Structure of rganic Molecules Ref. books: 1. A text book of rganic Chemistry - B.S. Bahl and Arun Bahl 2. rganic Chemistry - R.T. Morrison and R. N. Boyd Atom: The smallest part of an element that can exist chemically. Atoms consist of a small dense nucleus of protons and neutrons surrounded by moving electrons. No. of electrons = no. of protons. So the overall charge is zero. The electrons are considered to move in circular or elliptical orbits. Maximum no. of electrons in orbitals: 2, 8, 18, 32, (2n 2 ) utermost orbit of electrons is incomplete (except inert gases), which are known as valence electrons. Atoms combine to form molecule with the urge of atoms to complete their outermost orbits of electrons as in the inert gases. There are three basic ways in which chemical combination occurs: 1. Ionic or electrovalent bond 2. Covalent bond 3. Coordinate bond Ionic bond Ionic bond is a type of chemical bond that involves the electrostatic attraction between oppositely charged ions. Neutral atoms come near each other. Electron(s) are transferred from the metal atom to the non-metal atom. They stick together because of electrostatic forces, like magnets. In an ionic bond, electrons are lost or gained, resulting in the formation of ions in ionic compounds. K F K + F The compound potassium fluoride consists of potassium (K+) ions and fluoride (F-) ions The ionic bond is the attraction between the positive K + ion and the negative F - ion _ 1
CVALENT BND FRMATIN When one nonmetal shares one or more electrons with an atom of another nonmetal so both atoms end up with eight valence electrons Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons Both end with full orbitals 8 Valence electrons F F 8 Valence electrons Water is formed with covalent bonds Water H Each hydrogen has 1 valence electron Each hydrogen wants 1 more The oxygen has 6 valence electrons The oxygen wants 2 more They share to make each other happy Put the pieces together The first hydrogen is happy The oxygen still wants one more H 2
Water So, a second hydrogen attaches Every atom has full energy levels H H Note the two unshared pairs of electrons Another way of indicating bonds ften use a line to indicate a bond Called a structural formula Each line has 2 valence electrons H H = H H A single covalent bond Sharing of two valence electrons. nly nonmetals and hydrogen. Multiple bonds Sometimes atoms share more than one pair of valence electrons. A double bond is fromed when atoms share two pair (4) of electrons. A triple bond is formed when atoms share three pair (6) of electrons. C Carbon dioxide C 2 - Carbon is central atom Carbon has 4 valence electrons Wants 4 more xygen has 6 valence electrons Wants 2 more 3
Carbon dioxide Attaching 1 oxygen leaves the oxygen 1 short and the carbon 3 short Carbon dioxide Attaching the second oxygen leaves both oxygen 1 short and the carbon 2 short C C Carbon dioxide A coordinate covalent bond The only solution is to share more Requires two double bonds Each atom can count all the atoms in the bond When one atom donates both electrons in a covalent bond. Carbon monoxide (C) is a good example: C Both the carbon and oxygen give another single electron to share C 4
Coordinate bovalent bond Coordinate covalent bond When one atom donates both electrons in a covalent bond. Carbon monoxide (C) is a good example: This carbon electron moves to make a pair with the other single. C xygen gives both of these electrons, since it has no more singles to share. When one atom donates both electrons in a covalent bond. Carbon monoxide (C) The coordinate covalent bond is shown with an arrow as: C C Covalent Bonding Covalent bonding results when two electrons are shared in an orbital between two atoms The pair of electrons used are called the shared or bonding pair. The electron pairs that are not involved in bonding are called lone pairs. BND RDER - When only one pair of electrons are shared between two atoms, it s called a single bond. If two pairs of electrons are shared covalently between two atoms, it s called a double bond; three pairs, triple bond. Sigma Bonding Electron density lies between the nuclei. A bond may be formed by s-s, p-p, s-p, or hybridized orbital overlaps. The bonding M is lower in energy than the original atomic orbitals. The antibonding M is higher in energy than the atomic orbitals. 5
H 2 : s-s overlap s-p overlap verlap of an s orbital with a p orbital also gives a bonding M and an antibonding M p-p overlap Sigma bonds occur when the orbitals of two shared electrons overlap head-to-head. Pi Bonding Pi bonds form after sigma bonds. Sideways overlap of parallel p orbitals. 6
