Inorganic Pharmaceutical Chemistry 2024-2025 PDF

Al -Maaref University College of Pharmacy Assist. Prof Dr.Sameerah Fenjan Hasan Ministry of Higher Education and Scientific Research Al -Maaref University College of Pharmacy Al-An...

Al -Maaref University College of Pharmacy Assist. Prof Dr.Sameerah Fenjan Hasan Ministry of Higher Education and Scientific Research Al -Maaref University College of Pharmacy Al-Anbar, Iraq The name of the course Inorganic. Pharmaceutical Chemistry. 2024--2025 Atomic &Molecular Structure, Quantum Numbers And Hyperization PREPARED BY: Assist. Prof. Dr. Sameerah Fenjan Hasan, B.Pharm., Ph.D. 1 Al -Maaref University College of Pharmacy Assist. Prof Dr.Sameerah Fenjan Hasan 2 Al -Maaref University College of Pharmacy Assist. Prof Dr.Sameerah Fenjan Hasan Electronic Structure of Atoms The fundamental unit of all matter is the atom. The various chemical and physical properties of matter are determined by its elemental composition, and elements are composed of like atoms and their isotopes. In order to be able to predict the properties of matter, molecules, or elements, it is important to understand the structure of atoms. Subatomic Particles 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. These particles are held within the nucleus by various nuclear forces. One of these particles is called a neutron. It is an uncharged species with a mass of 1.675×10-24 g or approximately 1.009 mass units on the atomic scale. The other particle is termed a proton. This particle has a positive charge of essentially one electrostatic unit (e.s.u.). Its mass is close to that of the neutron at 1.672×10- 24 g or approximately 1.008 atomic units Every stable nucleus contains a certain 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 most of the atomic mass (or weight) of the element, and the number of protons is equal to the atomic number. Isotopic forms of a particular element differ in the number of neutrons, and therefore, in the atomic mass. A third subatomic particle is the electron, which has a negative charge of one e.s.u. and a mass of 9.107×10-28 g or approximately 0.0006 a.m.u. Its charge is opposite in sign and equal in magnitude to that of a proton, so a neutral atom will have the same number of electrons as protons. The mass of the electron is about 1/1840 of that of the proton, thus providing only a small contribution to the atomic mass. Electrons occupy regions of extranuclear space at various distances from the nucleus according to the laws of quantum mechanics. 3 Al -Maaref University College of Pharmacy Assist. Prof Dr.Sameerah Fenjan Hasan Atomic Orbitals The early quantitative description of electronic structure came from Niels Bohr in 1913, and involved a planetary picture of the atom. Electrons were considered as particles which revolved around the nucleus in stationary planar orbits and which had definite energies. Bohr's model involved the use of classical Newtonian mechanics, and suitably predicted the electronic spectrum of hydrogen. However, atoms with more than one electron did not yield satisfactory results. The arrangement of the electrons The electrons are found at considerable distances from the nucleus in a series of levels called energy levels. Each energy level can only hold a certain number of electrons. The first level (nearest the nucleus) will only hold 2 electrons, the second holds 8, and the third also seems to be full when it has 8 electrons. These levels can be thought of as getting progressively further from the nucleus. Electrons will always go into the lowest possible energy level (nearest the nucleus) - provided there is space. The electronic structures of hydrogen and carbon, for example, drawn as: The circles show energy levels - representing increasing distances fromthe nucleus. You could straighten the circles out and draw the electronic structure as a simple energy diagram. 4 Al -Maaref University College of Pharmacy Assist. Prof Dr.Sameerah Fenjan Hasan Carbon, for example, would look like this: Thinking of the arrangement of the electrons in this way makes a usefulbridge to the A ‘level view. In the 1920s the theory of quantum mechanics for the description of ultrasmall particles was developed, as was the quantum theory of atomic structure. Relevant to the present discussion is the description of electrons, placing them in discrete volume of space about the nucleus. These volumes of space are referred to by the term atomic orbitals, and the electrons contained within their boundaries are described by a set of four numbers called quantum numbers. Quantum Numbers The four quantum numbers set the probability limits within which an electron can be found. The first three quantum numbers refer to some property of the space or orbitals, while the fourth quantum number describes the spin of the electron. (1) The Principal Quantum Number. This number is given the symbol n. electrons in atoms exist in discrete energy levels. The energy associated with the electron increases as it locates farther from the nucleus. The principal quantum number describes the relative position of these energy levels, their distance from the nucleus and the possibility of discontinuities or points of zero probability in the levels. The value this number can assume are integers from n = 1, 2, 3, ….., ∞. When n = 1 the electron is found in the energy level closest to the nucleus. 