Medical Chemistry Lecture 2024 - Types of Chemical Bonds (PDF)

All your classes are cancelled this Wednesday (September 18th) from 4 pm. Schedule of seminars on Wednesday till 4 pm will not change. Since there will be no further chance to discuss the topic of week2 lecture in Medical Chemistry, we would like to offer the opportunity to make up for the missed me...

All your classes are cancelled this Wednesday (September 18th) from 4 pm. Schedule of seminars on Wednesday till 4 pm will not change. Since there will be no further chance to discuss the topic of week2 lecture in Medical Chemistry, we would like to offer the opportunity to make up for the missed medical chemistry seminars. Attendance is not compulsory at these seminars. Time and location of the make-up seminars: September 17th (Tuesday) at 4 p.m. LC1.05 September 17th (Tuesday) at 4 p.m. LC2.14 September 17th (Tuesday) at 6 p.m. LC1.05 September 17th (Tuesday) at 6 p.m. LC2.16 September 18th (Wednesday) at 8 a.m. LC1.05 September 18th (Wednesday) at 8 a.m. LC1.13 September 18th (Wednesday) at 10 a.m. LC1.13 September 18th (Wednesday) at 10 a.m. LC2.16 September 18th (Wednesday) at 2 p.m. LC1.05 September 23rd (Monday) at 8 a.m. Lecture Hall, Department of Obstetrics and Gynecology Week 2 Lecture 3-4 Medical Chemistry Lecture 2024 Types of chemical bonds. Covalent bonding. Intermolecular forces. Beata Lontay Department of Medical Chemistry UD Faculty of Medicine  Oral exam topics A3. Types of chemical bonds. Polar and nonpolar molecules. VSERP theory A4. Valence Bond theory of covalent bonds. Hybrid orbitals A5. Molecular Orbital theory of covalent bonding A6. Intermolecular forces (types and consequences). Properties of water  Bonding in Chemistry Central theme in chemistry: Why and How atoms attach together This will help us understand how to: 1. Predict the shapes of molecules. 2. Predict properties of substances. 3. Design and build molecules with particular sets of chemical and physical properties. NaCl C12H22O11  Chemical bonds  Atoms or ions are held together in molecules or compounds by chemical bonds.  The type and number of electrons in the outer electronic shells of atoms or ions are instrumental in how atoms react with each other to form stable chemical bonds.  Over the last 150 years scientists developed several theories to explain why and how elements combine with each other.  Formation of chemical bonds All chemical reactions involve breaking of some bonds and formation of new ones which yield new products with different properties.  Types of chemical bonds Type of bonds Interaction Bond energy Strength Example (kJ/mol) IONIC Cation-anion 400-4000 Strong NaCl(s) Primary COVALENT Electron pairs 150-1100 Strong F2(g), bonds CH4(g) METALIC Cation and valence 75-1000 Strong Na(s), electrons Mg(s) ION-DIPOLE Charge and permanent 40-600 Strong NaCl (aq) dipole HYDROGEN Polar hydrogen and 10-40 Medium H2O-H2O BONDING non-bonding electron H2O- pair CH3OH DIPOLE-DIPOLE Permanent dipoles 5-25 Medium HCl-HCl Secondary bonds ION-INDUCED Ion and the dipole 3-15 Weak Fe2+-O2 DIPOLE induced by that DIPOLE-INDUCED Permanent dipole and 2-10 Weak HCl-Cl2 DIPOLE the dipole induced by that LONDON FORCE Temporary dipoles 0.05-40 Weak Ar(g)-Ar(g) 1. Ionic bonding Ionic Bonds The ionic compound NaCl Metal to nonmetal. Metal loses electrons to form cation. Nonmetal gains electrons to form anion. The electronegativity between the metal and the nonmetal must be > than 2. Ionic bond results from + to − attraction. – Larger charge = stronger attraction. – Smaller ion = stronger attraction. Lewis theory allows us to predict the correct formulas of ionic compounds.  