Chapter 6 Electronic Structure of Atoms

Transcription

1 Chapter 6. Electronic Structure of Atoms NOTE: Review your notes from Honors or regular Chemistry for the sequence of atomic models and the evidence that allowed scientists to change the model. If you no longer have your notes, they are available on Dr. Hart s Chemistry weebly: tamchemistryhart.weebly.com See your current textbook for details of the material outlined below. Note that a number of topics in this chapter are no longer required for the AP Chemistry exam and therefore are not addressed in this outline. You are welcome to read about them if you are particularly interested in the related physics. 6.1 The Wave Nature of Light Electromagnetic radiation is characterized by its wave nature. All waves have a characteristic wavelength, λ (lambda), and amplitude, A. The frequency, ν, are Hertz (1 Hz = 1 s 1 ). The speed of a wave is given by its frequency multiplied by its wavelength. For light, speed, c = λν; ΕΜ moves through a vacuum with a speed of approximately 3.00 x 10 8 m/s. The electromagnetic spectrum is a display of the various types of electromagnetic radiation arranged in order of increasing wavelength. E.g. visible radiation has wavelengths between 400 nm (violet) and 750 nm (red)

2 Two electromagnetic waves are represented below. Sample Exercise 6.1 (p. 215) a) Which wave has the higher frequency? b) If one wave represents visible light and the other represents infrared radiation, which wave is which? Practice Exercise 1 (6.1) A source of electromagnetic radiation produces infrared light. Which of the following could be the wavelength of the light? a) 3.0 nm b) 4.7 cm c) 66.8 m d) 34.5 µm e) 16.5 Å Practice Exercise 2 (6.1) If one of the waves in Sample Exercise 6.1 represents blue light and the other red light, which would be which? Sample Exercise 6.2 (p. 216) The yellow light given off by a sodium vapor lamp used for public lighting has a wavelength of 589 nm. What is the frequency of this radiation? (5.09 x s -1 ) - 2 -

3 Practice Exercise 1 (6.2) Consider the following three statements: (i) For any electromagnetic radiation, the product of the wavelength and frequency is a constant. (ii) If a source of light has a wavelength of 3.0 Å, its frequency is 1.0 x Hz. (iii) The speed of ultraviolet light is greater than the speed of microwave radiation. Which of these three statements is or are true? a) Only one statement is true. b) Statements (i) and (ii) are true. c) Statements (i) and (iii) are true. d) Statements (ii) and (iii) are true. e) All three statements are true. Practice Exercise 2 (6.2) a) A laser used in orthopedic spine surgery produces radiation with a wavelength of 2.10 µm. Calculate the frequency of this radiation. (1.43 x s -1 ) c) An FM radio station broadcasts electromagnetic radiation at a frequency of MHz (megahertz; 1 MHz = 10 6 s -1 ). Calculate the wavelength of this radiation. (2.901 m) 6.2 Quantized Energy and Photons Read about black body radiation and the photoelectric effect in Section 6.2 in either your textbook or the earlier edition (blue), then complete the following problems

4 Sample Exercise 6.3 (p. 218) Calculate the energy of one photon of yellow light whose wavelength is 589 nm. (3.37 x J) If one photon of radiant energy supplies 3.37 x J, then how many photons will one mole of these photons supply? (2.03 x 10 5 J/mol) Note: this amount is the magnitude at which chemical reactions take place. Radiation break bonds, i.e. photochemical reactions. Practice Exercise 1 (6.3) Which of the following expressions correctly gives the energy of a mole of photons with wavelength λ? a) E = h b) E = N A λ c) E = hc d) E = N A h e) E = N A hc λ h λ λ λ Practice Exercise 2 (6.3) a) A laser emits light with a frequency of 4.69 x s -1. What is the energy of one photon of the radiation from this laser? (3.11 x J) b) If the laser emits a burst or pulse of energy containing 5.0 x photons of this radiation, what is the total energy of that pulse? (0.16 J) c) If the laser emits 1.3 x 10-2 J of energy during a pulse, how many photons are emitted during the pulse? (4.2 x photons) - 4 -

5 6.3 Line Spectra and the Bohr Model Line Spectra Sample Exercise 6.4 (p. 223) Using the figure above, predict which of the following electronic transitions produces the longest wavelength spectral line: n = 2 to n = 1, n = 3 to n = 2, or n = 4 to n = 3. Practice Exercise 1 (6.4) In the top part of Figure 6.11, the for lines in the H atom spectrum are due to transitions from a level for which n i > 2 to the n f = 2 level. What is the value of n i for the blue-green line in the spectrum? a) 3 b) 4 c) 5 d) 6 e) 7 Practice Exercise 2 (6.4) Indicate whether each of the following electronic transitions emits energy or requires the absorption of energy: b) n = 3 to n = 1; b) n = 2 to n =