5 Al -Maaref University College of Pharmacy Assist. Prof Dr.Sameerah Fenjan Hasan (2) The Suborbital Quantum Number. This number is given the symbol l. The volume of space which represents the region of greatest probability of finding an electron varies in shape and size, depending upon the energy level. The suborbital quantum number may be said to describe the shape of the orbital, or the "electron cloud". This number can assume integer values limited by the corresponding value of n such that l = 0, 1, 2, … (n -1). Thus, when n = 1, the only permissible value is l = 0. When n = 2, l can take two values, l = 0 and l = 1. The value of l is generally designated by the letters s, p, d, and f for the orbitals as follows: l 0 1 2 3 orbital s p d f (3) The Magnetic Quantum Number. The symbol for this quantum number is ml. Basically, this number describes the spatial orientation of the orbital. For any value of l, there are (2l + 1) allowed values of ml. (The allowed values are restricted by the value of l and can be positive or negative integer values according to: ml = -1, …., 0, …., +1. Obviously, when l = 0 the orbital is a (spherical) and can only have one orientation, ml = 0. When l = 1, there are three possible orientations of the associated p orbital, ml = -1,0,+1. These correspound to the three p orbitals shown in Figure 1-1along the x, y and z axes. Likewise, when l = 2 there are five possible orientation for the d orbital, ml = -2, -1, 0, +1, +2. (4) The Spin Quantum Number. This number is represented by the symbol ms. The electron can be envisioned in its particular state as a spinning mass. Since it is charged, it will have a magnetic moment which is directionally oriented. Depending upon the direction of spin, there are two orientations of the magnetic moment, +½ or -½. These are the only two allowed values of ms. The significance of this is that for two electrons to occupy the same orbital they must have opposing spin. If one has ms = + ½, the other must have ms 6 Al -Maaref University College of Pharmacy Assist. Prof Dr.Sameerah Fenjan Hasan SUMMARRY Quantum Numbers n => principal quantum number, quantized energy levels, which energy level n = 1, 2, 3, 4, 5, 6, 7, etc. l => secondary quantum number, quantized orbital angular momentum, which sublevelor type of orbital s type orbital l = 0 p type orbital l = 1 d type orbital l = 2 f type orbital l = 3 g type orbital l = 4 m => magnetic quantum number, quantizedorientation of angular momentum, whichorbital within sublevel. s type orbital m=0 p type orbital m = +1, 0 or -1 one value for each of the three p orbitals d type orbital m = +2, +1, 0, -1 or -2 one value for each of the five d orbitals f type orbital m = +3, +2, +1, 0 -1,-2 or -3 one value for each of the seven f orbitals 7 Al -Maaref University College of Pharmacy Assist. Prof Dr.Sameerah Fenjan Hasan Electronic Structure of Molecules Most of the electrons are in atomic orbitals surrounding the individual nuclei, and the remainder (valence electrons) are in more generalized multinuclear molecular orbitals. When atoms are incorporated into molecules, there are three major forces that are involved in the overall combination. Coulombic attraction occurs between the negatively charged electrons in the valence orbitals on one atom and the positive charged nucleus of another atom. As the atoms approach each other, there are two repulsive forces that tend to "push" the atoms away: (1) electron-electron repulsion between valence electrons on neighbor atoms. (2) nuclear repulsion between neighboring nuclei. A stable molecule is possible when the proper balance between these forces exists, and the energy of the resulting system of atoms is less than the sum of the energies of the "isolated" atoms. The equilibrium distance that are evident between atoms in molecules (bond distance) are established largely by the interaction of these forces. However, simply having an overriding attractive force between two atoms does not assure the formation of a bond between them. Depending upon the type of bonding interaction which is likely to occur (eg., ionic, covalent, etc.), there are other criteria based on the differences in electronegativity, availability of electrons, and the nature of the valence state atomic orbitals. Bond formation is also affected by the number of electrons in the valence shell orbitals, and by their orbital distribution. The bonding types that are possible vary with the amount of "sharing" of electrons between the two atoms participating in the bond. Covalent bonding ranges from an equal sharing of a pair of electrons in homonuclear diatomic molecules (eg., H2, Cl2, I2, etc.) to a polar or unequal sharing of the electron pair in heteronuclear diatomic molecules (eg., HCl). Ionic bonding is more of an electrostatic interaction resulting from the transfer of an electron from an electropositive atom to an electronegative atom (eg., Na+Cl-). 