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    0 introductory recapitulation 0 introductory recapitulation Presentation Transcript

    • Medical Chemistry Introductory recapitulation to lectures 2006 (J.S.) The basic chemical terms Electron configuration in atoms Chemical bonds and weak intermolecular interactions
    • Material objects (bodies, systems) Substances are the material contents of material objects Heterogeneous matter (multiphase systems) M ixtures Pure substances CHEMICAL INDIVIDUALS Homogenous mixtures SOLUTIONS Gaseous mixtures Alloys Glasses COMPOUNDS ELEMENTS NUCLIDES Homogenous samples of matter
    • Properties of single particles (atoms, molecules, ions, etc.) : Mass m Electric charge of an ion - the number of elementary charges The actual mass of this standard of reference m ( ) = 1.66 × 10 –24 g is also taken as 1 atomic mass unit (1 u). The crucial quantity in every chemical reaction is the number of atoms or molecules involved in the reaction. A r ( ) = 12.000 Because it is not only inconvenient but impractical to express the actual masses of single particles in grams, the relative atomic mass scale have been devised: the mass of a carbon-12 atom was chosen as the standard of reference for atomic as well as molecular masses M r . A r ( ) = 1.000
    • The properties of substances that are measureable in reality contain extremely large numbers of single atoms or other particles. Amount of substance (or substance amount, too) is the quantity proportional to the number of defined entities (atoms, molecules, ions, electrons, and other particles) comprising any system. The symbol for this quantity is n . The unit of amount of substance is the mole . The symbol for the mole is mol .
    • The definition of the unit amount of substance: One mole is an amount of a substance containing the same number of particles (defined entities as atoms, molecules, ions, formula units, electrons, etc.) as there are atoms in carbon with the mass exactly 12.000 g , i.e. approximately 6.022  10 23 . The number of particles comprising one mole of particles is the Avogadro ´s constant N A . N A = ( 6.022 136 7 ± 0.000 003 6 )  10 23 mol –1
    • The quantities related to the unit substance amount and describing so the properties of 1 mole of particular substance (or particles) are molar quantities . Be careful to note the distinction: Relative molecular or atomic mass is a ratio indicating the mass of a single particle relative to the mass of the 1/12 of carbon - 12 atom. Molar mass is the mass of 1 mole of those particles expressed in grams , in spite of both numerical values are the same. Molar mass M is the mass of 1 mole of a specified substance or particles. When the unit g mol –1 is used for a molar mass, then its numerical value is equal to the relative mass of the specified entity (A r , M r ) .
    • Avogadro ´s Law : Equal volumes of different gases at the same temperature and pressure contain the same number of molecules. Don't forget that the volume of gases increases with increasing temperature. If the pressure remains constant, V 2 = V 1  T 2 / T 1 . E.g., at 20  C and 200 m above sea level the molar volume V m equals 24.733 litres per mole. Molar volume of gases V m According to Avogadro ´s law, at the same temperature and pressure, one mole of any gas occupy the same volume as one mole of any other gas. V m,STP = 22.414 l mol –1 ( t = 0  C ; p = 101.3 kPa) Standard molar volume V m,STP is defined as the volume occupied by one mole of gases that are close to ideal at standard temperature and pressure (STP).
