Lec. 1 | Atomic & Molecular Structure PDF

‫‪Inorganic medicinal and‬‬ ‫‪pharmaceutical chemistry‬‬ ‫‪by : Block. Roche‬‬ ‫‪lec 1: Atomic and molecular Structure /‬‬ ‫‪complexation‬‬ ‫ﺻ ــﻮروا اﻟ ــﺒﺎرﻛ ــﻮد او اﺿ ــﻐﻄﻮا ﻋ ــﻠﻴﻪ ﻟ ــﻠﻮﺻ ــﻮل اﻟ ــﻰ‬ ‫ﻓـﻴﺪﻳـﻮ ﻳـﻨﻄﻴﻜﻢ ﻓـﻜﺮة ﻋـﺎﻣـﺔ ﻋـﻦ اﻟـﻤﺤﺎ...

‫‪Inorganic medicinal and‬‬ ‫‪pharmaceutical chemistry‬‬ ‫‪by : Block. Roche‬‬ ‫‪lec 1: Atomic and molecular Structure /‬‬ ‫‪complexation‬‬ ‫ﺻ ــﻮروا اﻟ ــﺒﺎرﻛ ــﻮد او اﺿ ــﻐﻄﻮا ﻋ ــﻠﻴﻪ ﻟ ــﻠﻮﺻ ــﻮل اﻟ ــﻰ‬ ‫ﻓـﻴﺪﻳـﻮ ﻳـﻨﻄﻴﻜﻢ ﻓـﻜﺮة ﻋـﺎﻣـﺔ ﻋـﻦ اﻟـﻤﺤﺎﺿـﺮة ﻳﺴﻬـﻞ‬ ‫ﻋﻠﻴﻜﻢ دراﺳﺘﻬﺎ ‪.‬وﻻ ﺗﻨﺴﻮﻧﺎ ﻣﻦ دﻋﺎﺋﻜﻢ 🙏‬  Electronic structure of atoms ATOMS ARE COMPOSED OF A CENTRAL NUCLEUS SURROUNDED BY ELECTRONS WHICH OCCUPY DISCRETE REGIONS OF SPACE. THE NUCLEUS IS CONSIDERED TO CONTAIN TWO TYPES OF STABLE PARTICLES WHICH COMPRISE MOST OF THE MASS OF THE ATOM. 1- NEUTRON: IT IS AN UNCHARGED SPECIES WITH A MASS OF 1.675 X I0-24 2- PROTON: THIS PARTICLE HAS A POSITIVE CHARGE OF ESSENTIALLY ONE ELECTROSTATIC UNIT (E.SU.). ITS MASS IS 1.672 X I0-24. 3- ELECTRON : WHICH HAS A NEGATIVE CHARGE OF ONE E.S.U AND A MASS OF 9.107 X 10-28 THE NUMBER OF PROTONS (EQUAL TO THE NUMBER OF ELECTRONS IN THE NEUTRAL ATOM) AND A PARTICULAR NUMBER OF NEUTRONS. THE SUM OF THE MASSES OF THE PROTONS-AND NEUTRONS ACCOUNTS FOR THE ATOMIC MASS OF THE ELEMENT AND THE NUMBER OF PROTONS IS EQUAL TO THE ATOMIC NUMBER.  ATOMIC ORBITALS Atomic orbital is the volume of space that contain the electrons about the nucleus and they described by a set of 4 quantum no. : 1- The principal quantum no. ( n ) :which state electrons exist in discrete energy levels. The energy associated with electron increases as it locate farther from the nucleus. 2- The suborbital quantum no. ( l ) :which represent the region of greatest probability of finding an electron varies in shape and size , depending upon energy level. l = 0,1,2,3,….(n-1 ) , so when n=1 the l value =0 and when n=2 the l take 2 values = 0,1 l = 0 is s orbital , l=1 is p orbital , l = 2 is d orbital , l = 3 is f orbital 3- The magnetic quantum (m l ) : describes the spatial orientation of the orbital 4- The spin quantum no. (m s) : represent the magnetic moment which is directionally oriented +1/2 , -1/2 , so if 2 electrons occupy the same orbital they must have an opposing spin. - The aufbau principle : is used to determine the electron configuration of an atom, molecule or ion. The principle postulates a hypothetical process in which an atom is "built up" by progressively adding electrons. - The pauli exclusion principle : in any atom , no 2 electron may be described by the same set of values of the quantum no.s. So a maximum of 2 electrons may occupy a single orbital and must be of opposed spin. - Hund’s rules : 1-lower energy orbitals must be filled before higher energy orbitals 2-electrons must enter degenerate orbials singly and with parallel spins and remain unpaired as long as possible.  1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ In certain elements in the transition series where d orbitals are being filled, an ns level will be only half- filled, and the (n — 1)d orbital will be either half-filled or full. These represent more stable configurations than may be achieved by simply filling orbitals. For example, chromium (At. No. 24) has an outer structure of 3d54s1 and copper (At. No. 29) has 3d104s1. It is possible to use the inert gas “core” that precedes the element being considered. For example, sodium (At. No. 11), which has the electronic configuration Na: 1s22s22p63s1, can be written using the neon core for the first ten electrons: Na: [Ne] 3s1. Similarly, manganese (At. No. 25) can be written using the argon core for the first 18 electrons: Mn: [Ar] 3d54s2.  Ionization - Ionization : is The process of losing one or more electrons by chemical or physical means , and the positive ion produced is termed a cation. This process is based in physical reality, and should not be taken as the exact opposite of the process of atom buildup. - It is always the most loosely “held” electrons which are lost first when an atom ionizes. - The electronic structure of the ion may not reveal the level from which the electron was lost. This is particularly true for transition elements. Because Relative orbital energies are subject to change as electrons are “placed” in them. This means that a high energy orbital in one atom may be of lower energy in a neighboring atom where it might be completely filled. Also the possibility of rearrangement of the