Electron Configurations of Ions Na: 1s2 2s2 2p6 3s1 -1 e- Na+:1s2 2s2 2p6 Cl: 1s2 2s2 2p6 3s2 3p5 +1 e- Cl-:1s2 2s2 2p6 3s2 3p6 Group 6a atom: [Noble gas] ns2 np4 +2 e- Group 6a ion2-: [Noble gas] ns2 np6 Group 7a atom: [Noble gas] ns2 np5 +1 e- Group 7a ion-: [Noble gas] ns2 np6 Electron Configurations of Ions Electron Configurations of Ions Atoms Ions Fe: [Ar] 4s2 3d6 - 2 e- Fe2+: [Ar] 3d6 Fe: [Ar] 4s2 3d6 - 3 e- Fe3+: [Ar] 3d5  Ionic Bonds and the Formation of Ionic Solids 1s2 2s2 2p6 3s1 1s2 2s2 2p6 3s2 3p5 Na + Cl Na+ + Cl- 1s2 2s2 2p6 1s2 2s2 2p6 3s2 3p6  Ionic Bonds and the Formation of Ionic Solids Born-Haber Cycle Electron affinity of chlorine First ionization of sodium Dissociation of Cl2 Formation of solid NaCl Sublimation of Na  Ionic Bonds and the Formation of Ionic Solids Born-Haber Cycle Step 1: Na(s) Na(g) +107.3 kJ/mol 1 Step 2: Cl2(g) Cl(g) +122 kJ/mol 2 Step 3: Na(g) Na+(g) + e- +495.8 kJ/mol Step 4: Cl(g) + e- Cl-(g) -348.6 kJ/mol Step 5: Na+(g) + Cl-(g) NaCl(s) -787 kJ/mol 1 Na(s) + Cl2(g) NaCl(s) -411 kJ/mol 2 Lewis structure of ionic compounds Ions that pack as spheres in a very regular pattern form crystalline substances Lattice Energies in Ionic Solids Lattice Energy (U): The amount of energy that must be supplied to break up an ionic solid into individual gaseous ions 2. Covalent Bonds and Molecular Structure Covalent Bonding in Molecules Covalent Bond: A bond that results from the sharing of electrons between atoms. © 2012 Pearson Education, Inc. Chapter 7/21 Strengths of Covalent Bonds A Comparison of Ionic and Covalent Bonds  Polar Covalent Bonds: Electronegativity Electronegativity: The ability of an atom in a molecule to attract the shared electrons in a covalent bond. © 2012 Pearson Education, Inc. Chapter 7/24 Polar Covalent Bonds: Electronegativity © 2012 Pearson Education, Inc. Chapter 7/25 Polar Covalent Bonds: Electronegativity NONPOLAR COVALENT BOND POLAR COVALENT BOND Bond polarity Bond polarity vs molecular polarity Lewis Bond Molecular Molecular Physical structure polarity shape polarity properties Bond polarity Molecular polarity  Bonding Theories and Models Lewis bond Theory Valence-Shell Electron-Pair Repulsion model Valence Bond Theory Molecular Orbital Theory Gilbert Newton Lewis  A. Lewis Bonding Theory Atoms ONLY come together to produce a more stable electron configuration. Atoms bond together by either transferring or sharing electrons. Many of atoms like to have 8 electrons in their outer shell. – Octet rule. – There are some exceptions to this rule—the key to remember is to try to get an electron configuration like a noble gas. Li and Be try to achieve the He electron arrangement.  Lewis Symbols of Atoms Uses symbol of element to represent nucleus and inner electrons. Uses dots around the symbol to represent valence electrons. – Puts one electron on each side first, then pair. Remember that elements in the same group have the same number of valence electrons; therefore, their Lewis dot symbols will look alike. Li Be B C N O: :F: :Ne:  Octet Rule Octet rule: Main-group elements tend to undergo reactions that leave them with eight outer-shell electrons. That is, main-group elements react so that they attain a noble-gas electron configuration with filled s and p sublevels in their valence electron shell.  Octet Rule Octet rule: Main-group elements tend to undergo reactions that leave them with eight outer-shell electrons. That is, main-group elements react so that they attain a noble-gas electron configuration with filled s and p sublevels in their valence electron shell. Metals tend to have low Ei and low Eea. They tend to lose one or more electrons. Nonmetals tend to have high Ei and high Eea. They tend to gain one or more electrons.  Electron-Dot Structures Electron-Dot Structure (Lewis Structure): Represents an atom’s valence electrons by dots and indicates by the placement of the dots the way the valence electrons are distributed in a molecule. Electron-Dot Structures Electron-Dot Structures of Polyatomic Molecules The Octet Rule  The VSEPR Model VSEPR: Valence-Shell Electron-Pair Repulsion model Electrons in bonds and in lone pairs can be thought of as “charge clouds” that repel one another and stay as far apart as possible, thus causing molecules to assume specific shapes. It is used to predict the geometries of molecules formed from nonmetals. Postulate: the structure around a given atom is determined principally by minimizing electron pair repulsion. The bonding and nonbonding pairs should be positioned as far apart as possible.  