    • Molar electric charge The amount of electric charge represented by 1 mol of electrons (also protons or ions with one elemental charge, both negative and positive) is the Faraday's constant F F = 96 485.3 C mol –1 which for most calculations can be rounded to 96 500 C mol –1 . Multiples of this quantity have to be used for molar charges Q m of ions with two and more elemental charges. The amount of elemental electric charge can be obtained from F : q = F / N A = 1.602  10 –19 C
    • Electron configuration in atoms
    • Each atom consists of a nucleus surrounded by electrons. The nucleus contains positively charged protons and neutral neutrons . In a neutral atom, the positive charge of the nucleus is exactly set off by the negative electrons. A neutral atom must contain the same number of electrons as protons. This number is the proton number (also the atomic number ) of an element. Electrons are circling around the nucleus forming so an electron casing of the nucleus. Most properties of the elements and their compounds that are of interest in chemistry are intimately related to the arrangement of electrons in atoms or molecules
    • Electrons are particles that exhibit the particle-wave duality. They can be viewed as both particles and quantums of electromagnetic waves. However, according to the Heisenberg uncertainty principle, it is impossible to know simultaneously both the exact angular momentum and the exact position of an electron (or any particle that exhibits quantum-mechanical behaviour). The Bohr model of the hydrogen atom Electrons are viewed as particles circling around the nuclei in fixed orbits. The Bohr model was based on the assumptions (postulates) that – the electron is moving around the nucleus only in specific orbits, in which its energy is quantized (i.e. can have only specific values – energy states, energy levels); – the electron in the ground state is in the lowest energy level; when the atom absorbs energy, it becomes excited – the electron jumps to a higher energy level and falling to lower energy, the extra energy is emitted as electromagnetic radiation. These two of Bohr´s ideas remain as part of modern atomic theory.
    • The quantum-mechanical model of the atom In this model, electrons are not considered to be only particles revolving around the nucleus in orbits, clearly defined paths of motion. Electrons occupy atomic orbitals that are filled with an “electron cloud” (electron waves which share the same mathematical expression as standing waves). The four quantum numbers – n ,  , m , and s – specify the energy and probable location of each electron in an atom. An orbital is a region in space around the nucleus within which an electron of a given energy is most likely to be found. Electrons of different energies are found in different orbitals.
    • 1 Electrons exist in energy levels at different distances from the nucleus. The principal quantum number n indicates the energy levels of the electrons relative to their distance from the nucleus. The number n may have any positive integral value up to infinity, but only values of 1 – 7 have been established for atoms of known elements in their ground state (lowest energy state). The maximum number of electrons in each principal energy level is 2n 2 . 2 Electrons in the principal energy levels (except the first) exist in a number of closely grouped sublevels (subshells, the number of sublevels is equal to n) that consist of orbitals of specific shape. The azimuthal quantum number  (the subshell number) specifies the shape of the orbital . It takes values from 0 up to (n – 1). As a rule the sublevels (orbitals) are designated by the letters s (for  = 0), p, d, and f (for  = 3). Within a specific principal energy level, the energy of electrons of the sublevels increases slightly from s to f.
    • 3 Electron orbitals have specific orientation in space. The magnetic quantum number m designates the spatial orientation of orbitals . This number can take only the integral values from 0 to ±  so that there can be at most one s orbital (m = 0), three p orbitals (m –1, 0, +1), five d orbitals (m –2, –1, 0, +1, +2), five d orbitals, and seven f orbitals in any given principal energy level. The orbitals of the same subshell in the given principal energ level that differ only in magnetic quantum numbers m are called degenerated orbitals. 4 Each orbital can be occupied by two electrons that differ in their spin. The spin quantum number s relates to the direction of spin of an electron . There are only two possible values of spin : + ½ and –½. A maximum of two electrons can occupy each orbital, no matter what its designation. When an orbital contains two electrons differing in their spin, the electrons are said to be paired.
    • Sublevel electrons in each principal energy level (electron shell) and the maximum number of orbitals and electrons in each energy level
    • Perspective representation of the s, p, and d orbitals The three 2p orbitals The five 3d orbitals An s orbital
    • Order of filling electron orbitals The lowest energy principle (the "building-up" principle) Electrons occupy the lowest energy orbitals available to them; they enter higher energy orbitals only when the lower energy orbitals are filled. With increasing proton number, electrons occupy the subshells available in each principal energy level in the following order : 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d ……. This sequence can be read from the periodic table.
    • Orbital filling sequence in the periodic table Take notice of energy overlaps: 4s and 5s occupied before 3d and 4d, 5d 1 and 6d 1 before 4f and 5f.