Formation of covalent bonding Multiple Bonds A double bond (2 pairs of shared electrons) consists of a sigma bond and a pi bond. A triple bond (3 pairs of shared electrons) consists of a sigma bond and two pi bonds. Unsaturated Hydrocarbons Electronegativity Electronegativity is the power of an atom to attract electron density in a covalent bond. Pauling s electronegativity scale: The higher the value, the more electronegative the element. Fluorine is the most electronegative element. It has an electronegativity value of 4.0 Caesium is the least electronegative element 0.79 7
The Pauling electronegativity (EN) scale A nonpolar covalent bond occurs between nonmetals. is an equal or almost equal sharing of electrons. has almost no electronegativity difference (0.0 to 0.4). Atoms EN Difference Type of Bond N-N 3.0-3.0 = 0.0 Pure covalent C-H 2.5-2.1 = 0.4 Nonpolar covalent The elements whose natural state is diatomic: Hydrogen (H 2 ), Nitrogen (N 2 ), xygen ( 2 ), Fluorine (F 2 ), Chlorine (Cl 2 ), Bromine (Br 2 ), and Iodine (I 2 ) the electrons are shared equally. This type of bond is a pure covalent bond (equal sharing) A polar covalent bond occurs between nonmetal atoms. is an unequal sharing of electrons. has a moderate electronegativity difference (0.5 to 2.0). Atoms EN Difference Type of Bond Cl-C 3.0-2.5 = 0.5 Polar covalent H-Cl 3.0-2.1 = 0.9 Polar covalent An ionic bond occurs between metal and nonmetal ions. is a result of electron transfer. has a large electronegativity difference (2.0 or more). Atoms EN Difference Type of Bond N-Na 3.0 0.9 = 2.1 Ionic (non-covalent) LiF 4.0 1.0 = 3.0 Ionic (non-covalent) Equal sharing of electrons in a nonpolar covalent bond 8
Unequal Sharing in Polar Covalent Bonds The negative pole is centered on the more electronegative atom in the bond. This atom has a share in an extra electron. The positive pole is centered on the less electronegative atom. This atom has lost a share in one of its electrons. Because there was not a complete transfer of an electron, the charges on the poles are not 1+ and 1, but δ + and δ. Electronegativity and bond types Attraction and repulsion in covalent bond Formation of a covalent bond A molecular orbital is the region of high probability that is occupied by an individual electron as it travels with a wavelike motion in the 3D space around one of two or more associated nuclei. 9
Spatial orientation of orbitals is designated by subscripts for the p, d and f orbitals. s orbitals require no designation since there is only one possible s orbital. Three p orbitals, five d orbitals and seven f orbitals. Therefore in the cases of p, d and f orbitals each one requires specific identification in order to differentiate one orbital of the same type from another. S orbital shape 1S p orbital shapes d orbital shapes dxz dx 2 -y 2 P X P Z P Y dyz dz 2 dxz 10
Hybridization Electron configuration of carbon In 1931, Linus Pauling proposed that the wave functions for the s and p atomic orbitals can be mathematically combined to form a new set of equivalent wave functions called hybrid orbitals. In a hybridization scheme: o No. of hybrid orbitals = total no. of atomic orbitals o The symbols identify the numbers and kinds orbitals involved. 2p 2s only two unpaired electrons should form s bonds to only two hydrogen atoms sp 3 orbital hybridization sp 3 orbital hybridization 2p 2p 2p Promote an electron from the 2s to the 2p orbital 2s 2s 2 sp 3 Mix together (hybridize) the 2s orbital and the three 2p orbitals 4 equivalent half-filled orbitals are consistent with four bonds and tetrahedral geometry 2s 11
Characteristics of hybridization The formation of four sp 3 hybrid orbitals by combination of an atomic s orbital with three atomic p orbitals. Each sp 3 hybrid orbital has two lobes, one of which is larger than the other. The four large lobes are oriented toward the corners of a tetrahedron at angles of 109.5. Hybrids shown together Four tetrahedral sp 3 hybrid orbitals 1. nly orbitals of almost similar energies and belonging to the same atom or ion undergoes hybridization. 2. Hybridization takes place only in orbitals, electrons are not involved in it. 3. The number of hybrid orbitals produced is equal to the number of pure orbitals, mixed during hybridization. 4. Both half filled orbitals or fully filled orbitals of equivalent energy can involve in hybridization. 5. Hybrid orbitals form only sigma bonds. Characteristics of hybridization 6. rbitals involved in π bond formation do not participate in hybridization. 7. Hybridization never takes place in an isolated atom but it occurs only at the time of bond formation. 8. The hybrid orbitals are distributed in space as apart as possible resulting in a definite geometry of molecule. 9. Hybridized orbitals provide efficient overlapping than overlapping by pure s, p and d-orbitals. 10. Hybridized orbitals possess lower energy. Structure of Ethylene C 2 H 4 H 2 C=CH 2 planar bond angles: close to 120 bond distances: C H = 110 pm C=C = 134 pm 12
