Medical Chemistry L1 Lecture Notes PDF

Medical Chemistry L1 Lecturer: Prof. Dr. Giovanni N. Roviello Medical Chemistry Teaching for Geomedi University, Tbilisi, Georgia Definition Medical Chemistry (also referred to as Medicinal Chemistry): is the branch of science focused on the discovery, design, an...

Medical Chemistry L1 Lecturer: Prof. Dr. Giovanni N. Roviello Medical Chemistry Teaching for Geomedi University, Tbilisi, Georgia Definition Medical Chemistry (also referred to as Medicinal Chemistry): is the branch of science focused on the discovery, design, and development of new therapeutic chemicals, which are then formulated into useful medicines. Hoefker, Ashley, "Relating Chemistry to Society to Promote the Understanding of the World for High Schoolers" (2023). School of Education and Leadership Student Capstone Projects. 971. https://digitalcommons.hamline.edu/hse_cp/971 Who am I? Giovanni Roviello Assistant Professor at Geomedi University, Tbilisi, Georgia Senior Researcher at the Institute of Biostructures and Bioimaging (IBB-CNR*), Naples, Italy *Italian National Research Council https://www.ibb.cnr.it/?command=viewu&id=387 Research Areas Medicinal Chemistry Drug Discovery Peptide Chemistry Spectroscopic Studies DNA Structure Phytodrug Development Top 2% Scientist in Stanford University’s 2023 Author Database (Single Year, Ioannidis, J. P. A., "Updated Science-Wide Author Databases", 2024). https://elsevier.digitalcommonsdata.com/datasets/btchxktzyw/7 https://topresearcherslist.com/Home/Profile/974858  importance of general & inorganic chemistry fundamentals for a course in Medical Chemistry: Atomic structure and electron configuration essential for understanding molecular interactions. Basics of molecular structures and bonding, essential for drug interaction. Understanding the chemical behavior of biologically-active molecules  The electronic structure of Atoms Atoms: Democritus (around 470-370 B.C.) proposed that all matter is composed of tiny, indivisible particles, which he termed "atoms," derived from the Greek word meaning "indivisible.“ Scientific theory: Dalton, early 19th century A.D. Matter is made up of atoms (indivisible). Atoms of the same element are identical in size and properties. Different substances are made of different atoms. Chemical reactions involve the rearrangement of atoms, but not the creation or destruction of atoms (conservation of mass = conservation of the number and type of atoms!). 10 grams 80 grams 90 grams REAGENTS PRODUCTS The electronic structure of Atoms Various experiments conducted between the late 1800s and early 1900s demonstrated that atoms are not indivisible, but are made up of smaller (elementary) particles. SUBATOMIC PARTICLES: Fundamental particles The modern description of the atom, and thus of all matter, is based on three fundamental particles: electrons, protons, and neutrons. The electronic structure of Atoms Experiment with the Cathode Ray Tube (1890):Joseph John Thomson Discovery of a particle with a negative charge and mass, called the electron, with the same characteristics regardless of the cathode used.  The electronic structure of Atoms Ernest Rutherford (1871-1937) Most of the α particles, which are positively charged particles with mass, pass through the thin foil without any change in their trajectory. This indicated the presence of empty space within the atom! The nucleus, which is positive, is at the center and is as small as a marble in comparison to a soccer field. The atom's positive charge (protons) is concentrated at the center, in a very small volume (nucleus). Planetary Atomic Model: The positively charged nucleus is located at the center. The negatively charged electrons orbit around the nucleus. The electronic structure of Atoms Ernest Rutherford's Experiment: Few α particles hit the nuclei and are deflected back towards the source. Particle beam directed at atoms in a thin metal foil. Electrons occupy the space outside the nucleus. Most α particles pass through undisturbed or undergo only a very slight deflection. A small fraction of the particles undergoes significant deflection. Atomic Dimensions: Atomic size: approximately 1 Å = 10⁻¹⁰ m = 0.1 nm Nuclear size: approximately 10⁻⁵ Å Key Insights: The majority of the atom is empty space. Nearly all of the atom's mass is concentrated in the nucleus. The electronic structure of Atoms In the nucleus, positively charged particles with much greater mass than electrons: protons. Comparison between hydrogen and helium: Hydrogen: 1 proton Helium: 2 protons, but about 4 times the weight of hydrogen! Discovery of another subatomic particle: Neutron: A particle with mass similar to that of a proton, but without an electric charge (discovered by James Chadwick). Important observation: Elements are distinguished from each other based on the number of protons in the atomic nucleus.  