    • No two electrons can have the same four quantum numbers (n, l , m, s) – each electron must have a unique set of quantum numbers. The Pauli exclusion principle Hund´s rule ( the rule of maximum multiplicity ) In the lowest energy state (in the ground state), electrons occupy orbitals of given n values (degenerated orbitals) in such a way as to have a maximum number of unpaired electrons.
    • Writing electron configurations Electron configuration of phosphorus may serve as an example: Proton number 15 - 15 electrons surrounding the nucleus. The complete configuration: A simplified notation: [ Ne ] 3s 2 3p 3 Electron-dot representation (Lewis symbol) of a phosphorus atom: Symbol of the element represents the kernel (the nucleus and all the electrons other than those in the outermost energy level. Paired valence electrons represented by a short line, dots represent unpaired valence electrons. 1s 2 2s 2 2p 6 3s 2 3p 3 or Completely filled and therefore stable electron configuration of the noble gas neon The outermost electrons available to take part in chemical bonding – the valence electrons P · · ·
    • Periodic table of the elements The periodic law : The properties of the elements are a periodic function of their proton numbers
    • Representative elements (s block and p block) d-Transition elements f-Transition elements (lanthanides and actinides)
    • Orbital radius of an atom decreases Ionization energy increases Electronegativity increases Non-metallic character increases Oxidation effects of an element increases Orbital radius of an atom increases Metallic character increases Reducing properties of an element increases Ionization energy decreases Electronegativity decreases In groups In periods
    • The electronegativity scale given in colour
    • W eak intermolecular bonding interactions Types of chemical bonds
    • The ability of an atom to form chemical bond is related to the distribution of electrons in that atom. So knowledge of electron configuration of elements is the basis of understanding chemical bonding between elements. Through chemical reactions, atoms tend to attain more stable states at lower energy levels. Atoms react by sharing, losing, or gaining valence electrons . Forces arise from electron transferring and sharing interactions. Chemical bonds are the forces that bind atoms in molecules or in polyatomic ions or that hold oppositely charged ions together. There are two principal types of chemical bonds : the covalent bond and the ionic (or electrovalent) bond .
    • A rule that accounts for the formation of many compounds is based upon the observed stability of compounds in which atoms are associated with eight valence electrons (ns 2 np 6 ). Atoms tend to combine by sharing, gain, or loss of electrons so that the outer energy level of each atom holds or shares four pairs of valence electrons (an octet of valence electrons). The octet rule: Octet of valence electrons and duplet of helium are particularly stable electron configurations. For most of the elements, octet is the electron configuration of the nearest noble gas (except for helium with the configuration 1s 2 ). Some exceptions to the octet rule exist.
    • The covalent bond The covalent bond dissociation energies (required to break a covalent bond) are in the range 150 – 550 kJ mol –1 . Covalent bonds often form so that the bonded atoms achieve octet configurations. A covalent bond represents an overlap of two half-filled atomic orbitals which have turned into a bonding molecular orbital occupied by two electrons of opposite spins. A single covalent bond is a bond in which two atoms are held together by sharing two valence electrons (one bonding electron pair).
    • Nonpolar covalent bond A nonpolar covalent bond is a bond in which the bonding electron pair is shared equally (bonds of atoms of the same element or of atoms with a small difference in electronegativity, e.g. C and H). 1s 1 1s 1 Bonds, in which the region of highest electron density surrounds the bond axis , are called σ bonds ( sigma bonds ). Only one σ bond can form between any two atoms.
    • Polar covalent bond In a polar covalent bond binding electron pair is shared unequally ; one atom (of higher electronegativity) acquires a partial negative charge (–) or δ– and the other acquires a partial positive charge (+). The entire molecule remains electrically neutral, but it is a dipole . Whether or not a molecule is a dipole depends upon molecular geometry In addition to the presence of polar bonds.
