1.Structure and Bonding

Ionic bonds in salts form as a result of electrostatic attraction Organic compounds have covalent bonds from ... orbitals and are attracted to nuclei ...

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1. Structure and Bonding

Based on McMurry’s Organic Chemistry, 6th edition, Chapter 1

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 2  10-10 m (200 picometers (pm)) [the unit angstrom (Å) is 10-10 m = 100 pm; 1 pm = 10-12m] 


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


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 


describes where electron most likely to be

 Electron cloud has no specific boundary so we show most

probable area


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


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


p-Orbitals  In each shell there are

three perpendicular p orbitals, px, py, and pz, of equal energy


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

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 (2s2 2p2)

 Stable molecule results at completed shell, octet (eight dots) for

main-group atoms (two for hydrogen)


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


Examples  Oxygen has six valence

electrons (2s2 2p4) but forms two bonds (H2O)

 Nitrogen has five valence electrons

(2s2 2p3) but forms only three bonds (NH3)

 Carbon has four valence electrons (2s2 2p2), forming four bonds (CH4)


Non-bonding electrons  Valence electrons not used in bonding are called nonbonding

electrons, or lone-pair electrons

Nitrogen atom in ammonia (NH3) shares six valence electrons in three covalent bonds and remaining two valence electrons are nonbonding lone pair


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


Bond Energy  Reaction 2 H·  H2 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 = 0.2390 kcal; 1 kcal = 4.184 kJ)


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


Hybridization: sp3 Orbitals and the Structure of Methane  Carbon has 4 valence electrons (2s2 2p2)  But in CH4, all C–H bonds are identical (tetrahedral): how can

we explain this?

 sp3 hybrid orbitals: s orbital and three p orbitals combine (or

hybridize) to form four equivalent, unsymmetrical, tetrahedral atomic orbitals (sppp = sp3), Pauling (1931)


Tetrahedral Structure of Methane  sp3 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.


Hybridization: sp3 Orbitals and the Structure of Ethane  Two C’s bond to each other by s overlap of an sp3 orbital from

each  Three sp3 orbitals on each C overlap with H 1s orbitals to form

six C–H bonds  All bond angles of ethane are tetrahedral


Hybridization: sp2 Orbitals and the Structure of Ethylene  sp2 hybrid orbitals: one 2s orbital combines with two 2p

orbitals, giving 3 orbitals (spp = sp2)  sp2 orbitals are in a plane with 120° angles  Remaining p orbital (unhybridized) is perpendicular to the plane




Structure of Ethylene  Two sp2-hybridized orbitals on each C overlap to form a s bond  p orbitals overlap side-to-side to formation a pi () bond  sp2–sp2 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 sp2 orbitals of C atoms  H–C–H and H–C–C bond angles of about 120°


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


Orbitals of Acetylene  Two sp hybrid orbitals from each C form sp–sp s bond  pz orbitals from each C form a pz–pz  bond by sideways overlap

and py orbitals overlap similarly


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


Hybridization of Nitrogen and Oxygen  Elements other than C can have

hybridized orbitals  H–N–H bond angle in ammonia (NH3) 107.3°  N’s orbitals (sppp) hybridize to form four sp3 orbitals  One sp3 orbital is occupied by two nonbonding electrons, and three sp3 orbitals have one electron each, forming bonds to H


Hybridization of Oxygen in Water  The oxygen atom is sp3-hybridized  Oxygen has six valence-shell electrons but forms only two

covalent bonds, leaving two lone pairs  The H–O–H bond angle is 104.5°


Formal Charges  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  Neutral molecules with both

a “+” and a “-” are dipolar


Formal Charges: examples


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+


The Reaction of HCl with H2O  When HCl gas dissolves in water, a Brønsted acid–base

reaction occurs  HCl donates a proton to water molecule, yielding hydronium ion

(H3O+) 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


Quantitative Measures of Acid Strength  The equilibrium constant (Ke) 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, Ka for HA  Stronger acids have larger Ka  Ka ranges from 1015 for the strongest acids to very small values

(10-60) for the weakest


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  pKa = -log Ka


Acid and Base Strength


Predicting Acid–Base Reactions from pKa Values  A stronger acid (larger Ka) has a smaller pKa and a weaker acid

(smaller Ka) has a larger pKa

 The difference in two pKa 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


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


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


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











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)


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


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 pKa


Lewis Acids  The Lewis definition of acidity includes metal cations, such as

Mg2+  They accept a pair of electrons when they form a bond to a base

 Group 3A elements, such as BF3 and AlCl3, are Lewis acids

because they have unfilled valence orbitals and can accept electron pairs from Lewis bases

 Transition-metal compounds, such as TiCl4, FeCl3, ZnCl2, and

SnCl4, 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

Lewis Acids


Illustration of Curved Arrows Formalism: Lewis Acid-Base Reactions


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


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


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 = 2.5


The Periodic Table and Electronegativity

You can find a very nice, interactive table at: http://www.ptable.com/ 46

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

Electrostatic Potential Maps

 Electrostatic

potential maps show calculated charge distributions  Colors indicate

electron-rich (red) and electron-poor (blue) regions


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 = 3.336  1030 coulomb meter length of an average covalent bond), the dipole moment would be 1.60  1029 Cm, or 4.80 D.


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


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


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.


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.


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

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



stronger Boiling points comparation of alkanes, heters, alcohols, amines (°C) CH3CH2CH3 -42.1

CH3OCH3 -23.7


CH3CH2NH2 16.6

CH3CH2CH2CH3 -0.5

CH3OCH2CH3 10.8

CH3CH2CH2OH 97.4

CH3CH2CH2NH2 47.8

CH3CH2CH2CH2CH3 36.1

CH3CH2OCH2CH3 34.5

CH3CH2CH2CH2OH 117.3

CH3CH2CH2CH2NH2 77.8



 “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.


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 sp3 hybrid orbitals  In double bonds with planar geometry, carbon uses three equivalent sp2 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

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