PF1009 2024 6 Bond and Molecular Polarity PDF

Pharmaceutical Chemistry Bond and Molecular Polarity Dr. J.J. Keating 1  Electronegativity A measure of the tendency of an element to attract a pair of electrons. Pauling scale...

Pharmaceutical Chemistry Bond and Molecular Polarity Dr. J.J. Keating 1  Electronegativity A measure of the tendency of an element to attract a pair of electrons. Pauling scale 2  Bonding A chemical bond forms if, as a result, the energy of the bonded atoms is lower than that of the separate atoms. Atoms tend to arrange themselves in the most stable patterns possible, which means that they have a tendency to complete or fill their outermost valence shells. Need to take into account the electronegativity of atoms when forming bonds. Many different types of bonding occur: Nonpolar covalent Polar covalent Ionic Hydrogen-bonding Coordinate bonds Metallic bonds Single, double, triple bonds σ-bonds π-bonds increasing difference in electronegativity 3  Nonpolar Covalent Bonds Covalent Bonds Atoms share an electron pair, with no or charges or vary small partial charges on atoms. Give rise to molecular compounds. Solids, liquids or gases at normal temperatures and pressures. Generally formed between elements that have similar electronegativities. Nonpolar covalent bonds – difference in electronegativity < 0.5 3.16 3.16 H 2.20 H2.20 H C2.55 C 2.55 H C H H H H C2H4 Cl2 CH4 ethene chlorine molecule methane (C-H) 2.55 – 2.20 = 0.25 (< 0.5) (Cl-Cl) 3.16 – 3.16 = 0.00 (< 0.5) (C-H) 2.55 – 2.20 = 0.25 (< 0.5) (C-C) 2.55 – 2.55 = 0.00 ( 0.5 but < 2.0. Give rise to polar bonds. delta negative Solids, liquids or gases at normal temperatures and pressures. Bonding electrons shared unequally between two atoms. δ+ delta positive Partial charges on atoms. 2.20 δ+ 3.44 δ– O 3.44 δ + δ– C 2.55 2.20 2.20 3.16 H3C CH3 δ+ 2.20 C3H4O H2O HCl acetone (propan-2-one) water hydrogen chloride (C-H) 2.55 – 2.20 = 0.25 (< 0.5) (O-H) 3.44 – 2.20 = 1.22 (Cl-H) 3.16 – 2.20 = 0.96 (O-C) 3.44 – 2.55 = 0.89 5 (> 0.5, < 2) (> 0.5, < 2) (> 0.5, < 2)  Ionic Bonds Complete transfer of one or more valence electrons from one atom to another. Full charges on resulting ions with atomic difference in electronegativity > 2. Can have high aqueous solubility. Attraction between cations and anions. Give rise to ionic compounds. Solids at normal temperatures and pressures. Generally formed between elements that differ strongly in their electronegativity. Ionic compounds are formed between elements towards the left of the periodic table and elements on the right of the table. Ionic compounds typically form between the metallic and non-metallic elements. 0.93 3.16 NaCl sodium chloride (Cl-Na) 3.16 – 0.93 = 2.23 (> 2) 6  Polarisation – Cations In an ionic compound, due to the attraction of cations for anion electrons, the spherical electron cloud of the anion becomes distorted – polarisable. Cations become smaller, more highly charged and more strongly polarising from left to right across a period (Be2+ v Li+, Mg2+ v Na+, Al3+ v Mg2+). Cations become larger and less polarising down a group (K + v Na+ v Li+). Polarising power of diagonal neighbours are similar (Li + ≈ Mg2+, Be2+ ≈ Al3+). NaCl AlF3 MgCl2 AlCl3 AlCl3 AlBr3 sizes of atoms and their ions (pm) (1 pm = 10 –12 m) 7  Polarisation – Anions Anions become smaller, less highly charged and less strongly polarisable (electron cloud becomes less distorted) from left to right across a period (O2– v F–). Anions become more polarisable as the cation they are associated with becomes more positively charged. (NaCl v MgCl2 v AlCl3). Anions become larger and more polarisable (electron cloud becomes more distorted) down a group (F– v Cl– v Br– v I–). NaCl AlF3 MgCl2 AlCl3 AlCl3 AlBr3 sizes of atoms and their ions (pm) (1 pm = 10 –12 m) 8  Dipoles and Partial Charges Partial charges – small shifts in distribution of electrons. Electric dipole – positive charge next to an equal but opposite negative charge. Electric dipole moment – magnitude of the dipole (μ, debye). Molecule μ Molecule μ HF 1.91 H2O 1.85 HCl 1.08 NH3 1.47 HBr 0.80 O3 0.53 NaCl 9.00 CO2 0 BF3 0 CH4 0 9  Polar Bonds v Polar Molecules Polar bond – covalent bond between atoms with partial electric charges. Polar molecule – molecule with non-zero electric dipole moment. Polar bond v polar molecule – Although each bond in a polyatomic molecule may be polar, the molecule as a whole may be nonpolar if its shape is such that the dipoles of the individual bonds point in opposite directions and cancel each other. CH4 (methane) (μ = 0) CH3Cl (chloromethane) (μ = 1.87) CH2Cl2 (dichloromethane) (μ = 1.55) CHCl3 (chloroform) (μ = 1.04) CCl4 (carbon tetrachloride) (μ = 0) Cl H Cl H Cl Cl Cl H H Cl H H C2H2Cl2 C2H2Cl2 C2H2Cl2 polar molecule non-polar molecule polar molecule 10  Testing Molecular Polarity using Vectors Vectors can be used to determine if a molecule has a net overall polarity While compounds and drugs can contain individual polar bonds, the net polarity of a molecule could be zero. Vectors have length and direction and can be added to and subtracted from each other. From a vector perspective, a vector can represent a polar bond. The direction of the vector is the direction of the more polar atom along the bond axis. The length of the vector is a measure of the polarity of the bond. A δ– A+B B δ + Parallelogram method 11  Testing Molecular Polarity using Vectors – I 1 3 2.20 1 Cl H 4 5 6 2 Cl H 3 2 3.16 2.55 4 5 6 7 δ– C2H2Cl2 δ+ polar molecule 12 Testing Molecular Polarity using Vectors – II 1 3 2.20 3.16 5 Cl H 3 4 1 2 6 H 2.55 Cl 2 4 5 6 =0 C2H2Cl2 non-polar molecule 13 Testing Molecular Polarity using Vectors – III 1 3 3.16 5 Cl Cl 6 1 2 2.20 3 H 2.55 H 4 2 4 δ– 7 7 5 6 δ+ C2H2Cl2 polar molecule 14  Testing Molecular Polarity using Vectors – IV C2H2Cl2 C2H2Cl2 C2H2Cl2 polar molecule non-polar molecule polar molecule A B C Cl H Cl H Cl Cl Cl H H Cl H H A B=0 A C 0 B C A>C>B 15