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Chemistry 141 Test 2 Study Guide

by: Sara Auger

Chemistry 141 Test 2 Study Guide Chemistry 141

Marketplace > Emory University > Chemistry > Chemistry 141 > Chemistry 141 Test 2 Study Guide
Sara Auger
Emory University
GPA 3.8

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About this Document

This study guide covers what will be on the Chemistry 141 Test Two.
Chemistry 141
Dr. Egap
Study Guide
General Chemistry, naming, Molecules and Compounds, Naming Covalent Compounds
50 ?




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This 8 page Study Guide was uploaded by Sara Auger on Sunday October 16, 2016. The Study Guide belongs to Chemistry 141 at Emory University taught by Dr. Egap in Fall 2016. Since its upload, it has received 28 views. For similar materials see Chemistry 141 in Chemistry at Emory University.

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Date Created: 10/16/16
Egap Test Two Study Guide Chemistry 141, 2016 What’s going to be on this test:  Last part of Chapter 3 (Uncertainty Principle and beyond)  Chapter 4  Chapter 5 (up until whatever is covered in class on Monday.)  Material covered the day of the test (probably Chapter 6) will NOT be on the test. Answers and explanations will be listed at the end of the study guide. Any questions, if my email is accessible on the site or Blackboard, feel free to contact me! Disclaimer: some questions are based on those in Nivaldo’s Structure and Properties of Matter. For the best study experience, the odd problems in the book have answers in the back and can give you a wider variety of problems. Chapter 3 1. Explain, in your own words Heisenberg’s uncertainty principle. 2. T/F: The more accurately you know position of an electron, the more accurately you know velocity of an electron. 3. How do Newton’s deterministic Laws apply to electrons? 4. We can only describe electrons ___________. 5. How do we describe the position of an electron? 6. Is spatial distribution of electrons important? Why or why not? 7. Which of the following is not true of quantum numbers? a. N is the principle quantum number, which determines the overall size and energy of the orbital b. L is the angular momentum quantum number, can be greater than or equal to n-1 c. The magnetic quantum number can be equal to -L to +L d. The magnetic spin number is quantized. 8. Which set of three quantum numbers does not specify an orbital in an atom? a. n = 2; l = 1; ml = 1 b. n = 2; l = 2; ml = -1 c. n = 1; l = 0; ml = 0 d. n = 3; l = 2; ml = 2 9. Calculate the wavelength of a light emitted when an electron in an atom makes a transition from an orbital with n = 4 to an orbital with n = 3. 10. Which electron transition produces light of the highest frequency in the hydrogen atom? a. 5p  1s b. 3p  1s c. 4p  2s d. 3p  2s 11. Which of the following is the P orbital shape? A. B. C. D. 12. What is a trajectory? What information do you need to determine it? Chapter 4 1. What explains the observations organized in the Periodic Table? 2. Describe the periodic trends for density, atomic radius, and mass. 3. What are main-group elements and transition elements? What is a big difference between the two groups? 4. T/F: each row or period shares similar properties. 5. What is an electron configuration? 6. What is the lowest energy state in an electron? 7. What kind of energy states do electrons occupy? 8. Explain the Pauli Exclusion Principle. 9. Jesse thinks that because the charge of the nucleus of K is +19, the valence electron in a K atom feels an attractive force of +19. Why is this incorrect? 10. How does penetration effect the charge an outer electron feels? As in, if it penetrates spatially another orbital. 11. According to Coulomb’s Law, as a proton and an electron get closer to each other, what happens to Potential Energy and stability? 12. Which electron in S is most shielded from nuclear charge? a. An electron in the 2s orbital. b. An electron in a 3p orbital. c. An electron in a 4p orbital. d. None of the above. (All of these electrons are equally shielded from nuclear charge.) 13. What is the electron configuration of Sulfur? 14. What set of quantum numbers can describe an electron in a 3s orbital? 15. Pattern of filling orbitals from the lowest energy up is known as the __________. 16. Explain Hund’s Rule. 17. Which atom will have the biggest atomic radius? a. Rb b. Tc c. Xe d. Te 18. Which electrons efficiently shield which electrons? 19. Effective nuclear charge (increases/decreases) as you move to the right across a row in the Periodic Table. 20. What is the electron configuration (use noble gas configuration) of K+? 21. Which is diamagnetic? a. B b. Cr1+ c. Cr d. Zn 22. Order the following in order of increasing atomic radius: Cl-, Ar, K+ 23. Order the following in order of increasing ionization energy. Na, Mg, Cs 24. Describe the periodic trend for metallic character 25. For which element is gaining an electron most exothermic? Rb, Sr, Ag, I Chapter Five 1. T/F: H2 (hydrogen gas) is a natural part of air on earth. 