    • Single covalent bonds in polyatomic molecules In most molecules with more than two atoms, overlap of s and p orbitals (described in precious slides) cannot explain the bond angles and the observed geometry of those molecules. The concept of hybridization was introduced to explain molecular geometry in terms of atomic orbitals and valence bond theory. Hybridization applies only to the orbitals that form σ bonds of atoms in a molecule. In terms of energy, hybridization represents the blending of higher energy and lower energy orbitals to form orbitals of intermediate energy. Hybridization is the combining of the atomic orbitals of a single atom to give a new set of orbitals, called hybrid orbitals , on that atom.
    • Example: A carbon atom forms four equivalent covalent bonds in CH 4 or CCl 4 . This is accomplished by sp 3 hybridization , the combination of s and three p orbitals: The valence electrons of carbon and excited state of carbon atom:: Hybridization sp 3 : Four equivalent sp 3 orbitals Carbon atom also can be sp 2 or sp hybridized to form multiple bonds: - s and two p orbitals combine to give three hybrid orbitals sp 2 ; p z orbital unhybridized
      • s and one p orbitals combine to give
      • two hybrid sp orbitals; p z and p y unhybridized
      2s 2p 2s 2p x p y p z 2s 2p x p y p z 2 (sp 3 ) 2s 2p x p y p z 2s 2p x p y p z
    • Hybridized orbitals form σ bonds by overlap with other hybrid orbitals, and with s and p orbitals:
    • Multiple covalent bonds Only one σ bond (symmetrical about the bond axis) can form between any two atoms. Multiple covalent bonds arise when p or d orbitals on atoms that are σ-bonded to each other also overlap . A bond formed in this way is called a π bond ( pi bond ). Electron density of π bonds is concentrated above and below the bond axis and always have a plane of zero electron density passing through the bond axis.
    • Covalent bonds in ethene CH 2 =CH 2 The C–H bonds are σ bonds formed by overlap of sp 2 orbitals from C atoms and s orbitals from H atoms. The C=C double bond is formed as σ bond by sp 2 –sp 2 overlap and the π bond by p–p overlap. An sp 2 -hybridized carbon atom three sp 2 orbitals in plane can form σ bonds p x orbitals above and below plane can form π bonds C C H H H H
    • So far, covalent bonds have been considered to which each atom contributes one electron (overlap of two half-filled orbitals). Such a formation of covalent bonds is called colligation . Covalent bonds can also originate by coordination: The number of covalent bonds of electron pair donors is then greater than the typical one, for acceptors of an electron pair, however, it may be smaller. An electronegative atom with an unshared electron pair may use this filled valence orbital as an electron pair donor for forming a covalent bond by overlapping the vacant valence orbital of another atom – the electron pair acceptor . The bond is then called a dative or coordinate covalent bond . If not speaking of its origin, the coordinate bond does not differ from other covalent bonds. Ι NH 3 + H + -> NH 4 + -> + 3
    • Nonmetals have their standard numbers of covalent bonds However, the number of covalent bonds also may be higher or lower than the standard one , e.g., when the atom takes part in formation of coordinate covalent bond(s); then the atom gains an electric charge : that are equal the number of electrons required to complete the octet of valence electrons. Nitrogen Oxygen Number of bonds 4 3 2 1 0 Chlorine
    • Particle Electron formula Number of covalent bonds Oxidation number H–H H 2 CO NH 3 NH 4 + CaO HClO 4 H 1 H I C 3 O 3 C II O –II N 3 H 1 N –III H I N 4 H 1 N –III H I H 1 Cl 4 H I Cl VII O –II 3x O 1 1x O 2 Ca 0 O 0 Ca II O –II
    • Covalent compounds exhibit low polarity and exist in two types: Most of them are molecular compounds consisting of neutral molecules. They occur as gases, nonpolar liquids, and soft solids with low melting and boiling temperatures, soluble in nonpolar solvents. Some solid covalent compounds form no discrete molecules; covalent bonds join the atoms continuously throughout a three-dimensional crystalline lattice, creating a single huge molecule. These covalent network solids are very hard, with high melting temperatures (e.g., diamond, silicon dioxide SiO 2 , silicon or boron carbides SiC, B 4 C).
