Medical Chemistry L2 PDF

Medical Chemistry L2 Lecturer: Prof. Dr. Giovanni N. Roviello Medical Chemistry Teaching for Geomedi University, Tbilisi, Georgia Periodic Table of Elements It is a catalog of all known elements, organized according to their characteristics. Similar elements were pl...

Medical Chemistry L2 Lecturer: Prof. Dr. Giovanni N. Roviello Medical Chemistry Teaching for Geomedi University, Tbilisi, Georgia Periodic Table of Elements It is a catalog of all known elements, organized according to their characteristics. Similar elements were placed in the same column, and elements with increasing mass were placed one after the other. Dmitri Mendeleev and Lothar Meyer independently cataloged the known elements at the same time. Periodic Table of Elements  Periods: horizontal Increasing size; Groups: vertical Similar chemical properties (state of matter, reactivity,...)  Metals: Solids at room temperature; Good conductors of electricity; Malleable and shiny; When they react, they form ions with a positive charge Non-metals: Gases or opaque solids at room temperature Poor conductors of electricity When they react, they form ions with a negative charge Metalloids: Have intermediate properties Semiconductors, widely used in technological applications A fundamental property of the atom is its mass. The masses of elementary particles measured in kilograms are as follows: Proton: 1.672623 x 10⁻²⁷ kg Neutron: 1.674929 x 10⁻²⁷ kg Electron: 9.1093690 x 10⁻³¹ kg The mass of the proton is 1836 times greater than that of the electron!!! UMA The masses of atoms measured in kilograms or grams are very small values. For example, the mass of a carbon-12 (¹²C) atom is 1.99 × 10⁻²⁶ kg (6 protons + 6 neutrons). Therefore, it is more convenient to use relative atomic masses, which are atomic masses compared to a reference value. To this end, a reference mass was introduced: THE ATOMIC MASS UNIT (u or uma) (also called the dalton), which corresponds to one twelfth of the mass of a carbon-12 atom, or 1.66 × 10⁻²⁷ kg. The relative atomic mass of an atom is determined by comparing it to the atomic mass unit. Thus, the relative atomic mass indicates how many times the mass of an atom is greater than one twelfth of the mass of a carbon-12 atom. EXAMPLE: The relative atomic mass of oxygen = 15.9994 uma = 15.9994 times greater than 1/12 of the mass of a carbon-12 atom. Periodic Table Mendeleev 1869: Classification of known elements based on their chemical properties. Meyer 1869: Classification of known elements based on their physical properties. Regular periodic repetition of certain properties as atomic weight increases. Periodicity The modern periodic table of elements takes into account our understanding of atomic structure and follows the basic layout of Mendeleev's table, preserving its division into periods and groups, but is enriched with a block structure. The application of quantum theory to the periodic law led to the redesign of the periodic table into its extended form with 18 groups. This version reflects the diagram of electron filling in the various allowed atomic orbitals (the Aufbau principle).  In the periodic table, there are four blocks corresponding to the types of orbitals: s, p, d, and f. Each block has as many columns as the number of electrons that can be accommodated in the corresponding sublevel: 2 columns for the s orbitals (since each s orbital can hold 2 electrons), 6 columns for the p orbitals (since each p orbital can hold 6 electrons), 10 columns for the d orbitals (since each d orbital can hold 10 electrons), 14 columns for the f orbitals (since each f orbital can hold 14 electrons). This arrangement reflects the number of available electron positions within each type of orbital, aligning with the structure of electron configurations in atoms  The periodic table is organized into groups and periods: Groups (the columns of the table) consist of elements that have the same outer electron configuration. Within each group, the elements share similar chemical characteristics due to their similar electronic structure. Periods (the rows of the table) begin with an element whose atom has an outer electron configuration consisting of one electron. As you move from left to right across a period, the atomic number (Z) increases by one with each successive element, reflecting the addition of one proton and one electron in the atom's structure. blocks The s-block: This block is formed by groups I-A and II-A (alkali metals and alkaline earth metals). Here, the s orbital is being filled with electrons. Elements in this block have their outermost electrons in the s orbitals, with two elements in each group (for a total of two electrons in the s orbital). The p-block: This block is located at the far right of the table, and it consists of the remaining six A groups. The eighth A group is often called group 0 (zero), corresponding to the noble gases. The elements in the p-block distribute six electrons into the three p orbitals. These elements include metalloids, non-metals, halogens, and noble gases. The d-block: Located in the central part of the table, the d-block consists of 10 vertical columns, forming 10 groups. These are the transition metals, elements that are filling 10 electrons into the five d orbitals. These metals are typically characterized by their ability to form multiple oxidation states. The f-block: This block is located at the bottom of the periodic table and is composed of two horizontal rows known as the lanthanide series and the actinide series. These elements are filling 14 electrons into the seven f orbitals (specifically, the 4f and 5f orbitals).  Periodic Properties The atomic radius is half the distance between the nuclei of two identical atoms that are bonded together in a molecule. It increases from top to bottom within a group and decreases from left to right across a period. Along a period, the value of Z increases, and consequently, the nuclear charge increases. The attraction exerted by the increasingly positive nucleus on the outermost electrons, all belonging to the same shell, overcomes the electrostatic repulsion between the negatively charged electrons in the same shell. Along a group, electrons occupy new shells: the effective nuclear charge increases only slightly, due to shielding from electrons in inner shells. Elements of the d-block series do not show significant variations in atomic radii with atomic number.  