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Chemistry
Course
CHEM 1AA3
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David Brock
Semester
Spring

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Chemistry Notes 1AA3 Chapter 11: Chemical Bonding II – Additional Aspects 11.1 What a Bonding Theory Should Do When two atoms are brought together there are 3 types of interactions between them: 1. Electrons are attracted to the two nuclei 2. Electrons repel each other 3. The 2 Nuclei repel each other Potential Energy (of electrons): the net energy of these interactions between atoms Potential Energy vs. Distance Graph • Plots potential energy of atoms over the distance between the atomic nuclei • When the potential energy is negative, there is a net attractive force between the two atoms • When the potential energy is positive, there is a net repulsion between the two atoms • there is little interaction between the atoms when they are far apart • at intermediate distances the attractive forces are predominant over repulsive forces • at small distances the repulsive forces are predominant over attractive forces • when potential energy is lowest: o lowest potential energy = - bond dissociation energy o distance at lowest potential energy = bond length Bonding Theory should explain • bond-dissociation energies, bond lengths, bond angles etc. Positive Negatives Lewis • easy to apply • no quantitative information about bond Theory energies, lengths and angles • written quickly VSEP • easily proposes molecular • no quantitative information about bond R shapes energies, lengths and angles Theory • generally in good • resonance structures (not possible to agreement with experimental represent) results • odd-electron species problem 11.2 Introduction to the Valence-Bond Method Valence-Bond Method: • when atoms approach each other, areas of high electron probability begin to interpenetrate (or overlap) • thus a bond is formed due to the high electron probability found in the overlap region (and thus covalent bonds are formed) • valence method gives a localized electron model of bonding: o core electrons and lone-pair valence electrons retain the same orbital locations as in the separated atoms o charge density of the bonding electrons is concentrated in the region of orbital overlap • energies of atomic orbitals differ from one type to another so the valence-bond method implied different bond energies for different bonds (supported by experimental observations) • the Lewis structures provide no information about bond enerfies 11.3 Hybridization of Atomic Orbitals Expanding the Valence-Bond Method: • when we extend the unmodified valence-bond method to greater number of molecules our descriptions of molecular geometry based on the simple overlap of unmodified atomic orbitals do not conform to observed measurements (of bond angles and energies etc.) • for the method to work with greater number of molecules we infer that bonded atoms have different orbitals from isolated, non-bonded atoms • atomic orbitals of bonded atoms must be modified Hybrid Orbitals: new orbitals formed by hybridization For Example: • atomic orbitals are mathematical expressions of electron waves in an atom • algebraic combination of the wave equations of the 2s and three 2p orbitals of the carbon atom produces four new sp hybrid orbitals • found in the methane molecule • the angle between any two of the orbitals is the tetrahedral angle: 109.5 • the hybridization of the carbon orbitals creates 4 valence electrons, which then overlap with the s orbitals of the Hydrogen atoms Hybridization: mathematical process of replacing pure atomic orbitals with reformulated atomic orbitals for bonded atoms • the number of hybrid orbitals equals to the total number of atomic orbitals that are combined • symbols ident3fy the numbers and kinds of orbitals involved o an sp orbital combines one s orbitals and three p orbitals • the energy of the orbitals also change in hybridization o in the sp orbitals the three p orbitals move down by ¼ of the energy difference between the s and p orbitals o the one s orbital moves up by ¾ of the energy difference between the sp and orbitals o thus energy is conserved • the orbital(s) with the higher energy move down by 1/(total number of final hybrid orbitals) of the energy difference • the orbital(s) with the lower energy moves up by the rest of the energy difference • in the end, the energy of the similarly combined hybrid orbitals are the same 3 • in the sp orbitals on the central atom is hybridized, the rest remain unhybridized The Hybridization Method: • not an actual physical phenomenon • we cannot observe electron charge distributions changing from pure to hybrid orbitals • for some covalent bonds, no single hybridization scheme works well • works very well for carbon-containing molecules thus used a great deal in organic chemistry • hybridization schemes are well established and very commonly encountered, particularly among the second-period elements Bonding in H O 2nd NH 3 • both require a tetrahedral electron-group geometry for four electrons groups • thus requires an sp hybridization scheme • in NH one sp