SCH4U UNIT 1 - Structures and Properties Mia Wong, 2019 1 Atomic Theory to Bohr - Development of the Modern Atomic Theory Early History of Chemistry Fundamental Chemical Laws Early Atomic Theory 3 3 3 3 Dalton’s Atomic Theory Thomson’s Model of the Atom Discovering the Electron - J.J. Thomson’s Cathode Ray Experiment 4 4 5 Robert Millikan’s Oil Drop Experiment Rutherford’s Atomic Model Problems with Rutherford’s Planetary Model 5 6 7 Classical Theories of Light Wave Theory of Light Basic Wave Terms 7 7 7 Planck’s Quantum Hypothesis Photoelectric Effect Photons 8 8 8 Atomic Spectra Line Spectra of Hydrogen Bohr’s Model of the Atom 8 9 10 Success and Failures of Bohr’s Model Henri Burquel’s Discovery of Radioactivity Characteristics of Radioactive Emissions 10 10 11 Atoms and Isotopes 11 Electron Cloud Model Quantum Mechanics Quantum Numbers Electron Configuration 12 12 12 13 Writing Electron Configuration Energy Level Diagrams Pauli’s Exclusion Principle 13 13 14 Aufbau Principle Hund’s Rule 14 14 General Periodic Trends Factors Affecting Properties Coulomb's Law 15 15 15 Atomic Size Ion Sizes Ionization Energy (IE) 15 15 16 Trends in Ionization Energy Reactivity of Elements Reactivity of metals: 16 16 16 Reactivity of non-metals: 16 2 Electron Affinity (EA) Electronegativity (EN) Chemical Bonding Intermolecular Forces London Dispersion/Van der waals Forces 17 17 18 18 18 Dipole-Dipole Forces Ion-Dipole Forces Induced Dipole Forces 18 18 18 Hydrogen Bonding Intramolecular Forces Chemical Bonds 19 19 19 Electronegativity and Electron Sharing Crystalline Solids Amorphous Solids 20 20 20 Metallic Bonding Properties of Metals and Metallic Solids Alloys 20 20 21 Ionic Crystals Properties of Ionic Compounds Covalent Bonding 21 21 21 Atomic Solids Molecular Solids Polar Molecular Compounds 21 22 22 Properties of Molecular Compounds Unusual Properties of Water Network solids 22 22 22 Ionic Solids 22 Bonding Theory Free Radicals Resonance Valence Bond Theory 23 23 23 23 Sigma () bond Pi (π) bond Molecular Orbital Theory 23 24 24 Formation of Molecular Orbitals in CH4 sp3 Hybridization and Shape Hybrid Orbitals 24 25 25 Hybrid Orbitals of Exceptions to the Octet Rule 25 Atomic Theory to Bohr - Development of the Modern Atomic Theory 3 Early History of Chemistry - 1,000 B.C.: ore processing for weaponry and ornaments; use of embalming fluids 400 B.C.: alchemy, pseudoscience in which it was believed that metals could be turned to gold, was practiced 16th century: Georg Bauer refined metal extraction from ores; Paracelsus used minerals for medicinal application 17th century: Boyle discovered that pressure is inversely proportional to volume (Boyle’s Law) 17th and 18th centuries: Georg Stahl claimed “phlogiston” was contained in combustible materials (Phlogiston Theory) 1772: Oxygen was discovered by Joseph Priestley Late 18th century: combustion is studied intensively Fundamental Chemical Laws - - - Antoine Lavoisier, referred to as the “Father of Modern Chemistry”, discovered the nature of combustion, published the first modern chemistry textbook which included the Law of Conservation of Mass 1804: John Dalton states the Law of Multiple Proportions: When two elements combine to form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers. 1808: John Dalton states Law of Definite Proportions (developed by Joseph Proust): A given compound always contains exactly the same proportions of elements by mass. 1811: Avogadro’s Hypothesis: At the same temperature and pressure, equal volumes of different gases contain the same number of particles. Early Atomic Theory - Ancient Greeks (e.g. Democritus) proposed that when matter is divided into smaller & smaller pieces, a finite limit known as the atom (atomos) is ultimately reached Figure 1.1: Empedocles (circa 494 BC - circa 434 BC), a Greek philosopher, proposed that all structures of the Earth are created from 4 root elements: fire, air, earth, and water. Dalton’s Atomic Theory 1) All matter is made of atoms. These indivisible and indestructible objects are the ultimate chemical particles. 4 2) All the atoms of a given element are identical, in both weight and chemical properties. However, atoms of different elements have different weights and different chemical properties. 3) Compounds are formed by the combination of different atoms in the ratio of small whole numbers. 