Prof Gilbert LECTURE 20 CHEM 1211 10 27 10 Last time Chapter 7 Energy levels inside atoms ground state atom hass the lowest energy its electrons are in the lowest energy orbitals available if not the atom is in an excited state Calculating energy changes in H atoms Matter waves high speed electrons can behave like waves just like photons waves of electromagnetic radiation Wave like nature of electrons explains stability of electrons orbiting nuclei Quantum mechanics equations called wave functions define the orbitals that hold electrons each orbital is defined by a set of 3 quantum numbers Allowed combinations of Q N values This time Quantum Numbers The three kinds of quantum numbers are 1 2 3 Principal quantum number n which indicates the relative energy of an orbital and its distance to the nucleus that is its electron shell Angular momentum quantum number l indicates the shape of an orbital Orbital designations l 0 1 2 3 s p d f Magnetic quantum number ml indicates the orientation of an orbital Orbital shapes Shown by balloons that surround the region within which the electrons are likely to be 90 of the time s orbitals are spherical and so have no orientation ml 0 p orbitals have shape of two party balloons tied together with node in between There are 3 of them each consisting to 2 lobes aligned along the x y and z axes in each shell with n 1 There are 5 d orbitals in 3rd and higher shells with these designations and shapes Four of the 5 d orbitals in a subshell have 4 lobes shaped Like fat clover leaves They penetrate toward the nucleus even less than the p orbitals KEY POINTS The properties of the elements are linked to how electrons are distributed in the orbitals of their atoms Lowest energy orbitals always fill first Each orbital can hold 2 electrons Electron distributions are represented by orbital diagrams or by electron configurations In orbital diagrams individual electrons are represented by half arrows pointed up or down to indicate the orientation of each electron s spin Within an orbital electron spins are paired one up and one down No electrons are paired in a subshell until each orbital is half filled Sequence of filling orbitals in a shell is s first then if available p d and f This is the order in which their electrons can penetrate close to the nucleus where they are less shielded by inner shell electrons and so feel more effective nuclear charge More interaction with the positive nucleus means lower energy Within an orbital electron spins are paired one up ms and one down ms where ms is the spin magnetic quantum number No electrons are paired in a subshell until each orbital is half filled Hund s rule Orbital filling pattern lowest values of n and l first however there are some detours Example after Ar 1s22s22p63s23p6 the 19th electron in a K atom is in the 4s orbital not 3d Why ANSWER Energy levels are closer together as n increases so that the energies of electrons in 3d orbitals are slightly higher than those in 4s The 4s orbital fills first Therefore at the beginning of the 4th row of the periodic table Only after the 4s orbital is filled do the 3d orbitals begin to fill K Ca 1s22s22p63s23p64s1 1s22s22p63s23p64s2 Sc 1s22s22p63s23p64s23d1 Electron configuration shortcut replace filled shells with symbol of the next largest noble gas in brackets so that for Sc we have Ar 4s23d1 The shortcut emphasizes the importance of the valence shell electrons in predicting chemical properties Important trend the similarities in chemical properties among elements in the same group is reflected in similar electron configurations in their outermost valence shells and subshells KEY POINTS Chemical stability comes with 1 2 Filled valence shell s and p orbitals Group 8A noble gases Half filled p and d subshells Cr is Ar 4s13d5 Transition metals d orbitals are the last filled but their s electrons are the first to go when atoms form positively charged ions Periodic trends in the properties of the elements 1 Atomic Radii 2 Ionic Radii 3 Ionization energies the energy needed to remove 1 mole of electrons from a mole of free gas phase atoms in their ground states