orbital is occupied by a lone electron pair and the rest are involved in bond formation: this 3 agrees with the VSEPR theory • however H-O-H and H-N-H bond angles expected for 1s and 2p atomic orbital overlaps are 90 o o evidence to suggest unhybridized p orbitals based on spectroscopic evaluation o maybe be because of the ionic character of O-H and N-H bonds that cause repulsions between the positive partial charges associated with the H atoms and force the H-O-H and H-N-H bonds o to open up greater than 90 • bonding in H O2and NH is l3rgely unsettled and underscores the occasional difficulty in finding a single theory that is consistent with all the available evidence 2 sp Orbitals: • sp orbitals are formed with most boron compounds, which corresponds to the trigonal planar electron geometry and observed experimental bond angles 2 • combines one 2s and two 2p orbitals into three sp hybrid orbitals and leaves one p orbitals unhybridized • the angle between any two of the orbitals is 120 o • boron has four orbitals but only three electrons in its valence shell sp Hybrid Orbitals: • Beryllium has four orbitals and only two electrons in its valence shell • Most beryllium compounds hybridize one 2s and one 2p orbital of Beryllium into two sp hybrid orbitals and leave two 2p orbitals unhybridized o • The angle between the two orbitals is 180 sp d and sp d Orbitals: • To describe hybridization that correspond to the 5 and 6 electron group geometries we need to go beyond the s and p subshells of the valence shell • sp d orbitals are found in P-Cl bonds in PCl wh5ch hybridize 1 s orbital, 3 p orbitals and 1 d orbital into 5 3 half-filled sp d orbitals • sp d orbitals are found in S-F bonds in SF whi6h hybridize 1 s orbital, 3 p orbitals and 2 d orbitals into 5 3 half-filled sp d orbitals • recent theoretical considerations cast serious doubt on d-electron participation • the five sp d orbitals are directed to the corners of a trigonal bipyramid • the six orbitals are directed to the corners of a regular octahedron Hybrid Orbitals and the Valence-Shell Electron-Pair Repulsion (VSEPR) Theory • Hybridization was introduced by Linus Pauling in 1931 before the VSEPR Theory • The ideas expressed in the VSEPR Theory was first introduced by H.E. Powell and N.V. Sidgwick in 1940 and then developed into a set of rules by Ronald Gillespie and Ronald Nyholm in 1957 • The VSEPR Theory has a predictive capability based on Lewis structures whereas hybridization schemes require a prior knowledge of the molecular geometry Choosing the likely Hybridization Scheme for a central atom in a structure in the valence-bond method: • Write a plausible Lewis structure for the species of interest • Using VSEPR theory to predict the probable electron-group geometry of the central atom • Selecting the hybridization scheme corresponding to the electron-group geometry • The hybridization scheme adopted for a central atom should be the one producing the same number of hybrid orbitals as there are valence-shell electron groups and in the same geometric orientation 11.4 Multiple Covalent Bonds • Two different types of orbital overlap occur when multiple bonds are described by the valence-bond method Bonding in C H2: 4 o • Ethylene: Planar molecule with 120 H-C-H and H-C-C bond angles • VSEPR theory treats it as a C atom with three electron groups in a trigonal-planar fashion 2 • Hybridization scheme that produces a set of hybrid orbitals with a trigonal-planar is sp • One of the bonds between the carbon atoms results from the overlap of sp hybrid orbitals from each atom • This overlap occurs along the line joining the nuclei of the two atoms • Sigma (σ) bond: orbitals that overlap in an end-to-end fashion linking the nuclei of two atoms • A second bond between the C atoms results from the overlap of the unhybridized p orbitals • There is another region of high electron charge density above and below the plane of the carbon and hydrogen atoms • Pi (π) bond: bond produced by side-to-side overlap of two parallel orbitals Summary: • The shape of a molecule is determined only by the orbitals forming sigma bonds (the σ-bond framework) • Rotation about the double bond is severely restricted o Twisting one CH g2oup out of the plane of the other would reduce the amount of overlap of the p orbitals making the pi bond weaker o The double bond is rigid and the C H2m4lecule is planar • A σ + π bond is stronger than a single σ bond but not twice as strong Illustration: (pg. 435) Bonding in C H2: 2 • Acetylene: Similar to C2H4except: o The Lewis structure of acetylene features a triple covalent bond H-C---C-H o The molecule is linear as found by experiment and expected from VSEPR Theory o A hybridization scheme to produce hybrid orbitals in a linear orientation is sp • In the triple bond in acetylene one of the C-C bonds is a σ bond and two are π bonds • Drawing three dimensional sketches to show orbital overlaps is not easy • It is simpler to draw them two-dimensional and label the bonds as π or σ Chapter 8: Organic Chemistry 26.1 Organic Compounds and Structures – An Overview Organic compound: is made up of carbon and hydrogen or carbon, hydrogen and a small number