The dual nature of the electron Particle nature: Electrons can be thought of as particles with mass and charge, which explains their ability to collide with other particles and interact through forces like the electromagnetic force. Wave nature: According to de Broglie, electrons also behave like waves, having a wavelength associated with their motion. This explains phenomena like electron diffraction and the quantization of energy levels in atoms, where only certain wavelengths (or frequencies) are allowed. This concept was confirmed through experiments such as the electron diffraction experiment, which showed that electrons can form patterns typically associated with waves, like light waves. The wave-particle duality is fundamental in quantum mechanics, where the electron is described by a wave function that gives probabilities for its position and momentum rather than definite values.  Quantum mechanics Quantum mechanics is the branch of physics that deals with the behavior of matter and energy at very small scales, such as atoms and subatomic particles. Unlike classical mechanics, which describes the motion of larger objects, quantum mechanics explains phenomena that cannot be understood with classical physics, particularly at the level of individual particles. Here are some key principles of quantum mechanics: Wave-particle duality: Particles, like electrons, exhibit both wave-like and particle-like behavior. This is best exemplified by the wave-particle duality, where particles can act as waves under certain conditions (e.g., diffraction), and waves can exhibit particle-like properties (e.g., photon impact). Quantization: Certain properties, like energy, are quantized, meaning they can only exist in discrete amounts or "quanta." For example, electrons in an atom can only occupy specific energy levels, not any value in between. Uncertainty principle: Proposed by Werner Heisenberg, it states that the more precisely you know a particle's position, the less precisely you can know its momentum (and vice versa). This principle highlights the inherent limitations in measuring certain properties simultaneously.  Quantum Mechanics Application to the Hydrogen Atom The Hydrogen Atom Bohr's model predicted, based on ad hoc considerations, the quantization of atomic radii and energies in the hydrogen atom. Applying Schrödinger's equation to the hydrogen atom system gives us the same result, but this time as the outcome of using a precise formalism. Quantum Mechanics Application to the Hydrogen Atom Assumptions: The nucleus has infinite mass The nucleus is centered in the Cartesian reference system. We must determine the total energy Et of the system: r is the position vector that defines the electron's position. In this framework, the solution of Schrödinger's equation leads to specific quantum states for the electron. Solving this equation is not trivial. First, we note that an electron orbiting the nucleus presents a spherical symmetry problem. We can switch to radial or polar coordinates. Ψ(x, y, z) → Ψ(r, θ, φ) In such a situation, it can be demonstrated (although we won’t prove it here) that the radial part and the angular part can be separated and treated independently: Ψ(r, θ, φ) = Rn,l(r) Yl,m(θ, φ) We will define the meaning of these three quantum numbers later. Schrödinger's equation splits into two simpler equations. The eigenvalue of the energy (solution of the radial Schrödinger equation) is: We notice that this is the same expression obtained from Bohr's model! It results from Schrödinger's equation, which is completely general and takes into account the actual potential that applies in this case. E₁ = Ionization energy (with the sign reversed) The energy required to move the electron to a non-bound (free) state → ionization of the atom. Quantum Mechanics Application to the Hydrogen Atom quantum numbers Defining the two quantum numbers, n (the principal quantum number, which takes values n = 1, 2, 3,...) and l (the secondary or angular quantum number, which takes values l = 0, 1, 2,..., n – 1): For the ground state of hydrogen (n = 1), there will be only one value for l, with an energy of -13.6 eV. For the first excited state (n = 2), the values of l become two (l = 0 and l = 1). These correspond to a single energy level (since En​ depends only on n): in this case, we will talk about degenerate energy levels. These levels do not represent the same physical system, as they are represented by different wavefunctions. Key points: n (principal quantum number) determines the overall energy level of an electron in an atom. l (orbital angular momentum quantum number) defines the shape of the orbital. Degenerate energy levels refer to states with the same energy but represented by different wavefunctions (e.g., different orbital shapes). This description explains how the quantum numbers govern the behavior of electrons in different energy states and highlights the idea of energy degeneracy in excited states.  