6 The Uncertainty Principle Heisenberg s uncertainty principle: We cannot determine the exact position, direction of motion, and speed of subatomic particles simultaneously. For electrons: We cannot determine their momentum and position simultaneously. 6.5 Quantum Mechanics and Atomic Orbitals Schrödinger proposed an equation containing both wave and particle terms: Solving the equation leads to wave functions, ψ. The wave function gives the shape of the electron s orbital. The square of the wave function, ψ 2, gives the probability of finding the electron. That is, ψ 2 gives the electron density for the atom. ψ 2 is called the probability density. Electron density is another way of expressing probability. A region of high electron density is one where there is a high probability of finding an electron. Orbitals and Quantum Numbers If we solve the Schrödinger equation we get wave functions and energies for the wave functions. We call ψ orbitals. Schrödinger s equation requires three 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. Angular momentum quantum number, l. This quantum number depends on the value of n. Usually we refer to the s, p, d, and f orbitals. This quantum number defines the shape of the orbital. Magnetic quantum number, m l. This quantum number depends on l. Magnetic quantum numbers give the three-dimensional orientation of each orbital, e.g. p x, p y and p z. A collection of orbitals with the same value of n is called an electron shell. There are n 2 orbitals in a shell described by a the n value. For example, for n = 3, there are 3 2 = 9 orbitals. A set of orbitals with the same n and l is called a subshell. Each subshell is designated by a number and a letter. For example, 3p orbitals have n = 3 and l = 1. There are n types of subshells in a shell described by a the n value. For example, for n = 3, there are 3 subshells: 3s, 3p and 3d. Orbitals can be ranked in terms of energy to yield an Aufbau diagram. Note that this Aufbau diagram is for a single electron system. As n increases note that the spacing between energy levels becomes smaller. NOTE: Assignment of quantum numbers to individual electrons is no longer required for the AP Chemistry exam

7 6.6 Representation of Orbitals; 6.7 Many-Electron Atoms; 6.8 Electron Configurations; 6.9 Electronic Configurations and the Periodic Table: Read these sections in your textbook, then complete the following problems: Sample Exercise 6.7 (p. 239) Draw the orbital diagram representation for the electron configuration of oxygen, atomic number 8. How many unpaired electrons does an oxygen atom possess? Practice Exercise 1 (6.7) How many of the electrons in the second row of the periodic table (Li through Ne) will have at least one unpaired electron in their electron configurations? a) 3 b) 4 c) 5 d) 6 e) 7 Practice Exercise 2 (6.7) a) Write the electron configuration of phosphorus, element 15. b) How many unpaired electrons does a phosphorus atom possess? Note: Condensed Electron Configurations = Noble Gas configurations Sample Exercise 6.8 (p. 243) What is the characteristic valence shell electron configuration of the group 7A elements, the halogens? Practice Exercise 1 (6.8) A certain element has an ns 2 np 6 outer-electron configuration in its outermost occupied shell. Which of the following elements could it be? a) Be b) Si c) I d) Kr e) Rb Practice Exercise 2 (6.8) What family of elements is characterized by having an ns 2 np 2 outer-electron configuration? - 7 -

8 Sample Exercise 6.9 (p. 243) a) Based on its position in the periodic table, write the complete electron configuration for bismuth, element number 83. b) How many unpaired electrons does a bismuth atom have? Practice Exercise 1 (6.9) A certain atom has an [noble gas]5s24d10504 electron configuration. Which element is it? a) Cd b) Te c) Sm d) Hg e) more information is needed Practice Exercise 2 (6.9) Use the periodic table to write the condensed electron configurations for the following atoms: a) Co (element 27) b) In (element 49) Integrative Exercise 6 (p. 243) Boron, atomic number 5, occurs naturally as two isotopes, 10 B and 11 B, with natural abundances of 19.9% and 80.1% respectively. a) In what ways do the two isotopes differ from each other? Does the electronic configurations of 10 B differ from that of 11 B? b) Draw the orbital diagram for an atom of 11 B. Which electrons are the valence electrons (the ones involved in chemical reactions)? c) Indicate three major ways in which the 1s electrons in boron different from its 2s electrons. d) Elemental boron reacts with fluorine to form BF 3, a gas. Write a balanced chemical equation for the reaction of solid boron with fluorine gas. f) Will the mass percentage of F be the same in 10 BF 3 and 11 BF 3? If not, why is that the case? - 8 -