1. Structure and Bonding. Based on McMurry s Organic Chemistry, 6 th edition, Chapter 1

Transcription

1 1. Structure and Bonding Based on McMurry s Organic Chemistry, 6 th edition, Chapter 1

2 What is an Atom? Structure of an atom Positively charged nucleus (very dense, protons and neutrons) and small (10-15 m in diameter) Negatively charged electrons are in a cloud (10-10 m) around nucleus Diameter is about m (200 picometers (pm)) [the unit angstrom (Å) is m = 100 pm; 1 pm = m] 2

3 Atomic Number and Atomic Mass The atomic number (Z) is the number of protons in the atom's nucleus The mass number (A) is the number of protons plus neutrons All the atoms of a given element have the same atomic number Isotopes are atoms of the same element that have different numbers of neutrons and therefore different mass numbers The atomic mass (atomic weight) of an element is the weighted average mass in atomic mass units (amu) of an element s naturally occurring isotopes 3

4 How are the electrons distributed in atoms? Quantum mechanics: describes electron energies and locations by a wave equation Wave function solution of wave equation Each Wave function is an orbital, A plot of 2 describes where electron most likely to be Electron cloud has no specific boundary so we show most probable area 4

5 What shapes do orbitals have? Four different kinds of orbitals for electrons Denoted s, p, d, and f s and p orbitals most important in organic chemistry s orbitals: spherical, nucleus at center p orbitals: dumbbell-shaped, nucleus at middle 5

6 Orbitals and Shells Orbitals are grouped in shells of increasing size and energy Different shells contain different numbers and kinds of orbitals Each orbital can be occupied by two electrons The first shell contains one s orbital, denoted 1s, holds only two electrons The second shell contains one s orbital (2s) and three p orbitals (2p), eight electrons The third shell contains an s orbital (3s), three p orbitals (3p), and five d orbitals (3d), 18 electrons 6

7 p-orbitals In each shell there are three perpendicular p orbitals, p x, p y, and p z, of equal energy 7

8 Atomic Structure: Electron Configurations Ground-state electron configuration of an atom lists orbitals occupied by its electrons. Rules: 1. Lowest-energy orbitals fill first: 1s 2s 2p 3s 3p 4s 3d (Aufbau ( build-up ) principle) 2. Electron spin can have only two orientations, up and down. Only two electrons can occupy an orbital, and they must be of opposite spin (Pauli exclusion principle) to have unique wave equations 3. If two or more empty orbitals of equal energy are available, electrons occupy each with spins parallel until all orbitals have one electron (Hund's rule). 8

9 The Nature of the Chemical Bonds Atoms form bonds because the compound that results is more stable (has less energy) than the separate atoms Ionic bonds in salts form as a result of electrostatic attraction Organic compounds have covalent bonds from sharing electrons (G. N. Lewis, 1916) Lewis structures show valence electrons of an atom as dots Hydrogen has one dot, representing its 1s electron Carbon has four dots (2s 2 2p 2 ) Stable molecule results at completed shell, octet (eight dots) for main-group atoms (two for hydrogen) 9

10 How many Covalent Bonds does an atom form? Atoms with one, two, or three valence electrons form one, two, or three bonds (because that is the number of electrons they have available to share in bond formation) Atoms with four or more valence electrons form as many bonds as they need electrons to fill the s and p levels of their valence shells to reach a stable octet 10

11 Examples Oxygen has six valence electrons (2s 2 2p 4 ) but forms two bonds (H 2 O) Nitrogen has five valence electrons (2s 2 2p 3 ) but forms only three bonds (NH 3 ) Carbon has four valence electrons (2s 2 2p 2 ), forming four bonds (CH 4 ) 11

12 Non-bonding electrons Valence electrons not used in bonding are called nonbonding electrons, or lone-pair electrons Nitrogen atom in ammonia (NH 3 ) shares six valence electrons in three covalent bonds and remaining two valence electrons are nonbonding lone pair 12

13 Valence Bond Theory Covalent bond forms when two atoms approach each other closely so that a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom Electrons are paired in the overlapping orbitals and are attracted to nuclei of both atoms H H bond results from the overlap of two singly occupied hydrogen 1s orbital H-H bond is cylindrically symmetrical, sigma (s) bond 13

14 Bond Energy Reaction 2 H H 2 releases 436 kj/mol Product has 436 kj/mol less energy than two atoms: H H has bond strength of 436 kj/mol. (1 kj = kcal; 1 kcal = kj) 14