8 Al -Maaref University College of Pharmacy Assist. Prof Dr.Sameerah Fenjan Hasan Orbital Hybridization The process of orbital hybridization may be envisioned as a "mixing" of the atomic orbitals to provide a new set of degenerates (energetically equivalent) orbitals having different spatial orientation 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 Hand's rules.Examples using Be, B and C including shapes and properties, Ground State of Be, B, & C The presumed mechanism allowing these and other elements to increase covalent bonding capacity involves promotion to the valence state, a situation requiring energy. This is a no observable hypothetical state of the atom which is justifiable on certain theoretical ground and on the basis of the stability of the resulting molecules. Figure 1-3 a show a portion of an energy diagram for the promotion of gaseous beryllium atoms (Be(g)) to the divalent valence state. This state is usually viewed as hybrid orbital state and is labeled sp in this particular case. Similar diagrams are shown for boron (Figure 1-3 b) and carbon (Figure 1-3 c) indicating that their tri- and tetravalent states also involve hybrid orbitals designated as sp2 and sp3, respectively. The tetravalent state for carbon is limited to saturated molecules. 9 Al -Maaref University College of Pharmacy Assist. Prof Dr.Sameerah Fenjan Hasan Figure 1-3 Partial energy diagrams showing the promotion of atoms in their gas phase to their valence state configurations. The distances between states are not meant to indicate relative energy barriers. Directing attention to the construction of hybrid orbitals, the first case to be considered is that of sp orbitals. This pair of hybrids is often referred to by the term digonal to describe the fact that they are two opposing orbitals. This type of hybridization (sp) is evident in the covalent compounds of Group II. Linear covalent molecules of gaseous halides of Be, Mg, and Ca, such as MgCl2, and 10 Al -Maaref University College of Pharmacy Assist. Prof Dr.Sameerah Fenjan Hasan the solid divalent compounds of Cd and Hg are indicative that the bonds are formed through sp hybrids on the Group II element. In a manner similar to that described above, elements of Group III may be promoted to a valence state in which singly occupied s and two p orbitals combine to form three equivalent sp2 hybrid orbitals (Figure 1-3b) The overall appearance of one of these hybrids is similar to the sp orbital except that the inclusion of an additional p orbital causes the hybrid to be somewhat more elongated. The three hybrids are located in the same plane, and are oriented toward the points of an equilateral triangle, 120o apart. Figure 1-5 shows the three trigonal hybrids on the same set of axes. The monomeric covalent compounds of boron, aluminum, and other Group III elements, as well as unsaturated "ethylenic" compounds of carbon, shown sp2 hybridization. The empty p orbital remaining in the valence state (Figure 1-3b) of the Group III elements leaves their compounds electron-deficient. These molecules, therefore, react as Lewis acids. The final extension of hybridization between s and p orbitals is involved in the tetravalent state of Group IVA elements (Figure 1-3c). 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 109o Hybridization schemes involving d orbitals are involved in transition metal complexes. Two of the more important hybrids in this group include a set of six orbitals with octahedral geometry termed d2sp3 orbitals, and a set of four orbitals with square planar geometry termed dsp2 orbitals. sp, sp2, sp3, d2sp3. The effect of ligand strength and the magnetic properties of the complex in determining shape e.g., octahedral, tetrahedral or square planar. Some general points about hybrid orbitals can now be made 1- Although the term "equivalent" has been used in referring to sets of orbitals, it should be noted that the equivalency is destroyed when the orbitals on a particular atom form bonds to different elements in the formation of molecules. 2- In the considerations above, the orbitals promoted to the valence state were singly occupied. Promotion to hybridized states is presumed to occur with doubly occupied orbitals as well. In other words, it is possible to have a nonbonded pair of electrons in a hybrid orbital. 11 Al -Maaref University College of Pharmacy Assist. Prof Dr.Sameerah Fenjan Hasan 3- Bond strengths tend to increase as the amount of s character in a hybrid decreases. This should be understandable since the p character (or d character) would alternately increase, thus providing better directional properties to the hybrid. 4- Finally, bonds formed with hybrid orbitals will not necessarily have the exact geometries presented above. In many molecules, the bond angles will vary from that predicted because repulsive forces will cause changes in the s and p character leading to some contribution intermediate between those discussed above. 12