of other elements such as oxygen, nitrogen and sulfur Carbon: • Special because of the ability of carbon atoms to form strong covalent bonds with one another • This allows them to join together into straight chains, branched chains and rings • Nearly infinite number of possible bonding arrangements of C atoms accounts for variety of organic compounds Hydrocarbon: a compound containing the two elements carbon and hydrogen • Can be found in straight or branched chains or ring structures • Simplest organic compounds • The simplest hydrocarbon is methane – CH 4 o Tetrahedral with four equivalent H atoms: equidistant fro the C atom and attached to tit by covalent bonds of equal strength o o Angle between any two C-H bond is 109.28 o Can also be described as four σ bonds formed by the overlap of four hydrogen 1s orbitals with four equivalent sp hybrid orbitals on the carbon atom • By increasing the number of C atoms in the chain, we can obtain still more hydrocarbons Dashed-Wedged line notation: method of conveying a three-dimensional perspective to a structure plotted in a plane • Ordinary lines are used to show bonds that lie in the plane of the paper • Solid wedge lines are used to show bonds that stick out toward the viewer; that is, in front of the plane of the paper • Dashed lines are used to show bonds directed away from the viewer; that is, behind the plane of the paper Isomers: compounds that have the same molecular formula but different structural formulas Skeletal Isomerism: results from differences in the skeletal structures (carbon atom frameworks) of molecules having the same composition • For example butane and isobutene (methylpropane) are isomers where butane is a straight chain and isobutene is a branched chain Positional Isomers: differ in the position on a hydrocarbon chain or ring where a functional group(s) is attached Saturated Hydrocarbons: hydrocarbons with all carbon to carbon bonds as single bonds Nomenclature: • System to name organic compounds developed by IUPAC (International Union of Pure and Applied Chemistry) 1. Select the longest continuous carbon chin in the molecule and use the hydrocarbon name of this chain as the base name a. Methane, ethane, propane, butane, pentane, hexane, heptanes, octane, nonane, decane 2. Consider every branch of the main chain to be a substituent derived from another hydrocarbon and change the ending of its name from –ane to –yl a. Alkane becomes alkyl group and methane becomes methyl group b. Methyl, ethyl, propyl, butyl/butyr, pentyl, hexyl, heptyl... 3. Number the C atoms of the continuous base chain so that the substituents appear at the lowest numbers possible 4. Name each substituent according o its chemical identity and the number of the C atom to which it is attached a. For identical substituents use di, tri, tetra, penta, hexa, hepta, octa, nona,deca 5. Separate numbers from one another by commas, and from letters by hyphens 6. List the substituents alphabetically by name Functional Groups: an atom or grouping of atoms attached to a hydrocarbon residue, R. The functional group often confers specific properties to an organic molecule • The symbol, R refers to the remainder of the molecule, which is usually an alkyl group • In some cases these groupings of atoms are substituted for H atoms in hydrocarbon chains or rings • In other cases they may build from the C atom itself • For example a carbonyl group consists of a C atom in the skeletal structure to which an O atom is attached by a double bond • The physical and chemical properties of organic molecules generally depend on the particular functional groups present • The R of a molecule often has little effect on these properties • Functional groups are a useful way to classify organic compounds with similar properties when studying organic chemistry 26.2 Alkanes Alkane: hydrocarbon molecules that have only single covalent bonds between carbon atoms. In their chain structures, alkanes have the general formula C Hn 2n + 2 • The bonds in these compounds are said to be saturated Alkyl Groups: alkane hydrocarbon molecules from which one hydrogen atom has been extracted (for example a methyl or ethyl groups are examples of alkyl groups) Methylene group: -CH - 2 Homologous Series: substances whose molecules differ only by a constant unit (such as -CH -) 2 • Members of the same homologous series usually have closely related chemical and physical properties • For example straight chain alkane molecules are more easily polarized than their branched-bhain isomers • Chains with greater molecular masses have higher boiling points (due to Van der Waal’s forces) • straight-chain molecules have the strongest intermolecular attractions and thus the highest boiling points o isomers with more compact structures have lower boiling points Conformations: refer to the different spatial arrangements possible in a molecule. Examples are “boat” and “chair” forms of cyclohexane • with a ball and stick model we can visualize an important type of motion in alkane molecules: rotation of groups with respect to one another about the σ connections them • for example in an ethane molecule o eclipsed conformation: when the ethane