Skill Practice 10 1. Define the terms ground state and excited state. Ground state: the normal energy level that an electron occupies. Excited state: when an electron has absorbed energy to occupy a higher energy level. 2. What is the wavelength of light that has a frequency of 4.22 x 10 15 Hz? 7.11x10-8 m 3. What is the energy of light that has a frequency of 1.30 x 10 14 Hz? 8.62x10-20 J 4. A certain atom has a green spectrum line of about 540 nm. What is the difference in energy between the two energy levels responsible for producing the line? 3.68x10-19 J 5. The wavelength of a certain beam of light was 3.52x10-7 m. a) Find the frequency of this light. 8.52x10 14 Hz b) Calculate how much energy this light has. 5.65x10-19 J 6. What is the frequency and wavelength of light that has energy of 5.09 x 10-19 J? f = 7.68x10 14 Hz λ = 3.91x10-7 m
Skill Practice 11 1. What is wrong with the following notation? There is no such thing as an f sublevel in the third energy level. Also, an f sublevel if it existed would have five orbitals instead of four. 2. How many sublevels would you expect in the 8 th energy level? 8 3. What is the maximum number of electrons that can fit in the 3d sublevel? 10 4. How many electrons can fit in a 2p orbital? 2 5. In the 5 th energy level, there is a fifth sublevel called the g sublevel. Considering the trend in number of orbitals and electrons in the s, p, d, and f sublevels, predict how many orbitals and how many electrons can fit in a g sublevel. Orbitals = 9 Electrons = 18 6. Considering your answer to question 5, how many electrons can fit in the entire 5 th energy level? 50 7. Write the notation for an electron spinning clockwise in a p sublevel in the 4 th energy level.
Skill Practice 13 1. Using arrows, write complete orbital diagrams for a) Scandium 1s 2s 2p 3s 3p 4s 3d b) Molybdenum 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d c) Selenium 1s 2s 2p 3s 3p 4s 3d 4p 2. Write the complete electron configuration (no arrows) for a) Chromium b) Antimony c) Calcium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 3 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3. What is wrong with the following electron orbital diagram? What is the name of the rule that allows you to identify the error? The 2p sublevel should look like 2p because of Hund s Rule. 4. How many unpaired electrons does cobalt have? Three
Skill Practice 15 1. Use the noble gases to write abbreviated electron configurations for a) Germanium b) Barium c) Bromine d) Bismuth [Ar]4s 2 3d 10 4p 2 [Xe]6s 2 [Ar]4s 2 3d 10 4p 5 [Xe]6s 2 4f 14 5d 10 6p 3 e) Manganese f) Gold [Ar]4s 2 3d 5 [Xe]6s 2 4f 14 5d 9 2. What column of the periodic table contains elements whose electron configurations end with d 4? Column 6 (or VIB) 3. What row of the periodic table contains elements with 4d electrons? 5 th row 4. What row of the periodic table contains elements with 3p electrons? 3 rd row 5. In each row of the d block there are only 10 elements. Why is this? Each d sublevel can hold 10 electrons. 6. In each row of the p block there are only 6 elements. Why is this? Each p sublevel can hold 6 electrons.
Skill Practice 16 1. What force of attraction does the second energy level of a phosphorus atom feel from the nucleus? Draw a Bohr diagram and use it to explain your answer. The nucleus of a phosphorus atom has 15 protons (a +15 charge) which gets weakened by 2 electrons in the first level and 8 electrons in the 2 nd level. Thus the outer level feels a +5 charge from the nucleus. 2. Using the concepts of shielding and attraction, explain why sulfur is smaller in radius than silicon. Sulfur s outer level of electrons feels a +6 charge pulling on it from the nucleus. Silicon s outer level feels a weaker charge of +4 and therefore it is not pulled as close to the nucleus as sulfur s outer level. 3. Why can t you tell by looking at the periodic table whether chlorine or lithium is larger? Chlorine is lower in the periodic table which would indicate that it is larger than lithium, but it is also located further to the right which would indicate that it is smaller. 4. Order the following elements from smallest to largest. A) Al, Na, S, Mg B) C, Sn, Pb, Si S, Al, Mg, Na C, Si, Sn, Pb C) K, Se, Ca, Br D) Be, Ca, C, B, Mg Br, Se, Ca, K C, B, Be, Mg, Ca E) Ga, Al, Cl, P F) O, Se, S, Ne Cl, P, Al, Ga Ne, O, S, Se
Skill Practice 17 1. If an atom has a high first ionization energy does this mean that it is relatively easy or relatively hard to remove an electron from the atom? It is relatively hard to remove the first electron. 2. Arrange the following atoms in order from lowest to highest 1 st ionization energy. A) Ca, Se, As, Br B) As, N, P, Bi Ca, As, Se, Br Bi, As, P, N C) Ga, Al, S, Si D) Li, K, O, C Ga, Al, Si, S K, Li, C, O E) Te, O, S, Po F) In, Te, Sn, I Po, Te, S, O In, Sn, Te, I 3. A certain atom in the 2 nd period has an unusually high 3 rd ionization energy. Name this element. Draw a Bohr diagram and use it to illustrate why you were able to identify this atom. Beryllium; below is the Bohr diagram for beryllium. Notice that after its two outermost electrons are removed, the next electron would need to be removed from an inner energy level, which is much more difficult. 4. Compare the trends for size and for ionization energy. As the size of an atom increases, what happens to the ionization energy? Explain why the ionization energy seems to depend on the size. Larger atoms tend to have smaller first ionization energies because the outer electrons are farther from the nucleus. Since the electrons are farther, the force of attraction from the nucleus is less.