1. Chapter 5

2.  Light is electromagnetic radiation (energy that exhibits wavelike behavior).  ER moves at a constant speed (c) of 3.0 x 10 8 m/s (through a vacuum).

5.  Wavelength and frequency are inversely proportional to one another; as wavelength increases, frequency decreases.  Speed of light= C = 3.00 x 10 8 m/s  C= λ ν  Rearrange the eqn. to solve for frequency and wavelength!  Practice Problem  A certain green light has a frequency of 6.26x10 10 Hz. The speed of light is 3.00 x 10 8 m/s. What is the wavelength?  λ = c/v = 3.00x10 8 m/s= 4.79 x10 -3 m 6.26x10 10 1/s

7.  Violet light has a wavelength of 4.10x10 -7 m. The speed of light is 3.00 x 10 8 m/s. What is the frequency of the light?  λ = 4.10x10 -7 m  c = 3.00 x 10 8 m/s  V = ?  C = λ ν  v = c / λ = 3.00x10 8 m/s= 7.32x10 14 1/s or Hz 4.10x10 -7 m

8.  What is the energy of light of frequency 2.13x10 12 Hz? Plank’s constant = 6.626x10 -34 J/Hz.  v = 2.13x10 12 Hz  h = 6.626x10 -34 J/Hz  E = ?  E = hv  E = hv = (6.626x10 -34 J/Hz)(2.13x10 12 Hz)  E = 1.41x10 -21 J

10. This diagram is called a line-emission spectrum which shows the frequencies of light emitted by hydrogen when its electrons return to ground state.

11.  Linked the atom’s electron with photon emission.  Electrons are allowed to circle the nucleus only in fixed paths (orbits).  Atom is in the lowest energy level when the electron is closest to the nucleus.  An electron can move to a higher orbital by gaining energy.  When an electron falls from an excited state, a photon is emitted  Bohr’s model mathematically explained the line- emission spectrum of hydrogen.

15. Quantum Number TypeAbbrev.Meaning Principle Quantum NumbernMain energy level occupied by electrons (1, 2, 3, 4, etc) i.e. which row it’s in Angular Momentum Quantum Number lIndicates the shape of the orbital (s, p, d, f) Magnetic Quantum NumbermIndicates the orientation of an orbital around the nucleus (For p, -1, 0, +1) Spin Quantum NumbersTwo possible values (+ ½ or – ½) which indicate the spin value of an electron in an orbital.

16.  S orbital has one shape.  It can hold 2 electrons  For s, m=0

17.  3 different orientations of the p orbitals.  Each orientation can hold two electrons.  Each corresponds to an m value m = -1, 0, or +1

18.  5 different orientations  Each can hold two electrons  m = -2, -1, 0, +1, or +2

19.  The periodic table is organized based on electron configuration.

21. AUFBAU RULE: Add electrons to the lowest energy levels and sublevels first and then go to the next level until all of the electrons have been accounted for. Diagonal rule: the order of filling

22. Examples 1H1H n = 1s 1 (1 electron, close to the nucleus) 2 He 1s 2 (2 electrons, both in the s-sublevel) 3 Li1s 2 2s 1 10 Ne1s 2 2s 2 2p 6 28 Ni1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 Argon’s configuration 17 Cl1s 2 2s 2 2p 6 3s 2 3p 5 Neon’s configuration [Ar [Ar]4s 2 3d 8 18e - [Ne [Ne]3s 2 3p 5 10e - Short-cut electron configuration: build configuration on the noble gas gas that ends the previous previous row on the periodic table. Begin with the “n’s” (1,2,3 …), always followed by the s clouds first (1s, 2s, 3s etc.)and continue to fill electrons until you are at your given element. Short-Cuts Noble gases Ar Ne Xe Rn Kr [He]2s 2 2p 6 He Same row (n) as element given Row above element given

23. Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results. 1s 2 2s 2 2p 3 nitrogen

24. Each orbital can hold a maximum of 2 electrons. To occupy the same orbital, 2 electrons must spin in opposite directions. When 2 electrons share an orbital, they are called “paired”. 2s 2 2p 6 2s 2 2p 6 [Ne]