4) A chemical reaction involves only the combination, separation, or rearrangement of atoms; atoms are neither created nor destroyed in the course of ordinary chemical reactions. - Led to Billiard Ball Model of indivisible atoms of elements (these combined to make all known compounds) Thomson’s Model of the Atom - Detected and measured the mass of a beam of negatively charged particles in a vacuum tube/cathode ray tube (CRT) Called tiny particles that emerged from a metal cathode “electrons” Explained electrical conduction in metals by the movement of subatomic electrons in the solid Electrical conduction in solutions was explained by the existence of charged atoms called ions Since a negatively charged particle was found, a positive charge must also exist Created the “Plum Pudding” Model Thomson’s model of the atom was similar to plum pudding, in which raisins were electrons, suspended in pudding, a positively charged medium Figure 1.2: Visual representation of Thomson’s model. Electrons are particles suspended in a positively charged medium, similar to the raisins suspended in Christmas pudding. 5 Discovering the Electron - J.J. Thomson’s Cathode Ray Experiment Figure 1.3: CRT has a glass body with electrodes (cathode, -, and anode, +) had negatively charged particles shoot from the cathode towards the anode which had a hole in the centre. This caused a visible ray to shoot straight through the anode at a high speed which the fluorescent screen helped to amplify the point B at which the ray hit the glass. (Ray is visible as a result of high voltage.) When putting two charged plates on the top and bottom of the tube, it bent the ray towards point A. When a magnet was added, the point at which the ray hit the fluorescent screen became point C. - Since electrodes were composed of different kinds of metals, all atoms must contain electrons Since there is a negative charge, the atom must also contain a positive charge Conclusions resulted in the Plum Pudding Model (electrons or plums suspended in a positive medium or pudding) Experiment allowed him to find the charge (Coulombs) to mass (kilograms) ratio but not the exact charge nor mass of a singular electron Robert Millikan’s Oil Drop Experiment Figure 1.4: Atomizer sprayed oil droplets which fell through a pinhole in a metal plate. An X-Ray was used to ionize (-) the droplets, then the velocity at which they fell was measured. After, a voltage was applied with a positive voltage through the first metal plate and a negative voltage through the second metal plate until the droplets were suspended, meaning gravitational force and electrical force reached an equilibrium. - Found charge (1.6 x 10-19 C) and mass (9.11 x 10-31 kg) of one electron using charge to mass ratio 6 Rutherford’s Atomic Model - Discovery of radioactivity (α, β, and γ rays) allowed Rutherford to probe inside atom Based on Rutherford’s model, he proposed alpha (α) rays (high energy He2+ ions - which is simply 2 protons) should pass through the positive “pudding” in a very thin sheet of gold (Au) foil Figure 1.5: Radium has a radioactive nucleus which makes it emit alpha particles. Alpha particle beam was expected to shoot directly through the gold foil with little deflection, since Thomson’s model concluded that the electric fields in an atom would be too weak to affect the fast-moving alpha particles. The actual result was alpha particles mostly going straight through with some having a large degree of deflection. Figure 1.6: Thomson’s model would have resulted in the alpha particles passing straight through. This was the expected result. Figure 1.7: Rutherford observed that the alpha particles sometimes had a large degree of deflection. - Concluded that there is a dense positive nucleus because of the deflection degree Proton was found in 1919 Since the positive charge is in the centre it means that the rest of the atom is 99.99% empty space with electrons orbiting the nucleus In 1932, Chadwick and Rutherford found that the nuclei masses did not match with masses of protons, this led to the discovery of neutral particles called neutrons 7 Problems with Rutherford’s Planetary Model 1) According to classical physics, electrons orbiting the nucleus should lose kinetic? energy and emit light. Loss of energy would cause electrons to spiral