sp 2 orbital hybridization sp 2 orbital hybridization 2p 2p 2p 2 sp 2 Promote an electron from the 2s to the 2p orbital 2s 2s Mix together (hybridize) the 2s orbital and two of the three 2p orbitals 3 equivalent half-filled sp 2 hybrid orbitals plus 1 p orbital left unhybridized 2 of the 3 sp 2 orbitals are involved in s bonds to hydrogens; the other is involved in a s bond to carbon 2s Structure of Acetylene C 2 H 2 HC CH linear bond angles: 180 sp 2 and pi orbital bond distances: C H = 106 pm CC = 120 pm 13
sp orbital hybridization sp orbital hybridization 2 p 2p 2p 2p Promote an electron from the 2s to the 2p orbital 2s 2s 2 sp Mix together (hybridize) the 2s orbital and one of the three 2p orbitals 2 equivalent half-filled sp hybrid orbitals plus 2 p orbitals left unhybridized 1 of the 2 sp orbitals is involved in a s bond to hydrogen; the other is involved in a s bond to carbon 2s NH 3 sp 3 orbital hybridization 14
Bond Lengths The distance between the nuclei of bonded atoms is called the bond length Because the actual bond length depends on the other atoms around the bond we often use the average bond length In general, the larger the atoms involved in a bond, the longer the bond length, multiple bonds result in stronger and shorter bonds. averaged for similar bonds from many compounds, Atomic size increases going down a group Bond length: S - Br > S - Cl > S - F Bond strength: S - F > S - Cl > S - Br Bond Energy Chemical reactions involve breaking bonds in reactant molecules and making new bonds to create the products. Bond energy is the energy required to break the bond(s) between two atoms. In general, the shorter the bond, the higher the bond energy. Using bond orders we get Bond length: C - > C = > C Bond strength: C > C = > C - 15
It is common practice that tabulated values of bond energy are termed as bond enthalpy Bond breaking is an endothermic process, so bond breaking enthalpies are positive. Bond angle A bond angle is an angle that is formed between three atoms across two bonds. 109 28 The overall shape of a molecule is determined by its bond angles. The angles made by the lines joining the nuclei of the atoms in the molecule. The bond angles of a molecule, together with the bond lengths accurately define the shape and size of the molecule. Factors affecting bond angle Lone pairs occupy more space than bonding electron pairs. Double bonds occupy more space than single bonds. LP-LP > LP-BP > BP-BP More electronegative attached atoms bond angle Cl.. P Cl Cl < I.. P I I decrease in Lone pairs are more repulsive than bonding pairs LP-BP repulsion H.. N H H 107o 31' decrease in bond angle > H 105 o 31' H More electronegative (central atom) H 105 o 31' > H H : S : 92 0 H increase in bond angle 16
Intermolecular/Intramolecular Forces Intermolecular forces are attractive forces between molecules. Intramolecular forces hold atoms together in a molecule (covalent bond). Generally, intermolecular forces are much weaker than intramolecular forces. Intermolecular vs Intramolecular 41 kj to vaporize 1 mole of water (inter) 830 kj to break all -H bonds in 1 mole of water (intra) Intermolecular Forces Strength of attractions between molecules. Generally much weaker than covalent or ionic bonds. Influence m.p., b.p., and solubility; esp. for solids and liquids. Classification depends on structure. Dipole-dipole interactions London dispersions Hydrogen bonding Dipole-dipole interactions Between polar molecules Positive end of one molecule aligns with negative end of another molecule. Repulsions less, net force is attractive. Larger dipoles cause higher boiling points and higher heats of vaporization. 17
