Line Spectra Radiation composed of only one wavelength is called monochromatic.

1. Line Spectra and the Bohr Model • Line Spectra • Radiation composed of only one wavelength is called monochromatic. • Radiation that spans a whole array of different wavelengths is called continuous. • White light can be separated into a continuous spectrum of colors. • Note that there are no dark spots on the continuous spectrum that would correspond to different lines.

2. Bohr Model • Colors from excited gases arise because electrons move between energy states in the atom.

3. Bohr Model

4. Types of Spectra • Emission Spectrum – a set of colored lines produced by “downward” transitions between energy levels. • produced when electrons are excited (by electricity or flame) and then return to lower energy levels. • Absorption Spectrum – a continuous spectrum with “dark lines” missing. It is produced by “upward” transitions between energy levels. • produced when white light (or IR, UV, other) is shown through a sample. Specific colors are absorbed. • The two spectra are “complimentary.” The lines and colors involved are exactly the same.

5. Electron Configurations • It is helpful to know how the electrons are arranged within an atom. • This explains and predicts much of the chemistry involved, such as reactions that will occur and ions that will be formed. • Energy levels – • Horizontal rows correspond to energy levels • 1-7

6. Energy Sublevels (s,p,d,f) • Within each energy level are sublevels • Energy level 1 has s only. • Energy level 2 has s and p. • Energy level 3 has s, p, and d. • Energy levels 4 and above have s, p, d, and f. • Orbitals – contain up to 2 electrons each. • s sublevels have 1 orbital • p sublevels have 3 orbitals • d sublevels have 5 orbitals • f sublevels have 7 orbitals

13. Writing Electron Configurations Examples – write electron congurations: • O • Al • Cl • Ni • Pu

14. Orbital Filling Diagrams • This is a visual representation of an electron configuration: • Each orbital is represented by a box () • Each electron is represented by an arrow ( or ) • All orbitals in the same sublevel are drawn together (ex. 3p )

15. Hund’s Rule • An electron configuration tells us in which orbitals the electrons for an element are located. • Three rules: • electrons fill orbitals starting with lowest n and moving upwards; • no two electrons can fill one orbital with the same spin (Pauli Exclusion Principle); • for degenerate orbitals, electrons fill each orbital singly before any orbital gets a second electron (Hund’s rule).

16. Examples Draw orbital filling diagrams: • O • Al • Cl • Ni • Pu

17. Valence Electrons – Electrons in the highest energy level of an atom. Examples: Write the electron configurations and determine the number of valence electrons: • K • S • C • Sb • Pb

18. Orbitals and Quantum Numbers • If we solve the Schrödinger equation, we get wave functions and energies for the wave functions. • We call wave functions orbitals. • Schrödinger’s equation requires 4 quantum numbers: • Principal Quantum Number, n. This is the same as Bohr’s n. As n becomes larger, the atom becomes larger and the electron is further from the nucleus.

19. Azimuthal Quantum Number, l. This quantum number depends on the value of n. The values of l begin at 0 and increase to (n - 1). We usually use letters for l (s, p, d and f for l = 0, 1, 2, and 3). Usually we refer to the s, p, d and f-orbitals. • Magnetic Quantum Number, ml. This quantum number depends on l. The magnetic quantum number has integral values between -l and +l. Magnetic quantum numbers give the 3D orientation of each orbital. • Spin Quantum Number, ms. This quantum number represents the “spin” or rotation of the electron. The possible values are +½ or -½.

20. Examples • Determine the quantum numbers of each electron: a.  4d •  2p • Describe (with drawings) the location of each electron: a. (4,0,0,-1/2) b. (6,2,-1,1/2)