into the nucleus, resulting in the collapse of the atom. Note: emission of light causes loss of energy 2) According to classical theories of light, excited electrons should emit a continuous spectrum of white light. Instead emission spectrum of all elements is unique (similar to a fingerprint). Classical Theories of Light - Light is a form of electromagnetic radiation Visible light is a portion of the spectrum ranging between 390 nm and 750 nm Wave Theory of Light - All electromagnetic radiation is made up of electric and magnetic fields which oscillate in a wave pattern as it moves through space Electromagnetic radiation travels at a constant speed in a vacuum but wavelength and frequency (f) varies Basic Wave Terms - c = speed of light λ (lambda) = wavelength (m) f = frequency/cycles per sec (s-1 or Hz) Figure 1.8: Typically measured in nanometers (nm), 1 nm = 1 x 10-9 m. This would mean that 1 nm wavelengths would fall under the X-Ray portion of the wave spectrum. Wavelength increases going right. Frequency increases going left. Planck’s Quantum Hypothesis - Studied radiant energy emitted by solid bodies (blackbodies) heated to incandescence 8 - Blackbodies always glow red, white, then blue Realized spectrum produced by white light was NOT continuous According to classical physics, energy should go up, but it reaches a peak and decreases Means that electrons can be lost OR gained E = hf Note: energy of mole (n) of photons can be found using E = nhf 2 Note: 1 Joule = 1 kg × ms - E = energy of a single photon (kJ) - n = integer - h (Planck’s constant) = 6.63 x 10-34 J x s - f = frequency (Hz, s-1, cycles/sec) - As atoms oscillate, radiation (energy) is emitted in bursts/a certain quantity separately instead of continuously - Proposed that light was composed of small packets (bursts) or quanta of energy called photons - Different colours of light are composed of photos with different “quanta” of energy - Energy of a particular photon is proportional to the frequency of radiation Photoelectric Effect - When light shines on metal, it emits electrons (not equivalent of β particles) Low frequency of light (e.g. red) shone on metal does not cause emission of electrons High frequency of light (e.g. violet) shone on metal causes emission of electrons Photons - Single unit of light energy, multiple of energy quanta Photon energy transfers to electron, then breaks away because remainder of energy becomes kinetic energy When frequency is below a certain threshold, electrons cannot escape Kinetic energy of electrons breaking away is dependent on photon energy Atomic Spectra - Spectroscopy: study of spectra (plural of spectrum) Spectrophotometers take light through a prism, then detects absorbed/transmitted waves Line Spectra of Hydrogen - Voltage runs through H2, H-H bonds break away and emit wavelengths from an excited state 9 - Emission spectrum: spectrum seen when electromagnetic radiation passes through a spectrophotometer 1) Continuous spectrum - all wavelengths are only at a particular electromagnetic spectrum region 2) Line spectrum - wavelengths are only at particular regions - Hydrogen has a line spectrum - Atomic spectra are unique to each element, it can be used to identify an atom - Since electrons can only exist at certain energy levels for each element, only certain quanta of energy are given off - Colour of light depends on the quantum’s energy Figure 1.9: The H2’s wavelengths pass through a slit which travels through a prism which emits the spectrum of H2 on a photographic plate. Note that this is not a diagram of a standard spectrophotometer’s inner workings. Rydberg Equation: 1 = R∞ n12 − n12 λ ( 1 2 ) Bohr’s Equation: E n = 2.18 × 10−18 J ( n12 ) Note: Bohr’s equation shows only valid energies of transitions for hydrogen’s single electron - As the electron gets closer to the nucleus, En becomes larger in absolute value but also increasingly negative Bohr’s Model of the Atom - As electrons move further away from the nucleus, energy increases Ground state: lowest energy state of an atom Excited state: each energy state in which n > 1 (gained energy) Electrons moving from ground state to higher states require energy; electrons falling from a higher state to ground state releases energy (photon emission) 10 - Bohr developed a model of the atom which explained line spectrum