    • The ionic bond (the electrovalent bond) Independent molecules do not exist in ionic compounds. Each ion is surrounded by other ions of opposite charge. Chemical formulas of ionic compounds show only the simplest ratios necessary for electroneutrality – the formula units –representing only parts of the entire crystalline lattice. Ionic compounds occur typically as pieces of crystalline solid. They are hard (much harder than covalent compounds), brittle, and have high melting temperatures. Many of them are water soluble. Ionic bonding is the attraction between positive and negative ions within a crystalline lattice.
    • There is no sharp dividing line between ionic and covalent bonds. Some bonds are mostly ionic, the other mostly covalent. The ionic bond is sometimes referred to as the extremely polarized type of the covalent bond . According to Pauling, the degree of ionic character of a covalent bond can be estimated approximately from the difference of electronegativities of the atoms bonded. ΔEN = (x A – x B ). Degree of ionic character equals 1 – e exp [– 0.21  (x A – x B ) 2 ] Polar covalent bonds are bonds between elements the ΔEN of which is in the range of 0.4 – 1.7. If ΔEN < 0.4, the degree of ionic character of the covalent bond is less than 5 % - it is a nonpolar covalent bond . If ΔEN > 1.7, the degree of ionic character of the bond is more than 50 % and the bond is taken as a ionic bond .
    • The metallic bond Metals, the electropositive elements having a small number of valence electrons, form a close-packed metallic lattice in the solid state. In this lattice, cations of metals are embedded and surrounded by an electron gas (electron sea) of free valence electrons that are free to move over the metal as a whole. It may be said that all the atoms in the metallic lattice share all of their valence electrons. In transition metals some of the (n–1)d electrons can also contribute to the electron gas. General physical properties of metals: Solids of high melting point, often hard, lustrous, malleable, ductile, strong conductors of heat and electricity. The metallic bond is sometimes considered the third type of chemical bond. Metallic bonding is the attraction between positive ions and surrounding electrons that can move throughout the lattice at continuous bands of energy levels.
    • Intermolecular forces in liquids and solids Weak intermolecular forces ( non-covalent interactions ) act between molecules, causing them to be attracted to each other in varying degrees. The strength of these forces at a particular temperature determines whether a molecular substance is a gas, a liquid, or a solid at that temperature. Non-covalent forces also contribute to the secondary, tertiary, or quaternary structures of biopolymers (namely proteins a nucleic acids), maintain the stability of biomembranes, and play important roles in highly specific biological interactions (enzyme – substrate, antigen – antibody, signal molecules – receptor. Types of intermolecular forces : – hydrogen bonds , – dipole–dipole interactions , – dispersion forces (London forces). Dipole–dipole and dispersion forces are sometimes referred collectively to as van der Waals forces .
    • Hydrogen bonds are the strongest intermolecular forces (8 – 40 kJ mol –1 ). When a hydrogen atom is covalently attached to a highly electronegative atom such as fluorine, oxygen, or nitrogen, it will also be attracted by an unshared electron pair of a highly electronegative atom of another molecule, forming a bond (or bridge) between the two molecules. Examples: Water in the liquid and solid states exists as aggregates (clusters) in which the water molecules are linked together by hydrogen bonds. In proteins and nucleic acids , there are many H-bonds that stabilize the secondary structures of the macromolecules. The usual type of these H-bonds is O Ι···H–O , O Ι···H–N , NΙ···H–O , or NΙ···H–N .
    • Van der Waals forces Dipole–dipole interactions are the attractive forces between permanent dipoles (orientation forces), or between a permanent dipole and an induced dipole (induction forces; the electric field of ions or permanent dipoles can induce a weak dipole moment in polarizable molecules). Hydrophobic interactions between nonpolar molecules in aqueous solutions depend on the formation of hydrogen bonds between water molecules. Dispersion forces (London forces) act on all atoms and all molecules, whether polar or nonpolar. They are the result of momentary shifts in the symmetry of the electron cloud of a molecule.