Fundamentals of Medicinal and Pharmaceutical Chemistry PDF

Fundamentals of Medicinal and Pharmaceutical Chemistry FUNCHEM.2 Electronic Configurations of Atoms and Ions of Physiological Importance D r. D a r r e n G r i ffi t h  General...

Fundamentals of Medicinal and Pharmaceutical Chemistry FUNCHEM.2 Electronic Configurations of Atoms and Ions of Physiological Importance D r. D a r r e n G r i ffi t h  General Chemistry - The Essential Concepts by Chang Recommended and Goldsby 7e reading  Sections 2.1, 2.2, 2.5, 7.8, 7.9 F U N C H E M. 2 E l e c t r o n i c C o n fi g u r a t i o n s o f A t o m s a n d I o n s o f P h y s i o l o g i c a l I m p o r t a n c e 2 FUNCHEM.2 Learning Outcomes Recall contributions by Thompson, Rutherford, Bohr, Einstein, deBroglie, Heisenberg, Schrodinger leading to the discovery of the structure of the atom. Define ‘atomic orbitals’ and be able to differentiate between the energy and shapes of the s, p and d orbitals. Define ‘Pauli Exclusion Principle’, ‘Aufbau Principle’ and ‘Hund’s Rule of Maximum Multiplicity’. Construct electronic configurations and abbreviated electronic configurations for the first 30 elements of the Periodic Table. Construct electronic configurations for simple cations and anions. 3 Four main points of Dalton's Atomic Theory Elements are composed of extremely small particles, called atoms. All atoms of a given element are identical, having the same size, mass, and chemical properties. The atoms of one element are different from the atoms of all other elements. Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction. A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction John Dalton 1766 – 1844 4 Early Atomic Structure - Thomson Credited with the discovery and identification of the first subatomic particle - the electron Using a cathode ray tubes and knowledge of electromagnetic theory, he determined the ratio of electric charge to the mass of an individual electron. J.J. Thomson −1.76 × 108 C/g, where C stands for coulomb, which is the unit of electric charge Nobel Prize 1906 Watch the animation titled ‘Cathode Ray Tube’ on the VLE 5 Early Atomic Structure - Thomson Matter is normally neutral. To maintain electrical neutrality, an atom must contain an equal number of positive and negative charges Must be another positive bit to equal the negative charge on the proposed electrons. Plum-pudding model proposed. J.J. Thomson Electrons are embedded in a uniform, positively Nobel Prize 1906 charged sphere. 6 Early Atomic Structure - Rutherford Alpha (α) particles are positively charged ‘This was as unexpected as if they had shot a bullet at a piece of tissue paper, and the bullet had bounced back’ The nucleus was born! Rutherford proposed, that the atom's positive charges are all concentrated in the nucleus, a dense central core within the atom. The positively charged particles in the nucleus are called protons. Rutherford 1910 Nobel Prize in Chemistry 1908 Watch the animation titled ‘Alpha-Particle Scattering’ on the VLE 7 Early Atomic Structure - Bohr How are electrons arranged outside the nucleus of an atom? Neils Bohr Nobel Prize 1922 Neils Bohr after studying the emission spectra of hydrogen proposed: (i) that electrons move in circular orbits around the nucleus (ii) the orbits have specific energies – i.e. energy levels 8 Photoelectric Effect - Einstein Albert Einstein (1879-1955) Nobel Prize 1921 Demonstrated that light, although a wave, could also have particle properties 9 Wave Particle Duality – de Broglie ‘If Einstein says that light (which we all knew was a wave) is really particles, then maybe all the things we think of as particles are really waves (or can behave as if they were waves!)’ Matter (small particles), as well as light, might have wave properties!!! deBroglie 1924 Therefore electrons, like light, also might have particle properties!! WAVE-PARTICLE DUALITY 10 Uncertainty Principle - Heisenberg "The more precisely the POSITION is determined, the less precisely the MOMENTUM is known" It is impossible to know simultaneously both the momentum (mass times velocity) and the position of a particle (such as an electron) with certainty 11 Schrodinger Wave Equation Schrodinger 1925 Developed mathematical functions which are solved by the Schrodinger Equation which refer to the probability of finding an electron within a particular position within an atom 12 Solvay Conference 1927 13 Solvay Conference 1927 14 Atomic Orbitals Knowing how electrons are arranged in an atom is very important as it helps us to understand how atoms interact with each other and even predict how atoms might interact with each other An atomic orbital is a region in space within an atom and around the nucleus where the probability of finding an electron is relatively high There are different types of atomic orbitals Please note we will not use quantum numbers to describe atomic orbitals and therefore have excluded section 7.7 from the recommended reading 15 Electronic Configurations Atomic orbitals: are available in different sizes and energies can have particular shapes have different orientation If required several of the same type of orbital can be used by the same nucleus to carry its electrons. 