The nucleus is larger and exerts a greater attraction on the outer electrons, which are not significantly shielded, and as a result, the size decreases. Increase in Atomic Radius within a Group: Along a group, the principal quantum number increases, and as the principal quantum number increases, the size of the outer orbitals and thus the size of the atom also increases. Decrease in Atomic Radius across a Period: The effective nuclear charge increases (more protons are added). The added outer electrons do not provide effective shielding (unlike the electrons in inner shells). As a result, the electron cloud "contracts." The ionic radius is measured from compounds in which the ions under examination are present. The addition of electrons (anions) increases the size of the atom (ionic radius) compared to the atomic radius. The loss of electrons (cations) causes a decrease in size, due to the increased effective nuclear charge and reduced repulsive effects between electrons. Ionic radius > Atomic radius > Cation radius As the charge increases, its effect on the radius also increases. In other words, adding electrons to form an anion increases the ionic radius, while removing electrons to form a cation decreases the ionic radius. Additionally, the greater the charge of the ion, the more significant the effect on the size of the ion. The consequence of the increased attraction between the nucleus and the remaining electrons is that the cation has a much smaller ionic radius compared to the atomic radius. For anions, the reverse is true. An increase in the electron cloud results in a decreased attraction between the nucleus and the electrons. Therefore, the anion has a much larger ionic radius compared to the atomic radius.  The first ionization energy is defined as the energy required to remove an electron from its position in a neutral isolated atom in the gas phase and move it to an infinite distance from the nucleus: This energy is measurable experimentally and is also a very useful parameter for investigating the energies of the electronic levels of an atom. It is also possible to measure the second, third, and subsequent ionization energies, as well as the energies associated with the subsequent ionization reactions. (Second ionization energy) (Third ionization energy) Increase in First Ionization Energy Across a Period: Across a period, the outermost electrons of the atom are in the same energy level but experience an increasingly effective nuclear charge. The energy required to remove an electron will be higher for elements on the right side of the period compared to those on the left. Decrease in First Ionization Energy Down a Group: As the size of the atom increases, the first ionization energy decreases because the distance from the positive nucleus increases. This reduces the electrostatic force that attracts the electron and decreases the effective nuclear charge.  Electron Affinity is the energy change associated with the acquisition of an electron by an atom in the gas phase: This is also an experimental measurement, but it can be either positive or negative (energy released or absorbed). A high electron affinity means that the process of forming a negative ion (gaining an electron) is energetically favorable. Increase in Electron Affinity Across a Period: The acquisition of an electron is favored by the increase in the effective nuclear charge experienced by the outer electrons. Electron Affinity Down a Group (small differences): As you move down a group, the electron experiences less attraction from the nucleus.  For two atoms connected by a chemical bond, electronegativity describes the ability of an atom to compete for the bonding electrons. There are several scales for measuring electronegativity, one of which is the Pauling scale, which is based on the values of first ionization energy and electron affinity. Most electronegative element: Fluorine (F) Least electronegative element: Cesium (Cs)  Based on their physical and chemical properties, elements can be classified into metals, non-metals, and metalloids. Non-metals are characterized by: High ionization energies Negative and large electron affinities High electronegativity As a result, they tend to gain electrons, forming monoatomic anions and oxyanions: Cl⁻, Br⁻, S²⁻ NO₃⁻, SO₄²⁻, ClO₄⁻ Metals are characterized by: Low ionization energies Small or positive electron affinities Low electronegativity As a result, they tend to lose their valence electrons, forming cations: Na⁺, Ca²⁺, Al³⁺ 1.Metals: Characteristics: Good conductors of heat and electricity, malleable, ductile, and have a shiny appearance. They tend to lose electrons to form cations. Examples: Iron (Fe), Copper (Cu), Gold (Au). 2.Alkali Metals (Group 1): Characteristics: Highly reactive, especially with water, soft, and have low melting points. They have one electron in their outer shell, which they readily lose to form +1 cations. Examples: Lithium (Li), Sodium (Na), Potassium (K). 3.Alkaline Earth Metals (Group 2): Characteristics: Reactive (but less so than alkali metals), harder, with higher melting points. They have two electrons in their outer shell, which they lose to form +2 cations. Examples: Magnesium (Mg), Calcium (Ca), Barium (Ba). 4.Transition Metals (Groups 3-12): Characteristics: Have multiple oxidation states, good conductors of heat and electricity, and are typically hard and have high melting points. They often form colorful compounds. Examples: Iron (Fe), Copper (Cu), Zinc (Zn), Gold (Au). 5.Halogens (Group 17): Characteristics: Highly reactive non-metals, especially with alkali metals. They have seven electrons in their outer shell and readily gain one electron to form -1 anions. Examples: Fluorine (F), Chlorine (Cl), Iodine (I). 6.Noble Gases (Group 18): Characteristics: Inert gases that are chemically stable due to having a full outer electron shell. They are colorless, odorless, and do not easily form compounds. Examples: Helium (He), Neon (Ne), Argon (Ar). inprotected.com