15 Bond Length Distance between nuclei that leads to maximum stability If too close, they repel because both are positively charged If too far apart, bonding is weak 15

16 Hybridization: sp 3 Orbitals and the Structure of Methane Carbon has 4 valence electrons (2s 2 2p 2 ) But in CH 4, all C H bonds are identical (tetrahedral): how can we explain this? sp 3 hybrid orbitals: s orbital and three p orbitals combine (or hybridize) to form four equivalent, unsymmetrical, tetrahedral atomic orbitals (sppp = sp 3 ), Pauling (1931) 16

17 Tetrahedral Structure of Methane sp 3 orbitals on C overlap with 1s orbitals on 4 H atoms to form four identical C-H bonds (sigma bonds) Bond angle: each H C H is 109.5, the tetrahedral angle. 17

18 Hybridization: sp 3 Orbitals and the Structure of Ethane Two C s bond to each other by s overlap of an sp 3 orbital from each Three sp 3 orbitals on each C overlap with H 1s orbitals to form six C H bonds All bond angles of ethane are tetrahedral 18

19 Hybridization: sp 2 Orbitals and the Structure of Ethylene sp 2 hybrid orbitals: one 2s orbital combines with two 2p orbitals, giving 3 orbitals (spp = sp 2 ) sp 2 orbitals are in a plane with 120 angles Remaining p orbital (unhybridized) is perpendicular to the plane

20 Structure of Ethylene Two sp 2 -hybridized orbitals on each C overlap to form a s bond p orbitals overlap side-to-side to formation a pi ( ) bond sp 2 sp 2 s bond and 2p 2p bond result in sharing four electrons and formation of C=C double bond H atoms form s bonds with four sp 2 orbitals of C atoms H C H and H C C bond angles of about

21 Hybridization: sp Orbitals and the Structure of Acetylene C-C a triple bond sharing six electrons Carbon 2s orbital hybridizes with a single 2p orbital giving two sp hybrids two p orbitals remain unchanged sp orbitals are linear, 180 apart on x-axis Two p orbitals are perpendicular on the y-axis and the z-axis 21

22 Orbitals of Acetylene Two sp hybrid orbitals from each C form sp sp s bond p z orbitals from each C form a p z p z bond by sideways overlap and p y orbitals overlap similarly 22

23 Bonding in Acetylene Sharing of six electrons forms C C Two sp orbitals form s bonds with hydrogens 23

24 Hybridization of Nitrogen and Oxygen Elements other than C can have hybridized orbitals H N H bond angle in ammonia (NH 3 ) N s orbitals (sppp) hybridize to form four sp 3 orbitals One sp 3 orbital is occupied by two nonbonding electrons, and three sp 3 orbitals have one electron each, forming bonds to H 24

25 Hybridization of Oxygen in Water The oxygen atom is sp 3 -hybridized Oxygen has six valence-shell electrons but forms only two covalent bonds, leaving two lone pairs The H O H bond angle is

26 Sometimes it is necessary to have structures with formal charges on individual atoms We compare the bonding of the atom in the molecule to the valence electron structure If the atom has one more electron in the molecule, it is shown with a - charge If the atom has one less electron, it is shown with a + charge Formal Charges Neutral molecules with both a + and a - are dipolar 26

27 Formal Charges: examples 27

28 Acids and Bases: The Brønsted Lowry Definition The terms acid and base can have different meanings in different contexts The idea that acids are solutions containing a lot of H + and bases are solutions containing a lot of OH - is not very useful in organic chemistry Instead, Brønsted Lowry theory defines acids and bases by their role in reactions that transfer protons (H + ) between donors and acceptors A Brønsted acid is a substance that donates a hydrogen ion (H + ) A Brønsted base is a substance that accepts the H + 28

29 The Reaction of HCl with H 2 O When HCl gas dissolves in water, a Brønsted acid base reaction occurs HCl donates a proton to water molecule, yielding hydronium ion (H 3 O + ) and Cl The reverse is also a Brønsted acid base reaction of the conjugate acid and conjugate base Acids are shown in red, bases in blue. Curved arrows go from bases to acids 29

30 Quantitative Measures of Acid Strength The equilibrium constant (K e ) for the reaction of an acid (HA) with water to form hydronium ion and the conjugate base (A - ) is a measure related to the strength of the acid The acidity constant, K a for HA Stronger acids have larger K a K a ranges from for the strongest acids to very small values (10-60 ) for the weakest 30