Molecular Shapes: The VSEPR Model Step 1 Write an electron-dot structure for the molecule and count the number of electron charge clouds surrounding the atom of interest. Step 2 Predict the geometric arrangement of charge clouds by assuming that the charge clouds are oriented in space as far away from one another as possible to minimize repulsions. The VSEPR Model: Two Charge Clouds LINEAR The VSEPR Model: Three Charge Clouds TRIGONAL PLANAR BENT The VSEPR Model: Four Charge Clouds The VSEPR Model: Four Charge Clouds TETRAHEDRAL TRIGONAL PYRAMIDAL BENT  The VSEPR Model: Five Charge Clouds TRIGONAL BIPYRAMIDAL The VSEPR Model: Five Charge Clouds TRIGONAL BIPYRAMIDAL SEESAW  The VSEPR Model: Five Charge Clouds T-SHAPED LINEAR The VSEPR Model: Six Charge Clouds  The VSEPR Model: Six Charge Clouds OCTAHEDRAL SQUARE PYRAMIDAL © 2012 Pearson Education, Inc. Chapter 7/47  The VSEPR Model: Six Charge Clouds SQUARE PLANAR Summary: Geometry around atoms with 2, 3, 4, 5 and 6 charge clouds Summary: Geometry around atoms with 2, 3, 4, 5 and 6 charge clouds The Valence Shell Electron Pair Repulsion Theory © 2012 Pearson Education, Inc.  Is the Molecule Polar? The more electronegative atom will pull the electron density of the bond closer to itself giving it a partial negative charge leaving the other atom with a partially positive charge. This is a dipole moment. Polar Covalent Bonds and Dipole Moments Polar Covalent Bonds and Dipole Moments C-Cl bond has a bond dipole because of a difference in electronegativities.  Polar Covalent Bonds and Dipole Moments The individual bond polarities do not cancel. Therefore, the molecule has a dipole moment. In other words, the molecule is polar.  Polar Covalent Bonds and Dipole Moments The individual bond polarities cancel. Therefore, the molecule does not have a dipole moment. In other words, the molecule is nonpolar.  Valence Bond Theory Valence Bond Theory: A quantum mechanical model which shows how electron pairs are shared in a covalent bond.  Valence bond theory Valence Bond Theory: A quantum mechanical model which shows how electron pairs are shared in a covalent bond.  B. Valence bond theory Covalent bonds are formed by overlap of atomic orbitals, each of which contains one electron of opposite spin. Each of the bonded atoms maintains its own atomic orbitals, but the electron pair in the overlapping orbitals is shared by both atoms. The greater the amount of overlap, the stronger the bond. Hybridization and sp3 hybrid orbitals How can the bonding in CH4 be explained? 4 valence electrons 2 unpaired electrons Hybridization and sp3 hybrid orbitals How can the bonding in CH4 be explained? 4 valence electrons 4 unpaired electrons Hybridization and sp3 Hybrid Orbitals How can the bonding in CH4 be explained? 4 nonequivalent orbitals Hybridization and sp3 hybrid orbitals How can the bonding in CH4 be explained? 4 equivalent orbitals Hybridization and sp3 hybrid orbitals Hybridization and sp3 hybrid orbitals sp2 hybrid orbitals sp hybrid orbitals  C. Molecular orbital theory: The hydrogen molecule Atomic Orbital: A wave function whose square gives the probability of finding an electron within a given region of space in an atom. Molecular Orbital: A wave function whose square gives the probability of finding an electron within a given region of space in a molecule. Molecular orbital theory: The hydrogen molecule σ bonding orbital σ* antibonding orbital  Molecular orbital theory: The hydrogen molecule σ bonding orbital σ* antibonding orbital (# bonding e– – # antibonding e–) Bond Order = 2  Molecular orbital theory: The hydrogen molecule 2–0 Bond Order = =1 2 Bond order, as introduced by Linus Pauling, is defined as the difference between the number of bonds and anti-bonds. The bond number itself is the number of electron pairs (bonds) between a pair of atoms. Bond number gives an indication of the stability of a bond.  