Principal Quantum Number (n): Specifies the energy level of an electron in the atom: determines the energy of the orbital. Secondary (or Angular) Quantum Number (l): Indicates the shape of the orbital in which the electron is found (s, p, d, f). Magnetic Quantum Number (m): Specifies the orientation of the orbital. m assumes integer values such that −l ≤ m ≤ +l (including 0). Spin Magnetic Quantum Number (ms): Indicates the direction of rotation of the electron in an orbital with values +1/2 and -1/2. To this, we add a final quantum number, called the spin quantum number, which can have two possible values: +1/2 and -1/2. This quantum number was introduced to explain an additional degeneracy observed in an appropriate experiment (Stern and Gerlach, 1924). Intuitively, this corresponds to the idea that the electron, in addition to moving around the nucleus, can also spin on its own axis, either in a clockwise or counterclockwise direction. Pauli’s Exclusion Principle states that two distinct electrons cannot be described by the same quantum state, i.e., by the same quadruple (n,l,m,s). summary orbitals Electrons arrange themselves around the nucleus according to three rules: The Aufbau principle, according to which electrons first occupy orbitals with the lowest energy, starting from the position closest to the nucleus. The Pauli exclusion principle, according to which an orbital can be occupied by no more than two electrons, and these electrons must have opposite spin quantum numbers The Hund's rule or the maximum multiplicity rule states that if there are multiple degenerate orbitals (orbitals with the same energy), electrons will first occupy each empty orbital singly, and then they will pair up in the orbitals. The most stable electron configuration in a subshell is the one with the maximum number of unpaired electrons with parallel spins. This arrangement allows the electrons to be as far apart as possible, minimizing mutual repulsion. electronic configuration The notation used to represent which orbitals are occupied is called the electronic configuration. The electronic configuration of an atom is typically represented graphically by indicating the electrons with arrows (representing electron spins), or symbolically by writing: The energy level number The symbol of the subshell The total number of electrons in that subshell as an exponent. For example, the notation 1s² means two electrons in the 1s orbital and is read as "one s two."  The shielding effect in many-electron atoms The shielding effect refers to the phenomenon where inner-shell electrons reduce the effective nuclear charge felt by outer-shell (valence) electrons in a multi-electron atom. Electron-Electron Repulsion: Electrons in inner orbitals shield outer electrons from the full positive charge of the nucleus due to repulsion between electrons. Effective Nuclear Charge (Z_eff): The effective nuclear charge is the net positive charge experienced by an electron in an atom. It is less than the actual nuclear charge (Z) due to the shielding by other electrons. Zeff=Z−S where S is the shielding constant (the effect of inner electrons). Shielding and Atomic Size: Greater shielding leads to a weaker attraction between the nucleus and the valence electrons, which can result in larger atomic radii. Impact on Chemical Properties: Shielding influences the reactivity of atoms. For example, the more shielded the outer electrons are, the easier it is for them to be lost in reactions. Order of Shielding: Electrons in the same shell (such as 2s and 2p) shield each other to a certain extent, but not as effectively as inner shell electrons (like 1s) shielding outer shell electrons.  Z and A The atomic number, or the number of protons within the nucleus, defines the element to which the atom belongs: atoms of the same element have the same atomic number. It is represented by the letter Z. Mass Number: The number of nucleons present in the atom: protons + neutrons.Indicated by the letter A. Number of neutrons = A – Z inprotected.com