p-orbitals; e.g B2, C2, N2 Flipped Order General molecular orbital energy level diagram for the second-period homonuclear diatomic molecules Li2, Be2, B2, C2, and N2.  Homonuclear Diatomic Molecules of Second-Period p-block elements Molecules with > half filled p-orbitals; O2 and beyond Order of molecular orbitals is as expected;  2py  2pz i.e not flipped  2px  Practice Exercise Use the molecular orbital model to predict the bond order and magnetism of each of the following molecules: HOMO = Highest occupied molecular orbital. LUMO = Lowest unoccupied molecular orbital.  Practice Exercise For each of these species; O2, O2+ and O2- give: The molecular electronic configuration & the bond order Determine which has the strongest bond & predict the magnetism of each. No. of valence O2 = 12 O2+ = 11 O2- = 13 electrons: (6+6) (6+6-1) (6+6+1) The simplified MO diagrams: O2+ = Strongest Bond Bond Order ½ (8-4) = 2 ½ (8-3) = 2.5 ½(8-5) = 1.5 Magnetism Paramagnetic Paramagnetic Paramagnetic  Heteronuclear Diatomic Molecules of Second- Period Elements E.g. CN-1 NO+1 No. of 10 10 valence (4+5+1) (5+6-1) electrons= Bond Order ½ (8-2) = 3 ½ (8-2) = 3 Magnetism Diamagnetic Diamagnetic Molecular Orbital Diagram of CN-1 & NO+1  References Chemistry, 6th ed., Zumdahl, Chapter 9. https://socratic.org/questions/5a64204ab72cff7b83554c78 General and Inorganic Chemistry PHCM101 Lecture 10 Redox reactions Principles of Electrochemistry Dr. Nesrine El Gohary 1  COMPETANCIES 1-1-1 Identify oxidation-reduction (Redox) reaction. 1-1-2 Articulate knowledge in assigning oxidation numbers to different atoms and in balancing redox-reaction equations. 1-1-3 Describe the components of a galvanic cell. 1-1-4 Define cell potential and standard electrode potentials. 2-2-1 Explain how electricity can be generated from a galvanic cell. 2-2-2 Explain the dependence of cell potential on concentration. (Nernst equation). 4-3-1 Self learning (general applications). 2  REDOX REACTIONS Oxidation – Reduction Reactions (electron transfer reactions) Oxidation Reduction Losing electrons Gaining electrons Which electrons are lost or gained? Electrons of the outermost energy level ( Valence electrons) Oxidation and reduction have to happen simultaneously (together). The element that LOSES electrons gets oxidized and is called the REDUCING AGENT The element that GAINS electrons gets reduced and is called the OXIDIZING AGENT Any Redox Reaction consists of One half is the oxidation and the other half is reduction 3 2 Half Reactions.  REDOX REACTIONS Remember LEO says GER LEO: Losing Electrons (Oxidation) GER: Gaining Electrons (Reduction) 4  OXIDATION NUMBERS (ON) Oxidations number are numbers (+ve or -ve) assigned to atoms in compounds that help us keep track of electrons in redox reactions and follow the changes that take place. 2 2 2 2 ON = 0 ON = 0 ON = +1 ON = -1 Increase in oxidation number Decrease in oxidation number OXIDATION of Na REDUCTION of Cl Redox reactions occur when: A substance reacts with O2. A metal combines with a nonmetal. In general, whenever electrons are transferred. 5 When there is a change in an atom oxidation number.  ASSIGNING OXIDATION NUMBERS (ON) Rules for assigning an oxidation state (O.N.) General rules 1) For an atom in its elemental form (Na, O2, Cl2, etc): O.N. = 0 - 2- + 2) For monatomic ion (F , S , K , etc) O.N. = ion charge 3) The sum of the O.N. values for the atoms in a neutral compound equal zero. 4) The sum of the O.N. values for the atoms in a polyatomic ion equal the ion’s charge. Rules for specific atoms 1. For alkali metals: Na, K, Li (Gp I) O.N. = +1 in all compounds. 2. For alkaline earth metals: Mg, Ca , Ba, Sr (Gp II) O.N. = +2 in all compounds. 3. For hydrogen O.N. = +1 in combination with nonmetals O.N. = -1 in combination with metals (NaH, CaH2) 6  ASSIGNING OXIDATION NUMBERS (ON) Rules for assigning an oxidation state (O.N.) [continue) 4. For fluorine O.N. = -1 in all compounds 5. For oxygen O.N. = -2 in all other compounds O.N. = -1 in peroxides (H2O2) O.N. = -1/2 in superoxides (KO2) 6. For halogens (Cl, Br, I) O.N. = -1 in combination with metals and nonmetals (except O) Example: Assign oxidation numbers to all atoms in the following species: CO2 SF6 NO3− ZnCl2 KMnO4 +4 -2 +6 -1 +5 -2 +2 -1 +1 +7 -2 X 2 X 6 X 3 X 2 X 4 -4 -6 -6 -2 -8 7  TEST YOURSELF Which atoms in the following reaction are reduced and which are oxidized? Mg( s ) + 2 HCl( aq) → MgCl2 ( aq) + H 2 ( gas) +1 −1 +2 −1 Mg + 2 H Cl → Mg Cl + H 0 (s) 2 0 2 2 Hydrogen gained 2 e-, it is reduced Magnesium lost 2 e-, it is oxidized Half reaction Half reaction Mg is the Reducing agent HCl is the Oxidizing agent It donates its electrons to other atoms so It accepts electrons from other atoms so causes their reduction causes their oxidation 8  HOW to Balance a Redox Reaction Balance the following redox reaction that takes place in acidic media: Fe2+ + MnO4- → Fe3+ + Mn2+ 1- Split the reaction into its two half reactions. Fe2+ → Fe3+ MnO4- → Mn2+ 2- Determine which half is oxidation and which half is reduction. Fe2+ → Fe3+ MnO4- → Mn2+ There is an increase in There is a oxidation number of Fe so +7 -2 decrease in this half is oxidation X 4 oxidation number -8 of Mn so this half is reduction 3- Balance each half reaction separately. 4- Combine or add the two half-equations. 9  BALANCING REDOX HALF REACTIONS MnO4- → Mn2+ (acidic solution) 1. Balance Atoms of the element being oxidized or reduced. There is 1 atom of manganese on both sides→ no adjustment is needed. 2. Balance Oxidation number by adding electrons on the appropriate side. Note that: For reduction half reaction, electrons are gained so are added in reactant side. For oxidation half reaction, electrons are lost so are added in product side. Mn+7O4 - + 5e- → Mn2+ Mn is reduced from +7 to +2→ 5 electrons must be added to the left. 3. Balance Charge by adding H+ ions in acidic solution, OH- ions in basic solution MnO4- + 8H+ + 5e- → Mn2+ On the left side, MnO4- is 1 negative charge and the electrons are 5 negative charges, giving a total of 6 negative charges. While on the right side, there are 2 positive charges on Mn ion. To balance charges on both sides, 8 positive charges need to be added to the left to make +8-6=+2. 4. Balance Hydrogen by adding H2O. MnO4- + 8H+ + 5e →Mn2+ + 4H2O To balance the 8H+ ions on the left, add 4 water molecules to the right. 10 5. Make sure that Oxygen is balanced → you did right ☺  BALANCING REDOX HALF REACTIONS Balance the second half reaction Fe2+ → Fe3+ 1. Balance Atoms of the element being oxidized or reduced. There is 1 atom of iron on both sides→ no adjustment is needed. 2. Balance Oxidation number by adding electrons on the appropriate side. This is an oxidation half reaction, electrons are lost so are added in product side. Fe2+ → Fe3+ + e 3. Balance Charge by adding H+ ions in acidic solution, OH- ions in basic solution The charge are identical on both sides of the reaction (2+) so there is no need to add anything and this half reaction is balanced. 11  WRITING THE FULL REDOX REACTION 1. Split the equation into two half equations. 2. Balance each half-equation, independently. 3. Combine or Add the two half-equations, when combining or adding the two half reactions, electrons must be eliminated and the final equation should not contain any electrons so sometimes necessary multiplication by a factor to eliminate electrons is needed. Fe2+ + MnO4- → Fe3+ + Mn2+ Oxidation: Fe2+ → Fe3+ Reduction: MnO4- → Mn2+ The balanced half equations are Fe2+ → Fe3+ + e- MnO4-+ 8H+ + 5e- → Mn2+ + 4 H2O To eliminate electrons: multiply the oxidation half equation by 5. Then, sum up the 2 half reactions 5 Fe+ → 5Fe3+ + 5e- MnO4-+ 8H+ + 5e-→ Mn2+ + 4 H2O --------------------------------------------------------- 5Fe2+ + MnO4-+ 8H+ → 5Fe3+ + Mn2+ + 4H2O 12  TEST YOURSELF Balance the following redox reaction in basic solution Cr(OH)3 + ClO- → CrO42- + Cl- 1- Split the reaction into its two half reactions. ClO- → Cl- Cr(OH)3 → CrO42- 2- Determine which half is oxidation and which half is reduction. ClO- → Cl- Cr(OH)3 → CrO42- +1 -2 -1 +3 -1 +6 -2 Reduction X 3 X 4 -3 -8 Oxidation 3- Balance each half reaction separately. 13  TEST YOURSELF CONT. Cr(OH)3 → CrO42- 1. There is 1 chromium atom on both sides→ atoms are balanced. 2. The oxidation number of chromium increases from +3 to +6, add 3 electrons to the right. Cr+3(OH)3 →Cr+6O4 2- + 3e- 3. There is a charge of zero to the left, -5 to the right→ add 5 OH- ions to the left. Cr(OH)3 + 5 OH- → CrO42- + 3e- 4. There are 8 H atoms to the left, add 4 H2O molecules to the right. Cr(OH)3 + 5 OH- → CrO42- + 4H2O + 3e- 5. There are 8 oxygen atoms on both sides → the half reaction is balanced. 14  TEST YOURSELF CONT. ClO- → Cl- 1. There is 1 chlorine atom on both sides→ atoms are balanced. 2. The oxidation number of chlorine decreases