31 Acid and Base Strength The ability of a Brønsted acid to donate a proton is sometimes referred to as the strength of the acid (imagine that it is throwing the proton stronger acids throw it harder) The strength of the acid is measured with respect to the Brønsted base that receives the proton Water is used as a common base for the purpose of creating a scale of Brønsted acid strength pk a = -log K a 31

32 Acid and Base Strength 32

33 Predicting Acid Base Reactions from pk a Values A stronger acid (larger K a ) has a smaller pk a and a weaker acid (smaller K a ) has a larger pk a The difference in two pk a values can be used to calculate the extent of transfer A proton always goes from the stronger acid to the stronger base The stronger base holds the proton more tightly 33

34 1- Electronegativity Predicting Acid Base Stength An anion is stabilized by having the negative charge on a highly electronegative atom. Higher electronegativity = lower basicity = higher acidity of the conjugated acid. 34

35 Predicting Acid Base Stength 2-Orbital surface Increasing the orbital surface decreases the electron density and, as a consequence, the basicity. 35

36 Predicting Acid Base Stength 3- Resonance effect The conjugated base can be stabilized by resonance effect. If electrons are delocalized the basicity decreases, so the acidity of the conjugated acid increases. -H CH CH 3 CH 3 CH 2 OH 2 O -H O O CH CH 3 C 3 COH O O CH 3 C O O CH 3 C O -H O O H 3 HC 3 CS SOH O O O O H 3 C S O O O H 3 C S O O O H 3 C S O O O H 3 C S O O OH > CH 3 C OH > CH 3 CH 2 OH 36

37 Organic Acids Those that lose a proton from O H, such as methanol and acetic acid Those that lose a proton from C H, usually from a carbon atom next to a C=O double bond (O=C C H) 37

38 Organic Bases Have an atom with a lone pair of electrons that can bond to H + Nitrogen-containing compounds derived from ammonia are the most common organic bases Oxygen-containing compounds can react as bases with a strong acid or as acids with strong bases 38

39 Acids and Bases: The Lewis Definition Lewis acids are electron pair acceptors and Lewis bases are electron pair donors Lewis acid include not only proton donors but many other species The Lewis definition leads to a general description of many reaction patterns but there is no scale of strengths as in the Brønsted definition of pk a 39

40 Lewis Acids The Lewis definition of acidity includes metal cations, such as Mg 2+ They accept a pair of electrons when they form a bond to a base Group 3A elements, such as BF 3 and AlCl 3, are Lewis acids because they have unfilled valence orbitals and can accept electron pairs from Lewis bases Transition-metal compounds, such as TiCl 4, FeCl 3, ZnCl 2, and SnCl 4, are Lewis acids Organic compounds that undergo addition reactions with Lewis bases are called electrophiles and therefore Lewis Acids The combination of a Lewis acid and a Lewis base can be shown with a curved arrow from base to acid 40

41 Lewis Acids 41

42 Illustration of Curved Arrows Formalism: Lewis Acid-Base Reactions 42

43 Lewis Bases Lewis bases can accept protons as well as Brønsted bases, therefore the definition encompasses that for Brønsted base Most oxygen- and nitrogen-containing organic compounds are Lewis bases because they have lone pairs of electrons Some compounds can act as both acids and bases, depending on the reaction 43

44 Polar Covalent Bonds: Electronegativity Covalent bonds can have ionic character These are polar covalent bonds Bonding electrons attracted more strongly by one atom than by the other Electron distribution between atoms in not symmetrical 44

45 Bond Polarity and Electronegativity Electronegativity (EN): intrinsic ability of an atom to attract the shared electrons in a covalent bond Differences in EN produce bond polarity Arbitrary scale. Electronegativities are based on an arbitrary scale F is most electronegative (EN = 4.0), Cs is least (EN = 0.7) Metals on left side of periodic table attract electrons weakly, lower EN Halogens and other reactive nonmetals on right side of periodic table attract electrons strongly, higher electronegativities EN of C =

46 The Periodic Table and Electronegativity You can find a very nice, interactive table at: 46

47 Bond Polarity and Inductive Effect Nonpolar Covalent Bonds: atoms with similar EN Polar Covalent Bonds: Difference in EN of atoms < 2 Ionic Bonds: Difference in EN > 2 C H bonds, relatively nonpolar C-O, C-X bonds (more electronegative elements) are polar Bonding electrons toward electronegative atom C acquires partial positive charge, + Electronegative atom acquires partial negative charge, - Inductive effect: shifting of electrons in a bond in response to EN of nearby atoms 47