2. What is the difference between a mixture and a compound? 3. Why do chemical bonds form? 4. What is the main difference between ionic bonds and covalent bonds? 5. Metals tend to _____ electrons. Nonmetals tend to ___ electrons. 6. When metals bind to nonmetals, they ______ electrons. When nonmetals bind to each other, they _____ electrons. 7. What is the difference between molecular and empirical formulas? 8. T/F: The Lewis Model and its implications aren’t very accurate at predicting properties and interactions of elements. 9. Describe the Octet Rule and its exceptions. 10. What shape do ionic bonds usually take? 11. Write the name of the following compound: NaI 12. Write the formula for the following compound: phosphorus pentoxide 13. Double bonds are ______ and ______than single bonds. 14.How do you find formula mass for a formula unit of a compound? 15. What is the empirical formula of C10H8? 16. Which of the following is an ionic compound? a. H2O b. KCl c. CCl d. HBr 17. What is the most likely formula for a compound of Calcium and Chlorine? 18. What is the formula for strontium iodide? 19. What are organic compounds? 20. What are inorganic compounds? 21. What is the key element in organic compounds? 22. General things I Would Expect to Need to Know Formulas (Heisenberg Uncertainty Equation) Facts Newton’s Laws are deterministic. We can’t know trajectory of electrons. Know how to read a probability distribution map. Know the properties etc. of the different quantum numbers. Know all the periodic table trends. Vocabulary Indeterminacy: you can do the exact thing over and over again and get different results. Deterministic: present determines future Probability Distribution Map: shows probability of electron being in certain area. Know Orbital and Sublevel (page 83). I’m not listing them here because I figure most people do know them. Periodic Property: one that is generally predictable based on element’s location. Periodic Law: when the elements are arranged in order of increasing mass, certain sets of properties occur… periodically. Know the details and implications of Coulomb’s Law. Degenerate: orbitals having the same energy Answers Chapter 3 1. Position and velocity of an electron can’t both be measured at the same time, because they come from particle and wave natures (respectively). Both natures can’t be observed at once. 2. False; more accurately you know one, less accurately you know the other. 3. They don’t. 4. Statistically. 5. We use orbitals. 6. Yes, because it is important in bonding due to electron sharing. 7. B; L can be less than or equal to n-1 8. B; L can’t be the same as n 9. Wavelength = 1.88E-6 10. A 11. B 12. It’s a path, you need to know particle’s velocity, position, and the forces acting on it. Chapter 4 1. Quantum Mechanics 2. They all increase as you go down a column and to the right. 3. Main group elements tend to have more predictable properties based on their location. Transition elements do not. 4. False; each group or family does. 5. Shows the particular orbitals that electrons occupy for that atom. 6. Ground State 7. The lowest energy level available. 8. No two electrons in an atom can have the same four quantum numbers. 9. Shielding; one electron can shield another electron form the full force of the nucleus, especially the closer the shielding electron is to the nucleus. 10. It is closer to the nucleus and feels less shielding from a more inner electron. 11. Potential Energy decreases and stability increases. 12. C 13. 14. N = 3, l= 0, ml = 0, ms = +/- ½. 15. Aufbau Principle 16. When filling degenerate orbitals, electrons first fill in singly and then in pairs. 17. Xe 18. Core electrons efficiently shield outermost electrons. 19. Increases 20. [Ar] 21. Zn 22. K=, Ar, Cl- 23. Cs, Na, Mg 24. Metallic character decreases as we move right across a period, and increases as we move down a group. 25. I Chapter Five 1. False 2. In a compound, elements combine in fixed, definite proportions; in a mixture, elements can mix in any proportions whatsoever. 3. They lower the potential energy of the charged particles that compose atoms. 4. Ionic bond forms between a metal and a nonmetal, covalent forms between two nonmetals. (Covalent Co.  Company  Companies are exclusive  Only form with nonmetals.) 5. Lose, gain. 6. Transfer, share. 7. Empirical formulas show the relative number of atoms of each element in a compound. Molecular formulas show the actual number of atoms of each element in a molecule of a compound. 8. False; they’re really accurate. 9. Usually stable atoms or molecules or lattices have 8 valence electrons, and so atoms bond to form those. Exception: Helium forms a duet, not an octet. 10. Lattices 11. Sodium iodite 12. P2O5 13. Shorter and stronger. 14. It’s the sum of all the atomic weights in the formula. 15. C5H4 16. KCl 17. CaCl2 18. SrI2 19. Composed of carbon, hydrogen and a few other elements. Come from living things. 20. Compounds that come from the earth (ex: salt) 21. Carbon


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