remaining electrons in an ion to a more stable configuration. - Atoms in the transition series with incompletely filled d orbitals will ionize to leave d ions in which may contain from one to 10 electrons ex: cobalt - This does not necessarily mean that both electrons were lost from the 4s orbital, even though the structure of the ion would seem to indicate that. one or both electrons could have been removed from the 3d orbital followed by rearrangement of all the valence electrons into this orbital. Elements in Groups VIA and VIIA w h i c h h ave l a r g e r n u m b e r s o f electrons in their’ p orbitals tend to ionize by accepting electrons to form anions. These ions have completely filled p orbitals so that the valence shell structure is the same as the inert gas in the same period as the neutral element. Examples of this can be seen in oxygen (At. No. 8) arid bromine (At. No. 35):  Periodic table - The periodic table is a tabular arrangement of the chemical elements, organized on the basis of their atomic numbers, electron configurations, and recurring chemical properties. Elements are presented in order of increasing atomic number. The standard form of the table consists of a grid of elements laid out in 18 columns and 7 rows. - The rows of the table are called periods; the columns are called groups, with some of these having names such as halogens or noble gases. - All elements from atomic numbers 1 (hydrogen) to 118 (ununoctium) have been discovered or reportedly synthesized, with elements 113, 115, 117 and 118 having yet to be confirmed. The first 98 elements exist naturally although some are found only in trace amounts and were initially discovered by synthesis in laboratories. A group :which are eight groups which correspond to the filling of s and p orbitals having principal quantum numbers equal to the number of the period or row of the table. Through Periods 2 and 3 s and p orbitals are filled in normal fashion. These are sometimes referred to as the “typical elements.” In the fourth and successive periods, d orbitals are filled “between” the s and p orbitals, giving rise to the “B groups” intervening between Groups IIA and IIIA. - In Periods 4, b, and 6, toward the center of the table, there is a triad of elements designated Group VIII. These are the “real” transition elements which occur between the group with elements having half- filled d orbitals (VIIB) and the group with elements having full d orbitals (IB). - The separated long periods below the main portion of the table contain the elements known as the lanthanides (atomic numbers 57 to 71) and the actinides (atomic numbers 89 to 103). These elements are formed by the filling of very low-lying f orbitals. - Electronegativity : is the affinity an element has for its electrons and the ability to take on additional electrons which will show that this property increase from left to right across any period and from bottom to top in any group (except VIIIA). The opposite concept is that of electropositivity which varies in directions opposite to those of electronegativity. - The ability to lose electrons increases as we descend in a particular group. This is due to the increased shielding effect of the inner electrons which diminishes the attractive force of the nucleus on the valence electrons.  Electronic structure of molecules - There are three major forces that are involved in the formation of molecules : (1) Coulombic attraction occurs between the negatively charged electrons in the valence orbitals on one atom and the positively charged nucleus of another atom.(2) the number of electrons in the valence shell orbitals, and (3) their orbital distribution. - There are 2 types of bonds : 1- Covalent bond which ranges from an equal sharing of a pair of electrons in homonuclear diatomic molecules (e.g., H2 , Cl2 , I2 , etc.) to a polar or unequal sharing of the electron pair in heteronuclear diatomic molecules (e.g., HCI). 2- Ionic bond which is more of an electrostatic interaction resulting from the transfer of an electron from an electropositive atom to an electronegative atom (e.g., Na+ Cl—). - Orbital hybridization is “mixing” of the atomic orbitals to provide a new set of degenerate (energetically equivalent) orbitals having different spatial orientations and directional properties than the original atomic orbitals. The number of hybrid orbitals produced is equal to the number of atomic orbitals involved in the hybridization, and the electrons contained in the original orbitals occupy the hybrids according to Huad’s rules. - The presumed mechanism allowing elements to increase covalent bonding capacity involves promotion to the valence state, a situation requiring energy.  ‫لحساب نوع التهجني بني الذرات‬ ‫ عدد االرتباطات لهذه الذرة )مجاميع( ناقص واحد يساوي نوع التهجني‬-1 (long pair)‫ املزدوج االلكتروني في الحسابات يعتبر مجموعة كاملة‬-2 There are 3 cases of hybrid orbitals 1- sp orbitals : The two sp orbitals are equivalent (degenerate), symmetrical about the bonding axis, and oriented 180° away from each other. The improved directional characteristics of the hybrid should also be noted. This type of hybridization (sp) is exist in the covalent compounds of Group II elements and the unsaturatedacetylenic compounds of carbon. Linear covalent molecules of gaseous halides of Be, Mg, and Ca, such as MgCI2, and the solid divalent compounds of Cd and Hg are indicative that the bonds are formed through sp hybrids on the Group II element. 2- sp2 orbitals : which involve singly occupied s and two p orbitals combine to form three equivalent sp2 hybrid orbitals. The three hybrids are located in the same plane, and are oriented toward the points of an equilateral triangle, 1200 apart. The monomeric covalent compounds of boron, aluminum , and other Group III elements , unsaturated “ethylenic” compounds of carbon, show sp2 hybridization. 3- sp3 orbitals : which include the tetravalent state of Group IVA elements. When one s and three p orbitals combine, the result is a set of four equivalent sp3 hybrid orbitals pointing to the four corners of a tetrahedron. Therefore, the geometry of a molecule formed through bonding with these orbitals is tetrahedral, and the bond angles are approximately 1090. - Promotion to hybridized states is presumed to occur with doubly occupied orbitals. In other words, it is possible to have a nonbonded pair of electrons in a hybrid orbital. ex H2O. In the ground state electronic structure of oxygen, there are two p orbitals containing one electron each. If the two hydrogen atoms became bonded to the oxygen through these two orbitals, the water molecule would be expected to have an H—O—H bond angle of 1800. - In fact, the bond angle is closer to 1040.The smaller angle cannot be explained on the basis of repulsion between the two polar O—H bonds, It is because that the valence shell orbitals on oxygen achieve a hybridized valence state. orbitals on oxygen achieve a hybridized valence state. Assuming sp3 hybridization, the water molecule would have two lone pairs of electrons in hybridized orbitals. Repulsive forces between these two orbitals will cause the angle between them to enlarge and the orbitals to lose p character. Such a change would in turn cause the H—O—H angle to become smaller, with a corresponding gain in p character for the bonding orbitals. Thus, the two lone pairs of electrons would occupy orbitals which are hybridized somewhere between sp2 and sp3, while the bonding orbitals on the oxygen are between sp3 and pure p orbitals.  Types of bonding interactions 1- Ionic bonding: is the electrostatic force that exists between two chemical entities of opposite charge (The cation : the least electronegative entity loses one or more of its valence electrons) and the negative species (the anion : the more electronegative entity ). It will be found that most stable ions have inert gas valence shell structures. Since the valence shell of all inert gases except helium contain eight electrons, this kind of structure is associated with stability, and has led to the octet theory of chemical bonding. Ionic bonding is usually found in associations between metallic, strongly electropositive elements (Groups IA and IIA) and nonmetallic, strongly electronegative elements (Group VIllA). It is also found in most salts where the anion is complex, such as SO4-2, PO4-3, NO3-. Ionic interactions are also present, in many cases, when polar compounds are dissolved in polar solvents. Hybridization is not involved in the formation of ionic compounds. - Generally speaking, metals lose electrons to form cations, and nonmetals attract electrons to form anions. - In the transition series, the octet theory is less obvious in ion formation. These metals will form variously charged cations , usually with electrons remaining in the d orbital valence shells. The low-lying d orbitals are responsible for the variable valences and electronic structures of these ions. 2- Covalent bonding : is the attractive force that exists between two chemical entities due to their “sharing” a pair of electrons. - The covalent can be either nonpolar or polar , in the nonpolar , the electron pair is shared equally by the two bonded atoms like in homonuclear diatomic molecules such as H2 , Cl2 , N2 , etc. Larger systems of atoms are also covalent with an equal distribution of the bonding electrons, e.g., S8, P4 and diamond carbon. The covalent bonds in saturated hydrocarbons approach an ideal sharing of electrons between the carbon atoms, but the carbon-hydrogen bonds depart from this to some extent. In The polar bond The electron density tends to be shifted toward the more electronegative member of the bond ex, Water which is strongly polar with the more electronegative oxygen taking the largest share of the bonding electron density. this polarity is indicated by noting partial charges. - there is another classification of covalent bond (sigma and pi): in sigma bonds the molecular orbitals (or electron distributions) are symmetrical about the bond axes. - In pi bond (the double and triple bonds) there are two and three pairs of electrons are being shared, respectively, between two atoms. only one of the bonds can be a alpha bond. The other pairs of electrons will occupy molecular orbitals which are distributed on both sides of and perpendicular to a plane passing through the bond axis. Carbon dioxide, CO2, is an illustration of a double covalently bonded molecule. Each C—O bond consists of two pairs of electrons.The sigma bonds are formed by overlapping sp orbitals on the carbon with singly occupied p orbitals on the oxygens. The oxygens can be rotated in turn to bring their p orbitals parallel to one of the carbon p orbitals, Overlapping of parallel porbitals on carbon and oxygen leads to 2 volumes of electron density separated by a nodal plane at the bond axis. The carbon thus forms a pi bond with each oxygen. Similarly, hydrogen cyanide, HCN , will serve as ex. of a triple bonded molecule where two pi bonds are formed between two atoms.The carbon is sp- hybridized as in CO2, but now both p orbitals overlap with two p orbitals on the nitrogen. 3 - Coordinate Covalent Bonding : is a covalent interaction which both electrons in the bond arise from a single orbital on one of the atoms forming the bond. It is found most frequently between complex chemical entities. The entity providing the pair of electrons is called as the donor species. The acceptor species is electron-deficient and has an empty orbital which can overlap with the orbital from the donor. ex; boron-trifluoride etherate complex. - This type of bond formation also occurs in acid-base chemistry, and is frequently the type of bonding one finds between sulfur and oxygen. In particular in the oxyacids, e.g., sulfuric, nitric, phosphoric. 4- Hydrogen bonding: it is a weak secondary interaction between the partial positive hydrogen atom that bonded to more electrongative element (e.g. O,N,F,CL) ,ect, with the nonbonding electrons on the other electronagative atoms in neighboring molecules. Hydrogen bonding is responsible for many of the physical and chemical ptoperties of water. For example, the relatively high boiling point. This type of association can also occur between unlike molecules, and plays an important role in solution formation. Hydrogen bonding is also important in interactions between complex molecules, and in the secondary structure of proteins. It is also a secondary binding force in drug-receptor interactions. 6- Van derwaals(London) Forces: These are very weak electrical dipole forces occur when the electrons in one atomic or molecular species inducing a repulsive distortion in the electron cloud of a neighboring species. The positive end of the dipole, which is essentially produced by protons in the nuclei has an attraction for the oppositely charged electrons in the same or in a neighboring species. Van der Waals forces are virtually the only attractive forces between nonpolar molecules ,The associations between aromatic hydrocarbon molecules such as benzene. These forces also function in the liquefaction and solidification of the inert gases, a process which requires extremely cold temperatures.  Resonance The resonance indicates a tendency for the electrons to be somewhat delocalized. Simple covalent molecules tend to be more localized than more complex systems. For example, the molecules HCL and CO2 , NO2 , benzen are represented by the following canonical (electronic picture) structures.