Name Date Due Atoimic Structure and the Periodic Table: Unit Objective Study Guide Part 2 Directions: Write your answers to the following questions in the space provided. For problem solving, all of the work leading up to the final answer must be shown in order to receive credit. 1. Describe the historical development of the periodic table (Döbereiner, Newland, Mendeleev, Moseley, and Seaborg). Describe the contributions made by each of the following scientists to the development of the periodic table. a. Döbereiner (What is a triad?) b. Newland (What was Newland s law of octaves?) c. Mendeleev (Why did he leave gaps in his periodic table? What were some limitations to his arrangement of the elements?) d. Moseley e. Seaborg 2. Describe the arrangement of the periodic table. The elements in the periodic table are arranged in order of increasing atomic. The vertical columns are called or. The horizontal rows are called. The elements within the same have similar chemical properties. There are periods and groups. The metals are located on the side of the periodic table. The nonmetals, with the exception of hydrogen, are located on the side of the periodic table. The metalloids are located the metals and nonmetals.
3. Describe the general properties of the alkali metals, alkaline earth metals, halogens, noble gases, inner transition metals (lanthanide and actinide series), and the transition metals. List at least 3 distinctive properties for each of the following. a. alkali metals b. alkaline earth metals c. halogens d. noble gases e. inner transition metals (lanthanide and actinide series) f. transition metals 4. Classify elements as metals, nonmetals or metalloids; solids, liquids or gases; alkali metals, alkaline earth metals, transition metals, halogens, noble gases or inner transition metals (lanthanides or actinides). Identify the classes of elements to which each of the following elements belongs. Each element is described by more than one term. Choose from the following classes: metal, nonmetal, metalloid; alkali metal, alkaline earth metal, transition metal, inner transition metal (lanthanide or actinide), halogen, noble gas; solid, liquid, gas (at room temperature). a. argon: b. germanium: c. mercury: d. phosphorus: e. magnesium:
5. Distinguish between principal energy level, energy sublevel, and atomic orbital. 1s 2 Principal energy level energy sublevel atomic orbital - 6. Use the Bohr model to draw representations for atoms. 7. Apply the aufbau principle, the Pauli exclusion principle, and Hund s rule in writing the electron configurations and orbital diagrams of elements. 8. Write the electron configuration and noble-gas notation of elements using the periodic table as a guide. Complete the following table. Nitrogen-14 Calcium-40 Xenon-131 Bohr Model Orbital Diagram Electron Configuration Bohr Notation Noble-Gas Notation
9. Distinguish between the ground state and the excited state of an atom. Ground state Excited state 10. Recognize the demarcation of the periodic table into an s block, p block, d block, and f block. Color the s block blue. Color the p block green. Color the d block red. Color the f block purple. For each of the following electron configurations of neutral atoms, determine if the configuration as written is the ground state, an excited state, or if it is an impossible configuration. a. 1s 2 2s 2 2p 4 b. 1s 2 2s 3 2p 6 c. 1s 2 2s 2 2p 5 3s 1 11. Use electron configurations to classify elements as noble gases, alkali metals, alkaline earth metals, halogens, transition metals or inner transition metals. Explain how an element s outer electron configuration is related to its position in the periodic table. Identify the ending electron configuration for each of the following classes of elements. The first one has been done for you. a. noble gases s 2 (helium), p 6 (other noble gases) b. alkali metals - c. alkaline earth metals - d. halogens - e. transition metals - f. inner transition metals - 12. Define valence electrons and determine how many are present in atom of each maingroup (representative) element. What are valence electrons? Indicate two different ways in which the valence electrons for a main-group element can be determined. Indicate the number of valence electrons for each of the following elements. a. potassium b. magnesium c. boron d. carbon e. phosphorus f. oxygen g. fluorine h. neon
13. Describe the octet rule and use it to predict the charge on a stable ion. State the octet rule 14. Define isoelectronic and give examples. What does the term isoelectronic mean? Complete the following table. Element Number of Valence Electrons Charge on Stable Ion Sodium 1 +1 Calcium Aluminum Nitrogen Sulfur Iodine Is it possible for two atoms to be isoelectronic? List four ions that are isoelectronic with argon. 15. Define periodicity. Periodicity - 16. Define atomic and ionic radii, ionization energy, and electronegativity. a. atomic radii - b. ionic radii c. ionization energy Which group of elements has the highest first ionization energy? Which group of elements has the lowest first ionization energy? Why is the third ionization energy for calcium much higher than the second ionization energy? d. electronegativity Which element has the highest electronegativity value?
17. Describe what is meant by nuclear charge (force) and the role it plays on periodic trends. What is nuclear charge (force)? Why does nuclear charge increase as you go from left to right across a period? Why are negative ions (anions) always larger than the neutral atom from which they came? (i.e. Why is the chloride ion larger than the chlorine atom?) Why are positive ions (cations) always smaller than the neutral atom from which they came? (i.e. Why is the magnesium ion smaller than the magnesium atom?) 18. Interpret and explain period trends in atomic radii, ionic radii, ionization energies and electronegativities. Explain how nuclear charge affects each of the following period trends: a. atomic radii - Arrange the following elements in order of increasing atomic radii: calcium, bromine, arsenic, and gallium b. ionic radii - Which ion is larger, the sulfide ion or the chloride ion? Explain your answer. c. ionization energy - Arrange the following elements in order of increasing first ionization energy: rubidium, xenon, strontium, tin, and iodine. d. electronegativity - Arrange the following elements in order of increasing electronegativity: lithium, fluorine, beryllium, boron, nitrogen, and oxygen.
19. Describe the shielding effect and the role it plays on periodic trends. What is the shielding effect? How does the shielding effect change as you go from top to bottom within a group? How does the shielding effect change as you from left to right within a period? 20. Interpret and explain group trends in atomic radii, ionic radii, ionization energies and electronegativities. Explain how the shielding effect affects each of the following group trends: a. atomic radii - Arrange the alkali metals in order of increasing atomic radii. b. ionic radii - c. ionization energy - Arrange the alkaline earth metals in order of increasing first ionization energy. d. electronegativity Arrange the halogens in order of increasing electronegativity.