Molecular orbital theory: The hydrogen molecule 2–1 1 2–2 Bond Order: = =0 2 2 2  Molecular orbital theory: Other diatomic molecules O2 O O Diamagnetic: All electrons are spin-paired. It is weakly repelled by magnetic fields. Paramagnetic: There is at least one unpaired electron. It is weakly attracted by magnetic fields. Oxygen, O2, is predicted to be diamagnetic by electron-dot structures and valence bond theory. However, it is known to be paramagnetic. Molecular orbital theory: Other diatomic molecules Molecular orbital theory: Other diatomic molecules Molecular orbital theory: Other diatomic molecules Combining Valence Orbital Theory and Molecular Orbital Theory  Molecular orbital theory: magnetic feature Diamagnetic: All electrons are paired. Weak repeal in magnetic field. Paramagnetic: Contains at least on unpaired electron. Weak attraction in magnetic field. O2 O O paramagnetic  Valence bond theory vs. Molecular orbital theory Valence bond theory Molecular orbital theory Localized electron model: Electrons Atomic orbitals no longer exist. in a molecule still occupy orbitals of Molecular orbitals are available for individual atoms. occupation by electrons. Bonds are localized between one Electrons are delocalized over the pair of atoms. entire molecule. Bonds formed from overlap of atomic Atomic orbitals are combined to and hybrid orbitals. form molecular orbitals. Predicts molecular shape based on Predicts arrangement of electrons regions of electron density. in molecules. Explains hybrid orbitals. Forms Explains and creates bonding and sigma and pi bonds. antibonding molecular orbitals. Explains the hybridization of atomic Does not explain the hybridization orbitals. of orbitals. 3. Metallic bond Sources of the Metallic Elements Sources of the Metallic Elements  Bonding in Metals Electron-Sea Model: A metal crystal is viewed as a three-dimensional array of metal cations immersed in a sea of delocalized electrons that are free to move throughout the crystal. © 2012 Pearson Education, Inc. Chapter 21/83  Bonding in Metals Molecular Orbital Theory for Metals (Band Theory)  Bonding in Metals Molecular Orbital Theory for Metals (Band Theory) Electrical Insulators: Materials that have only completely filled bands Metallic Conductors: Materials that have partially filled bands  Bonding in Metals Molecular Orbital Theory for Metals (Band Theory)  Semiconductors Semiconductor: A material that has an electrical conductivity intermediate between that of a metal and that of an insulator Valence Band: The bonding molecular orbitals Conduction Band: The higher energy antibonding molecular orbitals Band Gap: The energy difference between the valence and conduction bands Semiconductors Semiconductors  Types of chemical bonds Type of bonds Interaction Bond energy Strength Example (kJ/mol) IONIC Cation-anion 400-4000 Strong NaCl(s) Primary COVALENT Electron pairs 150-1100 Strong F2(g), bonds CH4(g) METALIC Cation and valence 75-1000 Strong Na(s), electrons Mg(s) ION-DIPOLE Charge and permanent 40-600 Strong NaCl (aq) dipole HYDROGEN Polar hydrogen and 10-40 Medium H2O-H2O BONDING non-bonding electron H2O- pair CH3OH DIPOLE-DIPOLE Permanent dipoles 5-25 Medium HCl-HCl Secondary bonds ION-INDUCED Ion and the dipole 3-15 Weak Fe2+-O2 DIPOLE induced by that DIPOLE-INDUCED Permanent dipole and 2-10 Weak HCl-Cl2 DIPOLE the dipole induced by that LONDON FORCE Temporary dipoles 0.05-40 Weak Ar(g)-Ar(g)  States of Matter and Their Changes Matter exists in any of three phases, or states— solid, liquid, and gas. The state depends on the relative strength of the attractive forces between particles compared with the kinetic energy of the particles. The greater the attractive forces between the molecules, the higher the boiling point.  