Dipole-dipole interactions attraction (common) Dispersion (London) forces Due to electron repulsion, a temporary dipole on one atom can induce a similar dipole on a neighboring atom, significant only when molecules are close to each other a) Normal Condition: A non-polar molecule has a symmetrical charge distribution b) Instantaneous Condition: A displacement of the electronic charge produces an instantaneous dipole with a charge separation represented as + and -. c) Induced Dipole: The instantaneous dipole on the left induces a charge separation in the molecule on the right. The result is a dipoledipole interaction. The two dipoles, in the two molecules, will attract each other, and the result is that the potential energy of the two is lowered. 18
Dispersion (London) forces Between nonpolar molecules Temporary dipole-dipole interactions Larger atoms are more polarizable. Branching lowers b.p. because of decreased surface contact between molecules. Dispersions CH 3 CH 2 CH 2 CH 2 CH 3 n-pentane, b.p. = 36 C CH 3 CH 3 CH CH 2 CH 3 isopentane, b.p. = 28 C H 3 C CH 3 C CH 3 CH 3 neopentane, b.p. = 10 C Hydrogen bonding A chemical bond in which a hydrogen atom of one molecule is attracted to an electronegative atom, especially a nitrogen, oxygen, or fluorine atom, usually of another molecule. rganic molecule having N-H or -H. The hydrogen from one molecule is strongly attracted to a lone pair of electrons on the other molecule. -H more polar than N-H, so stronger hydrogen bonding Hydrogen Bonding + - H H + 19
Dipole moments A dipole moment is simply the measure of net polarity in a molecule. The bond dipole moment is a measure of polarity of a chemical bond within a molecule. Bond dipole moments Molecular dipole moments Dipole moments are due to differences in electronegativity. They depend on the amount of charge and distance of separation (μ=q r) They are measured in debyes (D). The molecular dipole moment is the vector sum of the bond dipole moments. Depend on the bond polarity and bond angles. Lone pairs of electrons contribute to the dipole moment. Bond Dissociation Energy(D) (also bond strength) is how much energy it takes to break a specific bond into two radical fragments when the molecules is in the gas phase at 25 C. Bond dissociation energy It is characteristic of the particular bond Example: Methane, four carbon-hydrogen bonds with four different bond dissociation energies D(CH 3 H) = 104 kcal/mol D(CH 2 H) = 106 kcal/mol D(CH H) = 106 kcal/mol D(C H) = 81 kcal/mol = 397 kcal/mol Bond energy (E) is an average measured over many similar bonds in different molecules. CH 4 E(C H) =397/4 = 99kcal/mol In the case of diatomic molecules, bond energy (E) and bond dissociation energy (D) is same 20
Homolytic cleavage (Homolysis): When the bond breaks, each atom gets one electron. Bond formation occurs in two different ways: Two radicals can each donate one electron to form a two-electron bond. Alternatively, two ions with unlike charges can come together, with the negatively charged ion donating both electrons to form the resulting two-electron bond. Heterolytic cleavage (Heterolysis) : When the bond breaks, the most electronegative atom gets both electrons. Bond formation always releases energy. Resonance is the process whereby (generally) p-electrons can be delocalised by exchanging double bonds and single bonds. Resonance can be used to delocalise both lone pairs of electrons and cationic charges which are adjacent to double bonds. Delocalisation of positive and negative charges lead to relatively stable cations and anions, respectively. Benzene Bond Lengths Resonance structures of benzene The 6 p-electrons are able to flow (or resonate) continually around the p-molecular orbital formed from the six p atomic orbitals on each of the 6 carbon atoms on the ring structure. This is represented by the two resonance structures below (which are identical or degenerate). Canonical structurtes The p-electrons are referred to as being conjugated. 21
Resonance imparts stability to anionic and cationic structures H H H 3 C CH 2 Replaced by C= Relatively difficult to form H 3 C Relatively easy to form No adjacent double bond to the oxygen lone pair H 3 C The ability to be able to delocalise (spread out) charge via resonance allows an assessment of (i) the degree of ease of formation of the charged species, and (ii) the stability of the charged species H 3 C H 3 C The resonance arrow is not an equilibrium arrow The resonance arrow shows only the distribution of electrons. H 3 C Experimentally it is found that both C- bonds are the same length and are intermediate in length between the C- single and double bond, as are the C-C bonds in benzene. General structure that will display resonance of charges and lone pairs of electrons R 1 R 3 R 2 R 4 R 1 R 3 R 2 R 4 Me Me Me Me Note in a reaction mechanism we would not show the lone pairs on the carbons carrying the ve charge. 22