of hydrogen and why the atom does not collapse 1) Electrons can only move in certain fixed orbits. Each orbit corresponds to a specific energy level and an electron can move within an orbit without losing any energy 2) An electron can only move from one orbit to another when it gains/loses energy Success and Failures of Bohr’s Model - Correctly illustrates first 20 elements Cannot explain elements past calcium Henri Burquel’s Discovery of Radioactivity Figure 1.10: Since radioactive atoms (Uranium) emit alpha, beta, and gamma rays, the direction of rays depends on the electric field surrounding it. Supposing that there is a photographic plate in front of the Uranium’s emission direction (which is the method in which Burquel discovered radioactivity), the particle beams were shown on the plate. - Currently, it is known that radioactivity is the spontaneous decay of an atom’s nucleus (first proposed by Rutherford) Characteristics of Radioactive Emissions Alpha Particles Beta Particles Gamma Rays Symbol α/42α/42He β/β-/e γ Atomic Mass (u) 4 1/2000 0 Charge 2+ 1- 0 11 Speed Slow Fast Very fast (speed of light) Ionizing Ability High Medium None Penetrating Power Low Medium High Stopped by Paper Aluminum Lead Atoms and Isotopes - Radioactivity is caused by an unstable nucleus (result of certain isotopes) Isotopes: atoms of the same element with differing numbers of neutrons Radioisotopes’ nuclei decays and emits gamma rays and/or subatomic particles 12 Electron Cloud Model Quantum Mechanics - Electrons orbit the nucleus in probabilistic clouds Orbits are in complex shapes Based on the location of electrons within a given atom 90% of the time Quantum numbers system presents the electron distribution in an atom Quantum Numbers 1) Principal Quantum Number (n): positive whole number specifying energy level of an atomic orbital and its size, higher values indicate a higher energy level and size - n=1-∞ 2) Secondary Quantum Number (l) : “orbital angular momentum quantum number” or “orbital-shape quantum number”; indicates orbital shape - l = 0 - (n - 1) - 0 = “s” orbital (sphere) - 1 = “p” orbital (dumbell) - 2 = “d” orbital (clover) - 3 = “f” orbital 3) Magnetic Quantum Number (ml ): specifies the orientation of a given orbital around the nucleus - ml = ( - l) - (+ l) - (2l + 1) possibilities/orientations 4) Spin Quantum Number (ms): indicates whether an electron spins upward or downward - ms = 12 or (- 12 ) - First electron in a given subshell will spin upward, second will spin downward Figure 1.11: Visual representation of possible configurations of “s”, “p”, “d”, and “f” orbitals, effectively illustrating first three quantum numbers. 13 Electron Configuration - Shorthand to indicate electron arrangement around an atom’s nucleus Total of 7 energy levels denoted by the primary quantum number and secondary quantum number First energy level: - Represented by “1s” - Each “s” sublevel (including sublevels in higher energy levels) can hold 2 electrons Second energy level: - Represented by “2s” and “2p” - Each “p” sublevel (including sublevels in higher energy levels) can hold 6 electrons Third energy level: - Represented by “3s”, “3p”, and “3d” - Each “d” sublevel (including sublevels in higher energy levels) can hold 10 electrons - “3d” sublevel does NOT fill after “3p” sublevel, it fills after “4p” sublevel Fourth energy level: - Represented by “4s”, “4p”, “4d”, and “4f” - Each “f” sublevel (including sublevels in higher energy levels) can hold 14 electrons Subsequent energy levels: - Represented by “s”, “p”, “d”, and “f” - Sublevels are generally filled in the order “1s, 2s, 2p, 3s, 3p, 4s, 4p, 3d, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, 6f, 7d, 7f” Writing Electron Configuration 1) 2) 3) 4) Write energy level (primary quantum number) Write letter representing orbital shape (secondary quantum number) Write the number of electrons in sublevel as a superscript (ex. He - 1s2) Repeat for subsequent energy levels Energy Level Diagrams - Represents each electron’s placement in an atom Follows Pauli’s Exclusion Principle, Aufbau Principle, and Hund’s Rule Figure 1.12: Electron configuration for carbon is written, along with its energy level diagram. Pauli’s Exclusion Principle - No 2 electrons can have the same set of quantum numbers 14 - 2 electrons per orbital Aufbau Principle - Electrons fill according to orbital energies from lowest to highest “d” orbitals are exceptions to the Aufbau Principle, 3d will fill after 4p is full Hund’s Rule - Most stable arrangement for electrons in orbitals of equal energy is where the number of electrons with the same spin is maximized Electrons will not pair in an orbital unless required 15 General Periodic Trends Factors Affecting Properties - Elemental properties are related to attractive forces between the nucleus and electrons Force of attraction is dependent on 2 factors: distance and charge Coulomb's Law - Quantifies electrostatic attraction between charged particles F=k( - q1q2 r2 ) F = force of attraction between opposite charges q1 = charge of nucleus (# of protons) q2 = charge of electron r = distance from nucleus to electron (atomic radius/size) k = constant (integer) Atomic Size Increases down a group: - Electrons are added further from the nucleus, therefore there is less attraction - Additional energy levels and shielding effect - Additional energy levels “shields” electrons from being pulled in towards nucleus - Added energy levels - Going down a group increases atomic size more than going left across a period Decreases going right across a period: - Effective nuclear charge increases - Effective nuclear charge = # of protons - # of core electrons (all electrons except for the valence electrons) - Going right across a period decreases atomic size less than going up a group Ion Sizes - Cations get smaller because attraction increases Anions gets larger because attraction decreases Anions are larger than the atoms form in which they come Consider electron-proton attraction 16 Ionization Energy (IE) - Defined as the energy required to remove an electron from the highest energy level of an atom in a gas state Recall the photoelectric effect: electrons are rejected only when E = h υ exceeds a threshold First ionization energy: energy required to remove first electron Second ionization energy: energy required to remove second electron Etc. Trends in Ionization Energy - Increases going right across a period Decreases going down a group Energy required for removal of first electron is lower than the energy required from removal of the second electron Ex. Beryllium: When removing the second electron, the nucleus’ attractive force is spread across 3 electrons instead of 4. If a third electron is removed, ionization energy will have a huge jump because its orbital is closer to the nucleus than the 3rd and 4th. Figure 1.13: Electron configuration shows the orbital of each electron in beryllium. Proximity of electrons in orbital “1s” is greater than that of electrons in orbital “2s”. Reactivity of Elements - If an atom has a low ionization energy, it will have a high reactivity since it gives up electrons readily Reactivity of metals: - Increases towards francium Decreases across a period Increases down a group Reactivity of non-metals: - Increases across a period Decreases down a group Electron Affinity (EA) - Amount of energy released when an electron is added to a neutral atom Increases going across a period 17 - Decreases going down a group Electronegativity (EN) - Ability of an atom to attract electrons to itself in a covalent bond Atoms with fewer energy levels can attract electrons more strongly (smaller distance, less shielding) Increases going across a period Decreases going down a group Difference of electronegativity defines whether a given chemical bond is ionic, polar covalent, or non-polar covalent 18 Chemical Bonding Intermolecular Forces - Attractive forces holding molecules together “Inter”: between Weakened when melting point is reached Breakage occurs when boiling point is reached London Dispersion/Van der waals Forces - Result from random motion of electrons Occurs when electron concentration is higher in one region than another Forces at particular instant are known as instantaneous dipoles or dipole moments At low temperatures, dipoles can be induced in other atoms causing solidification in Helium Relatively weak Dipole-Dipole Forces - Polar molecules attempt to line up “positive” dipole ends with “negative” dipole ends Requires polar substance that has a “permanent dipole” Dipole aligning affects melting and boiling points Relatively weak Ion-Dipole Forces - Attractive force between an ion and