1. Electrons in Atoms

2. Models of the Atom

3. 5.1.1 – I can identify the inadequacies in the Rutherford atomic model. 5.1.2 - I can identify the new proposal in the Bohr model of the atom. 5.1.3 - I can describe the energies and positions of electrons according to the quantum mechanical model. 5.1.4 - I can describe how the shapes of orbitals related to different sub-levels differ.

4.  Elements Elements  Rutherford’s atomic model couldn’t explain the chemical properties of elements.

5. You may need to use your book to fill this in – Pages 128-129

7.  Niels Bohr (Danish 1885- 1962) a student of Rutherford saw that his model needed improvement.  Bohr proposed that an electron is found only in specific circular paths, orbits, around the nucleus.  Energy levels – the fixed energies an electron can have.

8.  Energy levels are like steps or rungs on a ladder.  Quantum – amount of energy required to move an electron from one energy level to another energy level.  Energies between levels are not all the same

9.  Erwin Schrödinger (Austrian 1887-1961) used math to describe the behavior of the electrons.  Quantum mechanical model – modern description of electrons in atoms based on mathematical solutions to the Schrödinger equation.

10.  The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus.  Think of as a fuzzy cloud of chance.

11.  Atomic orbitals – a region of space in which there is a high probability of finding an electron.  Distinguished by n (principle quantum number or energy level) and a number (n = 1, 2, 3…)

12.  In each energy level there are orbitals (shapes) called sublevels.  Each energy sublevel corresponds to a different shape, which describes where the electron is likely to be found.

16. Atomic Orbitals video

17.  Each energy level has as many sublevels as the level number (ex: level 1 has 1 sublevel, level 2 has 2 sublevel (shapes).

18.  To find the maximum number of electrons in an energy level use 2n 2.

19. 1s orbital = 1 total orbital 2s orbital 2p orbitals 3s orbital 3p orbitals 3d orbitals 4 total orbitals 9 total orbitals

20. Electron Arrangement in Atoms

21. 5.2.1 - I can describe how to write the electron configuration for an atom. 5.2.2 - I can explain why the actual electron configurations for some elements differ from those predicted by the aufbau principle.

22.  Electron configuration – way in which electrons are arranged in atoms.  Three rules – the aufbau principle, the Pauli exclusion principle, and Hund’s rule – tell you how to find the electron configuration.

23.  Aufbau principle – states that electrons occupy the orbitals of lowest energy first.  Orbitals on any sublevel are always the same energy.

25.  Pauli exclusion principle – an atomic orbital may describe at most two electrons.  When electrons pair they must have opposite “spins” so they don’t repel as much.

26.  Hund’s rule – electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible.  Basically singles in a sublevel until they have to double up.

27.  Shorthand for these are: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 …

29.  Some actual electron configurations differ from those assigned using the aufbau principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations.