1. Electron Configuration Atomic Models and Energy Levels

2. Models of the Atom-History in Models  Democritus (450BC, GR): smallest bit that is still that element.  Joseph Proust (1794 FR): proposed that compounds form in definite proportions. Proust’s experiments and law influenced Dalton.  Dalton (1808 UK): the atom is a solid ball of the element. Know the five parts of his theory.  Faraday (1830 UK): the atom has charged pieces-magnetic field affected the stream of particles in a cathode ray tube.  JJ Thomson (1903, UK): the atom is positively charged with negative bits embedded in it (plum pudding). Cathode rays are charged particles. Science Museum, UK Univ. of Oregon

3. Atomic Models 20 th Century  Rutherford (1911 NZ): the gold experiment demonstrated that most of the atom is empty space. It is a massive tiny nucleus with electrons in a cloud.  Niels Bohr (1913 DK): applied quantum theory of Einstein and others to atoms.  James Chadwick (1932 UK): proved the existence of the neutron with its slightly greater mass. Nucleus has both protons and neutrons. Bohr Model Updated Model This includes 1932 addition of neutrons

4. How to Know the Number of Electrons  Subatomic particle Atoms are composed of protons, neutrons and electrons.  The atomic number (Z) is equal to the number of protons.  The number of protons = number of electrons for neutral atom.  The charge of the atom or ion is equal to the number of protons – the number of electrons.  The number of electrons can be determined by using the atomic number and charge: number of electrons (negative charges) + ion charge

5. Energy levels, subshells, and orbitals, and spin: Quantum Numbers  Electrons are associated with energy, more specifically quantum energy, and exemplify wave-like and particle-like characteristics.  Quantum: the magnitude of a physical phenomena takes on only discrete values  The electrons of a ground state electron are in the lowest principal energy levels possible.  Electrons in an orbital that is higher than ground state are excited electrons.  The electrons exist in specific energy levels, not a continuous range.  The principal energy level, n, which is numbered 1, 2, 3, 4, …

6. Excited Electrons  When excited electrons drop back down to ground state, a photon of light is emitted.  The drop between two orbits has a specific wavelength or color.  Each element has a unique pattern, a signature.  This is used in mass spectrometry and in astronomy to determine the elements present in a sample. The visible spectrum of light and emission line spectra of hydrogen, neon, and iron. Note that the heavier an element is, the more spectral bands it has.

7. Subshells  Electrons are placed in energy levels.  These energy levels are sub-divided into subshells (labeled s, p, d or f ).  The s subshells is the lowest energy and begins in level 1.  The p subshells is higher energy and therefore doesn’t begin until level 2.  The d is higher energy and begins in level 3 and the f is even higher energy and begins in level 4.  The subshells are further sub-divided into orbitals (s has 1 orbital, p has 3 orbitals, d has 5 orbitals and f has 7 orbitals).  Each orbital can hold 2 electrons.

8. Valence Electrons: Element properties  The location of valence electrons that are available to participate in reactions can be predictive of properties and behavior of an element.  S orbital electrons—group 1 and group 2, the alkali metals and alkaline earth metals, have valence electrons in the s orbital.  The principal quantum number “n” fills the s orbital.  The maximum is two electrons for any one s orbital.  What are the ionic charges for elements of these groups?

9. P orbitals  The p block on the periodic table includes groups 13 through 18, excepting Helium which does not have occupied p orbitals.  There is a maximum of six electrons filling the three orbital shapes with a maximum of two electrons each.  Think of the shape of the orbital as a map of where an electron is most likely to be, a map of probability. It is not a route that the particle flies along.