a polar end of a molecule with the opposite charge Explains ionic compounds’ solubility Induced Dipole Forces - - An ion-induced dipole force occurs when an ion approaches a nonpolar molecule and shifts the electron density in the molecule, making it temporarily polarized and attracted to the ion A dipole-induced dipole force occurs when a permanent dipole approaches a nonpolar molecule and shifts the electron density in the molecule, making it temporarily polarized and attracted to the permanent dipole Hydrogen Bonding - A particularly strong dipole-dipole force - 19 Between a polarized hydrogen atom in a molecule and an electronegative atom (N, O, or F) from another molecule Hydrogen must be bonded to an electronegative atom such as N, O, or F 10X weaker than covalent bonds 10X stronger than regular dipole-dipole forces Increases melting and boiling points significantly Intramolecular Forces - Attractive forces bonding atoms together to create a molecule “Intra”: within Breakage results in decomposition of a molecule Chemical Bonds - Ionic bonds Covalent bonds Metallic bonds Influences shape of molecules, molecular shape, and properties Figure 1.14: Visual representation of methods of chemical bonding between atoms. Electronegativity and Electron Sharing - Δ EN between atoms can be ionic, intermediate or high Δ EN f rom 1.7 - 3.3 is mostly ionic Δ EN f rom 0.4 - 1.7 is polar covalent Δ EN f rom 0.0-0.4 is non-polar or mostly nonpolar 20 Crystalline Solids - Includes atomic solids, molecular solids, and network solids In organized particle arrangements Amorphous Solids - Indistinct shapes Disorganized particle arrangements Metallic Bonding - “Electron-sea” model Valence electrons move freely among ions Electrons do not belong to any single nucleus Delocalized electrons hold metal ions in place rigidly Strong bonding force Aggregates of crystals Properties of Metals and Metallic Solids - High melting point and boiling point Metals gain metallic character going up and across the periodic table q q Recall: F = k ( 1r2 2 ), a smaller radius results in a stronger force between opposite charges Free moving electrons cause electrical and thermal conductivity in solid and liquid forms Metal ions can slide by one another easily with electrons constantly surrounding it causing malleability and ductility Variation between metals is because of crystal sizes; the smaller the atomic size of a metal, the harder the pure metal will be Crystalline solid Few valence electrons Low ionization energies Lustrous because of delocalized electrons which emit and absorb various wavelengths Alloys - Can be composed of 2 or more metals or one or more metals and nonmetals Can significantly affect a metal’s properties If atoms of added element are of similar size to metal atoms, they will replace metal atoms in the structure If atoms of added element are smaller than metal atoms, they will fit in between metal atoms in the structure 21 Ionic Crystals - Crystal lattice composed of alternating positive and negative ions Unit cell is the smallest repeating ion groups Increasing ionic radius decreases melting and boiling points because of lower attractive forces Increasing ionic charge increases melting and boiling points because of higher attractive forces Properties of Ionic Compounds - High melting point and boiling point Attractive forces between ions and water is stronger than intermolecular forces between the ions themselves contributes to solubility Hard, brittle and breaks apart when struck because of its crystalline structure Conducts electricity in water because ions can move within the fluid Covalent Bonding - Polar: unequal sharing of electrons Non-polar: equal sharing or near equal sharing of electrons Atomic Solids - Crystalline solid Noble gases form liquids and solids at low temperatures due to weak bonds between atoms Since attractive forces are weak, they are weakened and broken at low temperatures As the number of electrons increases, the melting point and boiling point also increase As molar mass increases, London Dispersion Forces increase Very low melting and boiling points that increase going down a group No electrical conduction Molecular Solids - Crystalline solid Substances with covalent bonds or polar covalent bonds Can exist in all states at room temperature Nonpolar Molecular