Quick Review. 1. Kinetic Molecular Theory. 2. Average kinetic energy and average velocity. 3. Graham s Law of Effusion. 4. Real Gas Behavior.

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1 Quick Review 1. Kinetic Molecular Theory. 2. Average kinetic energy and average velocity. 3. Graham s Law of Effusion. 4. Real Gas Behavior.

2 Emission spectra Every element has a unique emission spectrum called a fingerprint.

3 Electromagnetic Radiation Energy travels through space in the form of electromagnetic radiation. All electromagnetic radiation travels in wave motion.

4 Electromagnetic Radiation What is the difference between wavelength and frequency? Wavelength (λ) distance between two consecutive peaks or troughs. Frequency (ν) number of waves per second that pass through a given point in space.

5 Electromagnetic Radiation How are frequency and wavelength related? Example: Are gamma rays (λ = m) a form of high or low frequency radiation?

6 Electromagnetic Radiation Electromagnetic radiation has the properties of waves but transfers energy in a manner similar to particles. Called wave-particle duality of light. Only certain quanta or packets of energy can be transferred by the various λ that make up the EMR spectrum. Packets of energy are called photons.

7 Electromagnetic Radiation How is the energy per photon quantified? What is the relationship between energy and frequency? What is the relationship between energy and wavelength?

8 Electromagnetic Radiation Calculate the wavelength and energy of UV irradiation (λ = 1 x 10-8 m).

9 Hydrogen Emission Spectrum H atoms are excited, they contain excess energy, which they release by emitting light of various wavelengths. Every element has a unique spectrum called a fingerprint. Hydrogen emission spectrum can be explained by the Bohr model.

10 Bohr Model 1. Electrons occupy circular orbits about the nucleus. 2. Only certain orbits are allowed (indicated by n = 1, 2, 3, ) 3. Each orbit has an energy associated with it: E n = -Z 2 R H /n 2 4. Electrons move from one orbit to another by absorbing or emitting photons of light. E = E 2 E 1 = -Z 2 R H (1/n 22 1/n 12 ) R H = x J (Rydberg constant) Z = atomic number Bohr Model ONLY works for one electron systems (H, He +,..)

11 Bohr Model n represents allowed energy states called principle quantum number, has values = 1,2,3,4,5, n =1 is the ground state (lowest energy, electron is closest to the nucleus). n = 2 and greater are called excited states. n = means the electron is completely removed from the atom. Absorption electron moves to a higher energy level, E is positive. Emission electron moves to a lower energy level, E is negative.

12 Think about it How does the Bohr model help explain the emission spectrum of hydrogen? Why don t we see more lines in the hydrogen emission spectrum since there are an infinite number of allowed transitions?

13 Bohr Model The red-orange line in the hydrogen emission spectrum corresponds to the n =3 to n = 2 transition (λ = nm). Are the other transitions of higher or lower energy?

14 Think about it Each atom/ion has its own unique emission spectrum, why?

15 Ionization Energy Ionization Energy the energy required to completely remove an electron from the ground state (n 1 = 1 to n 2 = ). Calculate the ionization energy for hydrogen.

16 Ionization Energy Calculate the ionization energy for Li 2+.

17 Supplemental Question An electron is excited from the n = 1 ground state to the n = 3 excited state in a hydrogen atom. Which of the following statements are true? a) It takes more energy to ionize the electron from n =3 than from the ground state. b) The electron is farther from the nucleus on average in the n = 3 state than in the n =1 state. c) The wavelength of light emitted if the electron drops from n =3 to n = 2 will be shorter than the wavelength of light emitted if the electron falls from n = 3 to n =1. d) The wavelength of light emitted when the electron returns to the ground state from n = 3 will be the same as the wavelength of light absorbed to for from n = 1 to n = 3. e) For n = 3, the electron is in the first excited state.

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