Intermolecular Forces Intermolecular Forces: Attractions between “molecules” that hold them together. These forces are electrical in origin and result from the mutual attraction of unlike charges or the mutual repulsion of like charges. Types of Intermolecular Forces: Ion-Dipole Forces van der Waals Forces dipole-dipole forces London dispersion forces hydrogen bonds Intermolecular Forces  Intermolecular Forces Ion-Dipole Forces: The result of electrical interactions between an ion and the partial charges on a polar molecule © 2012 Pearson Education, Inc.  Intermolecular Forces Dipole-Dipole Forces: The result of electrical interactions among dipoles on neighboring molecules © 2012 Pearson Education, Inc. Intermolecular Forces Chapter 10/96 Dipole-Dipole Forces As the dipole moment increases the intermolecular forces increase. As the intermolecular forces increase, the boiling point increases. © 2012 Pearson Education, Inc.  Intermolecular Forces London Dispersion Forces: The result of the motion of electrons which gives the molecule a short-lived dipole moment. This induces temporary dipoles in neighboring molecules. © 2012 Pearson Education, Inc.  Intermolecular Forces London Dispersion Forces Polarizability: the ease with which an electron cloud can be deformed. The larger the molecule (the greater the number of electrons) the more polarizable. The larger the molecular mass and/or the surface area, the greater the temporary polarization of a molecule. London Dispersion Forces vs Dipole-dipole Forces 99  Intermolecular Forces: hydrogen bond Hydrogen Bond: An attractive force between a hydrogen atom bonded to a very electronegative atom (O, N, or F) and an unshared electron pair on another electronegative atom.  The Hydrogen Bond water waterhas has water water thehashas water the has thehighest lowest theheat highest highest the heat highest of molarofmass fusion melting vaporization point boiling point The melting point, boiling point, heat of fusion and heat of vaporization of water are extremely high and do not fit the trend of properties relative to molar mass within Group VIA. Water exhibits these unusual properties because of hydrogen bonding between water molecules.  Formation of hydrogen bond 1. One molecule must contain at least one H atom attached to a highly electronegative atom (i.e. F, O or N). 2. The other molecule must contain an F, O or N atom that provides the lone pair of electrons. Types of Crystals Life would be impossible without H-bond It allows oxygen to dissolve in water Intermolecular forces define the structure of biomolecules © 2016 Pearson Education, Ltd. H-bonding - Properties of Water The effect of temperature on the density of water – ice floating Density maximum at 3.98 oC 1.0 0.999 Density, g/mL Supercooled water 0.998 WATER 0.997 ICE 0.910 0 5 10 Temperature, oC  H-bonding - Properties of Water H-bonding in water 117 pm Each and every one of the water molecules can form four hydrogen bonds, an elaborate network of molecules is formed. 96 pm In liquid water the less Ordered three- ordered arrangement of dimensional network of the molecules without hydrogen bonds in ice. fixed distances between them results in higher density at low Solid Liquid temperature H-bonding - Properties of Water Viscosity: The measure of a liquid’s resistance to flow. Surface Tension: The resistance of a liquid to spread out and increase its surface area. water Intermolecular Forces - Evaporation and Vapor Pressure Vapor Pressure: The partial pressure of a gas in equilibrium with liquid at a constant temperature Stronger intermolecular forces → smaller vapor pressure Higher temperature → larger vapor pressure