compounds only have London Dispersion Forces between molecules Forces of attraction are greater between molecules with a greater number of atoms Low melting and boiling points 22 Polar Molecular Compounds - Compounds with bond dipoles and molecular dipoles have high melting and boiling points due to intermolecular forces between permanent dipoles Crystalline solid Properties of Molecular Compounds - No electrical conductivity since electrons cannot move freely Relatively low melting and boiling points because of weak intermolecular forces Greater intermolecular force strength result in a higher melting and boiling point Unusual Properties of Water - Dubbed the “universal solvent” because it can dissolve polar molecules and ionic compounds Expands when frozen because of the organization of hydrogen bonds in the solid Network solids - Covalent network crystal Individual atoms bound in covalent bonds forming an interwoven network High melting and boiling points Extreme hardness Crystalline solid Ionic Solids - Hard and brittle solids formed by metal cations and nonmetal anions Held together by electrostatic forces (ionic bonds) High melting and boiling points Electrically conductive in liquid state Crystalline solid 23 Bonding Theory Free Radicals - Atoms and/or molecules with unpaired electrons are reactive substances Due to an odd number of electrons Resonance - When 2+ Lewis structures can be drawn, “hybrid” structure is assumed Electrons are assumed to be “delocalised” Figure 1.15: Resonance structures of nitrite ion. There are 2 resonance structures, neither of which are the actual structure. The actual structure is a resonance hybrid of both. Both resonance structures can also be presented as one structure as seen on the most right diagram. Valence Bond Theory - Quantum mechanical model of bonding Covalent bonds form when a pair of half-filled orbitals overlap to form combined (bonding) orbitals Contains 2 electrons with opposite spins Electron density is highest between the 2 nuclei Sigma ( σ ) bond - Result of direct orbital overlap Overlapping orbitals can form between s and p Combined orbitals represents a lower energy state Figure 1.16: Examples of sigma bonds resulting from overlapping orbitals. When put on the same axis specified and lined up, the unshaded parts of each orbital show the regions of direct overlap. The shaded areas are the regions that do not participate in the sigma bond. 24 Pi ( π ) bond - Sigma bonds result from direct orbital overlap and only one pair of electrons can overlap directly within an atom Explains double and triple bonds between atoms Occur when one or more pairs of electrons is shared between 2 atoms Require unhybridized orbitals because the space between orbitals is already taken up by a sigma bond Enclose a sigma bond on two sides “Sideways” bonding Figure 1.17: Bonding of ethylene is shown according to valence bond theory. Carbon-carbon double bond is pictured in which the sigma bond is enclosed on the top and bottom by one pi bond. Molecular Orbital Theory - Lewis Theory considers all valence electrons as identical Contradicted by Wave-Mechanical Model in which there are different atomic orbitals in each energy level Experimental evidence confirmed Lewis Model of covalent bonding in compounds States atomic orbitals combine to make molecular orbitals Molecular orbitals are combinations of Schrodinger’s equation containing multiple nuclei Formation of molecular orbitals involves electrons promotion and orbital hybridization Formation of Molecular Orbitals in CH4 1) 2nd “s” electron is “promoted” into empty “p” orbital 2) 2s and 2p atomic orbitals undergo hybridization to form sp3 orbitals (1s and 3p make 4 orbitals) 3) Each identical sp3 orbital can form sigma bond with another half-filled orbitals sp3 Hybridization and Shape - Electron repulsion creates a tetrahedral shape (AX4) with bond angles of 109.5° 25 Figure 1.18: Visual representation of sp3 orbital composed of 4 sp hybrid orbitals with one large lobe and one small lobe. Hybrid Orbitals - s + p → sp (linear) - s + 2p → sp2 (trigonal planar) - s + 3p → sp3 (tetrahedral) Note: sp hybridized orbitals create a dumbbell shape with one large lobe and one small lobe Hybrid Orbitals of Exceptions to the Octet Rule - s + 3p + d → sp3d (trigonal bipyramid) s + 3p + 2d → sp3d2 (octahedral)