10. d Orbitals  These are in groups 3 through 12, transition elements.  The Romans. This diverse group of metals have variable oxidation states with the exception of Ag, Zn, and usually Cd.  Five orbital shapes  10 electrons

11. f Orbitals  No group number.  The Actinides. These larger atoms exhibit some unusual properties (Neodymium and Praseodymium)  The Lantanides. These are the heavy elements which are unstable and radioactive.  Seven orbital shapes  14 electrons

12. Periodic Table and Orbitals

13. Aufbau Principle  Rules for writing electron configurations The Aufbau principle states that energy levels must be filled from the lowest to the highest and you may not move on to the next level unless the previous level is full. Use the periodic table as a guide (read left to right):

14. Aufbau Diagram (Diagonals)  It is important to keep in mind that the Aufbau principle represents an approximate trend that holds in most cases. There are however exceptions to these rules.  Why is gold, a noble metal, that is positioned in the middle of the transition elements?  Exceptions tend to be where shells can be filled (or half-filled- one electron in each orbital).

15. List the Levels and Sublevels in Order  Start from the lowest energy level to the highest.  Use the periodic table as a guide.

16. Energy Levels  Energy increases as electron positions move away from the nucleus.

17. Hund’s Rule: or I want my own orbit.  Hund’s Rule says that when placing electrons in orbitals of equal energy, place one in each orbital before doubling up in order to arrive at the lowest energy configuration.  The Pauli Exclusion Principle states that when electrons do share an orbital, they must be of different “spin.” No two electrons can have the same quantum numbers: each one has an unique number address.

18. The Elements  Each time an electron is added it goes to an open orbital in the same energy level FIRST.  In larger atoms, both the s and p may have one electron only to get a configuration with one electron in every orbital. See Cr and Cu.  This creates a more balanced geometry.

19. Periodic Table

20. Lewis Dot Table (s and p orbitals combined) s 1 s 2 p 1 p 2 p 3 p 4 p 5 p 6 Noble: complete shells

21. Lewis Dot with Levels  As electrons are added, they fill electron shells in an order determined by which configuration will give the lowest possible energy. The first shell (n=1) can have only 2 electrons, so that shell is filled in helium, the first noble gas.  In the periodic table, the elements are placed in "periods" and arranged left to right in the order of filling of electrons in the outer shell. So hydrogen and helium complete the first period.

22. Quantum Numbers  Principal quantum number (n) is the shell (energy level) that an electron belongs to.  If an electron is in the lowest possible energy state, then it is in its ground state.  If an electron is in a higher energy state, then it is in an excited state.  The total number of orbitals for n = n 2. 1 st level has 1*1 orbitals. 2 nd level has 2*2 orbitals or 4 total.  n may equal 1, 2, 3, 4, 5, 6, …  Angular momentum (secondary, azimuthal) quantum number ( l ) specifies the shape of the orbit or subshell.  l 01234  Letter spdfg

23. Quantum Numbers  Magnetic quantum number (m l ) specifies the orientation in space.  For s orbitals, l is zero. For s orbitals, m l is zero  For p orbitals, l is -1, zero, +1 (orientation to x, y, or z axis)  For d orbitals, l is -2, -1, zero, +1, +2 (five orientations)  For f orbitals, l is -3, -2, -1, zero, +1, +2, +3 (seven orientations)  m l = - l, …, 0, …, l  Spin quantum number (m s ) specifies one of two possible spin orientations within a single orbital.  This is often referred to as up or down. Pauli Exclusion Principle states that no two electrons can have exactly the same four values for their quantum numbers. Two electrons in the same orbital must have opposite spins.

24. Quantum Numbers (1 st Three)  n l m l Number of Orbital Number of (energy) (orientation) orbitals Name electrons (shape) 1 0 0 1 1s 2 2 0 0 1 2s 2 1 -1, 0, +1 3 2p 6 3 0 0 1 3s 2 1 -1, 0, +1 3 3p 6 2 -2, -1, 0, +1, +2 5 3d 10 4 0 0 1 4s 2 1 -1, 0, +1 3 4p 6 2 -2, -1, 0, +1, +2 5 4d 10 3 -3, -2, -1, 0, +1, +2, +3 7 4f 14  Spin (m s )is written as -½ or +½.

25. Resources  http://chemwiki.ucdavis.edu/Inorganic_Chemistry http://chemwiki.ucdavis.edu/Inorganic_Chemistry  http://www.colorado.edu/physics/2000/quantumzone/bohr.html Bohr model animation with photon interactions. http://www.colorado.edu/physics/2000/quantumzone/bohr.html