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- 1. PHY/CHEM ATOMIC STRUCTURE AND THE PERIODIC TABLE PHARMACEUTICAL CHEMISTRY I Dr. Kato Lodrick mps
- 2. ATOMIC STRUCTURE An atom is a basic building block of matter. • Contain negatively charged electrons, positively charged protons and neutral particles called neutrons. Electron arrangement • Electrons move around the nucleus and are arranged in shells at increasing distances from the nucleus. • Some electrons are closer to the nucleus than others • The different areas of an electron in an atom are called energy levels. • Each energy level can hold a maximum number of electrons • The further away the nucleus, the more the energy level can hold
- 3. Energy Level/shells Number of electrons 1 2 2 8 3 18 4 32
- 4. Bohr model •
- 5. Cont. • Bohr atomic model is the simplest, in which atoms are assumed to be positioned around the nucleus in discrete orbitals. • Bohr model describe electrons in terms of their positions (orbitals) and energy Limitations of the Bohr model? Other atomic models? Differences?
- 6. cont • An Orbital is a region in space that can hold 2 electrons. • electrons do not move freely around the nucleus, but are confined to regions of space called shells. • Each Shell contains (2n2) electrons. 1st shell has 1s orbital 2nd shell has one 2s and three 2p orbitals 3rd shell one 3s, three 3p, and five 3d orbitals 4th shell has one 4s, three 4p, five 4d and seven 3f orbitals
- 7. Periodic table and electronic configuration
- 8. Cont. • Elements are classified according to electron configuration in this table • Elements are placed with increasing atomic number in 7 horizontal rows called periods. • Elements in the column or group have similar valency electron configurations, as a result similar properties. Valence is the number of electrons an atom must gain or lose to gain the nearest noble or inert gas electronic configuration Group 0- inert gasses (filled electron shells) Group VIIA- halogens, one electron deficient and Group VIA , 2 electrons deficient from having stable configurations Group1A and 11A are alkali (except H) & alkali earth metals, with 1 and 2 excessive electrons respectively from stable configurations
- 9. Cont.. • Elements in three long periods, Groups from IIIB to IIB are transitional metals, with partially filled d electron states and in some cases one or two electrons in the next higher shell • Elements in Group IIA, IVA and VA have characteristics between metal and non metals because of their valence electron configurations Ground state electronic configurations of some elements First period Second period Third period H 1 1S1 He 2 1S2 Li 3 2S1 Be 4 2S2 B 5 2s2 2P1 Na 11 3S1 Mg 12 3S2 Al 13 3S2, 3P1 Si 14 3S2, 3P2
- 10. 10 • Why? Valence (outer) shell usually not filled completely. • Most elements: Electron configuration not stable. SURVEY OF ELEMENTS Electron configuration (stable) ... ... 1s22s22p 63s23p 6 (stable) ... 1s22s22p 63s23p 63d 10 4s24p6 (stable) Atomic # 18 ... 36 Element 1s1 1 Hydrogen 1s2 2 Helium 1s22s1 3 Lithium 1s22s2 4 Beryllium 1s22s22p 1 5 Boron 1s22s22p 2 6 Carbon ... 1s22s22p 6 (stable) 10 Neon 1s22s22p 63s1 11 Sodium 1s22s22p 63s2 12 Magnesium 1s22s22p 63s23p 1 13 Aluminum ... Argon ... Krypton Adapted from Table 2.2, Callister & Rethwisch 8e. • Why?
- 11. 11 Electron Configurations • Valence electrons – those in unfilled shells • Filled shells more stable • Valence electrons are most available for bonding and tend to control the chemical properties • example: C (atomic number = 6) 1s2 2s2 2p2 valence electrons
- 12. 12 Electronic Configurations ex: Fe - atomic # = 26 valence electrons Adapted from Fig. 2.4, Callister & Rethwisch 8e. 1s 2s 2p K-shell n = 1 L-shell n = 2 3s 3p M-shell n = 3 3d 4s 4p 4d Energy N-shell n = 4 1s2 2s2 2p6 3s2 3p6 3d6 4s2