molecule is viewed head-on one set of C-H bonds is directly behind the other  the distance between the H atoms on the adjacent C atoms is at a minimum leading to a maximum repulsion between the H atoms o staggered conformation: the H atoms are located a maximum distance apart o at room temperature the ethane molecules have enough thermal energy so that the –CH groups 3 can freely rotate about the C-C bond, making all conformations accessible o at lower temperatures ethane occurs mostly in staggered conformation (more stable) o similar situations are encountered in higher alkanes Ring Structures Aliphatic: hydrocarbon molecules that have their carbon atom skeletons arranged in straight or branched chains • have the general formula C Hn 2n + 2 Alicyclic: hydrocarbon molecules that have their carbon atom skeletons arranged in rings and resemble aliphatic (rather than aromatic) hydrocarbons • formed by joining two ends of an aliphatic chain after the elimination of a H atom from each end • have the general formula C Hn 2n • bond angles of clycoalkanes are more strained/stressed than aliphatic alkanes making them more unstable Cyclopropane, Cyclobutane, Cyclopentane, Cyclohexane Cyclopropane is the most reactive, similar to the reactivity of alkenes, and numerous reactions occur in which the ring breaks open to yield a chain molecule Cyclobutane buckles slightly so the four C atoms are not all in the same plane Cyclopentane has one of the C atoms buckled out of the plane of the other four Cyclohexane has numerous conformations, such as the “boat” and “chair” conformations, and flips between these and others fairly easily at room temperature Preparation of Alkanes: • chief sources of alkanes is petroleum • in the presence of a catalyst (such as Pt or Pd) unsaturated hydrocarbons, whether containing double or triple bonds, can be converted to alkanes by the addition of H atoms to the multiple bond systems • halogenated hydrocarbons react with alkali metals to produce alkanes of double the carbon content when under hear and pressure • alkali metal salts of carboxylic acids may be fused with sodium metal hydroxides forming sodium carbonate and an alkane with one less carbon than the carboxylate Reactions of the Alkanes: • saturated hydrocarbons have little affinity for most chemical reactants • non-polar substances that are insoluble in water and unreactive toward acids, bases or oxidizing agents • halogens react only slowly with alkanes at room temperature • At higher temperatures, particularly in the presence of light, halogenations occurs with a substitution reaction: typical of those involving alkane and aromatic hydrocarbons. In such a reaction a functional group replaces an H atom on a chain or ring • This reaction occurs by a chain reaction For example, the Chlorination of Methane: (pg. 1083) Initiation: Propagation: Termination: • Reaction is initiated when some Cl 2olecules absorb sufficient energy to dissociate into Cl atoms • Cl atoms collide with methane to produce methyl free radicals, which combine with Cl mol2cules to form the product CH 3l • When any or all of the last three reactions proceed to the extent of consuming the free radicals present, the reaction stops ⃗heat∨light • The overall equation for the formation of chloromethane is CH + C4 2 CH 3l + HCl • Multiple mixture of products produces • Polyhalogenation can also occur to form dichloromethane (CH Cl , 2et2ylene chloride, a solvent and paint remover) or trichloromethane (CHCl , 3hloroform, a solvent and fumigant) or tetrachloromethane (CCl , carbon tetrachloride, a solvent) 4 Burning of Alkanes: • Most common reaction with oxygen • Oxidation of hydrocarbons underlies their important use as fuels Alkanes from Petroleum: • Lower molecular mass alkanes (methane and ethane) are found principally in natural gas • Propane and butane are found dissolved in petroleum and can be extracted as gases and sold as liquefied petroleum gas (LPG) • Higher alkanes are obtained by fractional distillation of petroleum, a complex mixture of at least 500 compounds • Smooth burning components are preferred because explosive burning results in engine burning • The engine performance is given an octane rating: while 2,2,4-trimethylpentane has a 100, heptanes as 0 for poor performance Refining Petroleum: 1. Fractional Distillation: produces petroleum with octane numbers of near 50-55 and 90 is acceptable for automobiles 2. Thermal Cracking: large hydrocarbon molecules are broken down into molecules in the gasoline range, and the presence of special catalysts promotes the production of branched-chain hydrocarbons 3. Re-forming: isomerization, converts straight-chain to branched-bhain hydrocarbons and other types of hydrocarbons having higher octane numbers 4. Alkylation: In thermal and catalytic cracking, some low-molecular mass hydrocarbons are rejoined into higher molecular mass hydrocarbons • The octane rating can be further improved by adding antiknock compounds to prevent premature combustion such as tetraethhyllead (C H 2 P5 4 lead additives are now replaced by oxygenated hydrocarbons such as methanol and ethanol 26.3 Alkenes and Alkynes Unsaturated hydrocarbons: molecules that contain one or more carbon-to-carbon multiple bonds Alkenes (olefins): hydrocarbons have one or more carbon-to-carbon double bonds in their molecules • Simple alkenes have the general formula C H n 2n Alkyne: hydrocarbons that have one or more carbon-to-carbon triple bonds in their molecules • Simple alkynes have the general formula C H n 2n -2 1. Select as the base chain the longest chain containing the multiple bond 2. Number the C atoms of the chain to place the multiple bond at the lowest possible number 3. Use the ending -ene for alkenes and -yne for alkynes Properties: • Alkenes are similar to alkanes in physical properties • At room temperature, those containing 2 to 4 C atom are gases • Those with 5 to 18 are liquids • Those with more than 18 are solids • Alkynes have higher boiling points than their alkane and alkene counterparts Geometric Isomerism: in organic compounds refers to the existence of non-equivalent structures (cis and trans) that differ in the positioning of substituent groups relative to a double bond • In complexes, the nonequivalent structures are based on the positions at which ligands are attached to the metal center 2 • A double bond between C atoms consists of the overlap of sp hybrid orbitals to form a π bond • Thus rotation about this double bond is severely restricted and the functional are fixed in one place • When the functional groups are on the same side they are called “cis” • When the functional groups are on opposite sides they are called “trans” Stereoisomers: the number and types of atoms and bonds in molecules are the same, but certain atoms are oriented differently in space. Cis and trans isomerism is one type of stereoisomerism • The number and types of atoms and bonds in sterioisomers are the same, but certain atoms are oriented differently in space Optical Isomerism: results from the presence of a chiral atom in a structure, leading to a pair of optical isomers that differ only in the direction that they rotate the plane of polarized light (enantiomers) Preparation and Uses: Elimination reactions: is one in which atoms are removed from adjacent positions on a hydrocarbon chain to produce a small molecule (for example, water) and an additional bond between carbon atoms (from an alkane to an alkene) Preparation of ethylene (pg. 1086) Preparation of Acetylene Addition Reactions: two new groups become bonded to the C atoms of an alkene at the site of the double bond, and the carbon-to-carbon double bond is converted to a single bond Symmetrical Reactants: when the two new groups added are identical Unsymmetrical Reactions: when the two new groups added are different • Difficult to predict the products when the reactants are unsymmetrical Empirical Rule (by Vladimir Markonikov, 1871): When an unsymmetrical reactant (HX, HOH, HOSO H) is 3 added to an unsymmetrical alkene or alkyne, the more positive fragment (usually H) adds to the carbon atom that has the greatest number of attached atoms • The addition of H 2 to a double bond is the reverse of the reaction in which a double bond is formed by the elimination of H2O; the addition reaction is favored in dilute acid, and the elimination reaction is favored in concentration sulfuric acid • Addition reactions of alkynes also follow Markovnikov’s rule • Hydrogen can be added to both alkenes and alkynes in the presence of a metal catalyst such as palladium or platinum called hydrogenation; reaction is only stopped at alkane • If a heterogeneous catalyst known as Lindlar’s catalyst is employed on an alkyne which has a triple bond not on the terminus of the chain then a cis alkene is formed • If sodium in liquid ammonia is used as a hydrogenation agent then a trans alkene is formed 26.4 Aromatic Hydrocarbons Aromatic Hydrocarbons: have ring structures with unsaturation (multiple-bond character) in the carbon-to- carbon bonds in the rings • Most aromatic hydrocarbons are based on the molecule benzene, C H 6 6 Substituted Benzenes: a H from the benzene is substituted with another group (toluene, o-Xylene) Fused Benzenes: rings are fused together (naphthalene, anthracene) Characteristics of Aromatic Hydrocarbons: • Highly flammable and a carcinogen (found in cigarettes, polluated air and as a decomposition product of grease in the charcoal grilling of meat) • Prolonged inhalation of benzene vapor results in a decreased production of both red and white blood cells • They are planar (flat), cyclic molecules • They have a conjugated bonding system: a bonding scheme among the ring atoms that consists of alternating single and double bonds • The ring system must extend throughout the ring and the π electron clouds associated with the double bonds must involve (4n + 2) electrons, where n = 1,2,… • Insoluble in water but soluble in organic solvents • Boiling points of aromatic hydrocarbons are slightly higher than those of the alkanes of similar carbon content o This is because the planar structure and delocalized electron charge density of benzene increases the attractive forces between molecules • Phenyl Group: the resulting species when one of the six equivalent H a
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