16 Atomic Orbitals – s orbitals The nucleus of the The 3 axes represent atom is at 3-dimensional space the center of the three axes. The s-orbital is sphere shaped 17 Atomic Orbitals – s orbitals There are different types of s-orbitals Increasing number, therefore increasing energy and distance form the nucleus 18 Atomic Orbitals – p orbitals p-orbitals have different shapes than s-orbitals The 3 axes represent 3-dimensional space p orbital The p-orbital is dumbbell shaped 19 Atomic Orbitals – p orbitals p-orbitals occur in sets of three with: the same shape the same energy (degenerate). but different orientation px py pz 20 Atomic Orbitals – d orbitals d-orbitals have different shapes than s- and p-orbitals d-orbitals occur in sets of five with their own unique shapes The five d-orbitals have the same energy (degenerate). dx2-y2 dz 2 dxy dxz dyz https://www.youtube.com/watch?v=K-jNgq16jE Y 21 Energy Sublevels An energy sublevel is a group of atomic orbitals within an atom, all of which have the same energy For example the px, py and pz orbitals have the same energy Maximum of 2 electrons per orbital No. of Sublevel orbitals Maximum number of electrons s 1 2 p 3 6 d 5 10 22 Electronic Configurations Describes how electrons are organised in an atom – how electrons are distributed among the atomic orbitals and sublevels Electrons are assigned to orbitals and sublevels orbitals and sublevels have capacity and energy 23 Electronic Configurations 1. Pauli Exclusion Principle No more than 2 electrons can occupy an orbital – electrons must have opposite spin (electron spin) 2. Aufbau Principle Electrons fill the lowest available energy level first 3. Hunds Rule With orbitals of equal energy (px, py, pz), electrons fill the orbitals singly before filling in pairs. 24 Electronic Configurations spd(f) notation uses a prefix to designate energy and size of the orbital – the principal quantum number, recall the higher the number the further from the nucleus and generally the higher the energy Superscript indicates the number of electrons in the orbital Empty orbitals are not included prefix superscript 1s2 pronounced "one s two" 25 Electronic Configurations 1s 2s and 2p 3s, 3p and 3d 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p 26 Electronic Configurations Order of Filling Orbitals and Energy Sublevels 27 Electronic Configurations Examples Atomic orbital diagrams* Electronic Configurations 1 H 1s1 2 He 1s2 3 Li 1s2 2s1 4 Be 1s2 2s2 5 B 1s2 2s2 2p1 6 C 1s2 2s2 2p2 7 N 1s2 2s2 2p3 * Boxes represent orbitals and arrows electrons 28 Electronic Configurations Examples 8 O 1s2 2s2 2p4 9 F 1s2 2s2 2p5 10 Ne 1s2 2s2 2p6 11 Na 1s2 2s2 2p6 3s1 outermost e- called 11 Na : [Ne] 3s1 valence electron 29 Electronic Configurations In the same way: 18 Ar 20 Ca Noble Gas electron Ca: [Ar] 4s2 configuration 20 30 Transition Metals – d metals Remember: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p Energy Atomic orbital diagram - 21Sc Electronic Configuration - 21Sc 1s2 2s2 2p6 3s2 3p6 4s2 3d1 31 Transition Metals – d metals 22 Ti 1s2 2s2 2p6 3s2 3p6 4s2 3d2 27 Co 1s2 2s2 2p6 3s2 3p6 4s2 3d7 http://www.webelements.com/webelements/elements/text/periodic-table/econ.html Watch the animation titled ‘Electronic Configuratuions’ on the VLE 32 Transition Metals – d metals 24 Cr Expect: 1s2 2s2 2p6 3s2 3p6 4s2 3d4 Actually: 1s2 2s2 2p6 3s2 3p6 4s1 3d5 Half-filled Please note that we will not cover electron configurations of lanthanides or rare earth series 33 Transition Metals – d metals 29 Cu Expect: 1s2 2s2 2p6 3s2 3p6 4s2 3d9 Actually: 1s2 2s2 2p6 3s2 3p6 4s1 3d10 Completely-filled Cr and Cu are exceptions as there is extra stability associated with either a half-filled or completely filled d orbital 34 Formation of Ions An ion is an atom or a group of atoms that has a net positive or negative charge. The loss of one or more electrons from a neutral atom results in a cation, an ion with a net positive charge. an anion is an ion whose net charge is negative due to an increase in the number of electrons Nucleus not affected 35 The main positive ion (cation) in blood Na+ How is Na+ formed? 11 Na 1s2 2s2 2p6 3s1 Loss of one electron 11+ 11- 11 Na+ 1s2 2s2 2p6 (11+) + (10-) = +1 hence Na+ 36 The main negative ion (anion) in blood Cl- How is Cl formed? - 17 Cl Gains one electron 17+ 17- (17+) +(18-) = -1 hence Cl- 37 Try K+: the main cation in cell fluid 19 K 19 protons 19 electrons 19 K+ must have lost an electron 19 protons 18 electrons 19 K+ 1s2 2s2 2p6 3s2 3p6 38 FUNCHEM.2 Learning Outcomes Recall contributions by Thompson, Rutherford, Bohr, Einstein, deBroglie, Heisenberg, Schrodinger leading to the discovery of the structure of the atom. Define ‘atomic orbitals’ and be able to differentiate between the energy and shapes of the s, p and d orbitals. Define ‘Pauli Exclusion Principle’, ‘Aufbau Principle’ and ‘Hund’s Rule of Maximum Multiplicity’. Construct electronic configurations and abbreviated electronic configurations for the first 30 elements of the Periodic Table. Construct electronic configurations for simple cations and anions. 39 Thank you F O R M O R E I N F O R M AT I O N P L E A S E C O N TA N T D r. D a r r e n G r i ffi t h E M A I L : d g r i ffi t h @ R C S I. C O M 40