48 Electrostatic Potential Maps Electrostatic potential maps show calculated charge distributions Colors indicate electron-rich (red) and electron-poor (blue) regions 48

49 Polar Covalent Bonds: Dipole Moments Molecules as a whole are often polar :from vector summation of individual bond polarities and lone-pair contributions Strongly polar substances soluble in polar solvents like water; nonpolar substances are insoluble in water. Dipole moment - Net molecular polarity, due to difference in summed charges - magnitude of charge Q at end of molecular dipole times distance r between charges = Q r, in debyes (D), 1 D = coulomb meter length of an average covalent bond), the dipole moment would be C m, or 4.80 D. 49

50 Absence of Dipole Moments In symmetrical molecules, the dipole moments of each bond has one in the opposite direction The effects of the local dipoles cancel each other 50

51 Intermolecular interactions The strength of attractions between molecules regulates the physical properties, such as melting point, boiling point and solubility. There are various forms of interactions. Dipole-dipole interaction (depends on polarity) Van der Waals Interactions (or London forces) Hydrogen bonds 51

52 Dipole-dipole Interactions Interaction between polar molecules The positively charged side of a molecule attracts the negatively charged side of a another molecule. Stronger the dipole-dipole interactions, higher the melting and boiling temperatures. 52

53 Van der Waals interactions Always present, also in non-polar molecules Derive from temporary dipole Bigger atoms polarize easily The strength of van der Waals forces depends on the contact surface. Branched molecules have weaker interactions. 53

54 Hydrogen bond It is a special kind of dipole-dipole interaction, very strong. It occurs between a Hydrogen bonded to an electronegative atom (O, N, Halogen), and an electron pair of the electronegative atom of another molecule. Hydrogen bond in water 54

55 Hydrogen bond In organic molecules the hydrogen bond occurs when N-H (amines, amides), or O-H (alcohols, carboxilic acids) are present. O-H is more polar than N-H, so the hydrogen bond is stronger Boiling points comparation of alkanes, heters, alcohols, amines ( C) CH 3 CH 2 CH CH 3 CH 2 CH 2 CH CH 3 CH 2 CH 2 CH 2 CH CH 3 OCH CH 3 OCH 2 CH CH 3 CH 2 OCH 2 CH CH 3 CH 2 OH 78 CH 3 CH 2 CH 2 OH 97.4 CH 3 CH 2 CH 2 CH 2 OH CH 3 CH 2 NH CH 3 CH 2 CH 2 NH CH 3 CH 2 CH 2 CH 2 NH

56 Solubility Similia similibus solvuntur Polar molecules are soluble in polar solvents: solvatation. Non-polar molecules are soluble in nonpolar solvents. Molecules with similar intermolecular interactions are miscible. 56

57 Summary Atom: positively charged nucleus surrounded by negatively charged electrons Electronic structure of an atom described by wave equation Electrons occupy orbitals around the nucleus. Different orbitals have different energy levels and different shapes Covalent bonds - electron pair is shared between atoms Valence bond theory - electron sharing occurs by overlap of two atomic orbitals Sigma (s) bonds - Circular cross-section and are formed by head-on interaction Pi ( ) bonds dumbbell shape from sideways interaction of p orbitals Carbon uses hybrid orbitals to form bonds in organic molecules. In single bonds with tetrahedral geometry, carbon has four sp 3 hybrid orbitals In double bonds with planar geometry, carbon uses three equivalent sp 2 hybrid orbitals and one unhybridized p orbital Carbon uses two equivalent sp hybrid orbitals to form a triple bond with linear geometry, with two unhybridized p orbitals 57

58 Summary Organic molecules often have polar covalent bonds as a result of unsymmetrical electron sharing caused by differences in the electronegativity of atoms (+) and ( ) indicate formal charges on atoms in molecules to keep track of valence electrons around an atom A Brønsted( Lowry) acid donates a proton A Brønsted( Lowry) base accepts a proton The strength Brønsted acid is related to the -1 times the logarithm of the acidity constant, pka. Weaker acids have higher pka s A Lewis acid has an empty orbital that can accept an electron pair A Lewis base can donate an unshared electron pair (In condensed structures C-C and C-H are implied Skeletal structures show bonds and not C or H (C is shown as a junction of two lines) other atoms are shown Molecular models are useful for representing structures for study Some substances must be shown as a resonance hybrid of two or more resonance forms that differ by the location of electrons.) 58