First Semester Final Exam Study Guide

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1 First Semester Final Exam Study Guide Name: Chemistry multiple choice questions Chapter 4 1. Explain the law of conservation of mass. 2. What mass of product is produced in the following reaction? 2 grams H2 + 3 grams Cl2 3. Complete the chart. Subatomic Particle Charge Size Location in Atom 4. Which subatomic particle makes up most of an atom s mass? 5. Most of an atoms volume is really what? 6. Define isotope and give an example. 7. Why is the atomic mass of a given element usually NOT listed on the periodic table as a whole number? 8. What particle does the atomic number listed on the periodic table represent, and where does this particle reside? Page 1

9. Write the equation that relates the mass # to the number of protons and neutrons in the nucleus of an atom. 10. Calculate the number of protons, neutrons and electrons in a atom of Potassium, K?

2 9. Write the equation that relates the mass # to the number of protons and neutrons in the nucleus of an atom. 10. Calculate the number of protons, neutrons and electrons in a atom of Potassium, K? 11. What is the atomic mass and the atomic number of lead, Pb? 12. Which atom has a great atomic mass: Silver or Arsenic? 13. Which element on the periodic table has the atomic number of 17? 14. Which of the following are isotopes of each other? (CIRCLE ALL THAT APPLY) Ca-40 Cr-50 Cr-53 Cu-63 Fe What is the atomic mass of an atom that has 30 neutrons and 25 protons? Chapter Fill in the orbital diagram below for Tungsten and then write the full electron configuration below. W: Page 2

3 17. Which of the following is an incorrect orbital diagram? Why? 18. Write the unabbreviated electron configuration for the following atoms: a) Helium b) Oxygen c) Aluminum d) Neon 19. Write the abbreviated electron configuration for the following atoms: a) Silver b) Copper c) Tungsten d) Titanium 20. Define ground state. 21. Define excited state. Page 3

4 22. How do atoms jump from ground state to excited state and back? 23. Which atomic particle changes energy states? 24. Label the number total number of electrons at each energy level for each elements below: MAGNESIUM ARGON ZINC -which of the above is also an alkaline earth metal? -which of the above is unreactive? Chapter Define period. 26. Define group. What is another name for group? 27. Define electronegativity. How is it used to determine bonds? 28. Define atomic radius. 29. Define ionization energy. Page 4

5 30. Label the following on the attached periodic table (at the end of SG): 1) Atomic radius trend 2) Electronegativity trend 3) Ionization energy trend 4) Alkali metals 5) Alkaline earth metals 6) Halogens 7) Noble gases 8) Actinides 9) Lanthanides 10) Metals 11) Nonmetals 12) Metalloids Chapter 7 & How is an ion different from an atom? 32. Define cation and give an example. 33. Define anion and give an example. 34. Compare and contrast ionic and covalent bonds. 35. How many electrons are shared in a covalent bond? 36. Define octet rule. Which family on the period table has a complete octet and thus does not react to form compounds? 37. Define valence electron. 38. How many valence electrons do the following atoms have: a) Magnesium b) Bromine c) Calcium d) Phosphorous Page 5

6 39. The electron configuration of oxygen is 1s 2 2s 2 2p 4. How many more electrons does oxygen need to satisfy the octet rule? Chapter Name the following ionic compounds. a) Al(SO4)2 b) PbF4 c) KBr2 d) Ca(OH)2 e) Mg(NO2)2 41. Write the formula for the following ionic compounds. a) Magnesium chlorate b) Potassium cyanide c) Tin (IV) chloride d) Zinc nitrate e) Calcium chlorite 42. Name the following covalent (in other words molecular) compounds. a) As3P5 b) IF7 c) NO2 d) CO e) SiCl4 Page 6

7 43. Write the formula for the following covalent (molecular) compounds. a) dinitrogen pentoxide b) carbon monoxide c) trisulfur difluoride d) triarsenic pentanitride e) dihydrogen monoxide Chapter Circle the subscripts in the following chemical formula. K 2 SO What information does a subscript reveal and how would changing a subscript in a chemical formula change the chemical it represents? 46. Define molar mass. 47. Calculate the molar mass of the following: a) Au = b) H2O = c) Mg(MnO4)2 = d) CH3CH2COOH = Page 7

8 48. Define molecular formula. 49. Define empirical formula. 50. What is the empirical formula for C6H12? 51. What is the empirical formula for C11H22O11? 52. Calculate the percent composition of each element in the following compounds: a) potassium cyanide, KCN b) butane, C4H10 c) sulfuric acid, H2SO4 53. Find the empirical formula of a compound that is % iron and % sulfur. 54. Calculate the empirical formula of a compound containing 1.0 g K, 0.70 g Cr, and 0.82 g of O. Page 8

9 55. A white powder is analyzed and found to have an empirical formula of P2O5. The compound has a molar mass of g/mol. What is the compound s molecular formula? 56. A compound with the empirical formula CH4O was found to have a molar mass of approximately 192 g/mol. What is the molecular formula of the compound? 57. How many atoms are there in a mole of gold? 58. Convert 3.5 moles of nickel into atoms. 59. How many moles are in 1.20 x atoms of Zinc? 60. What is the mass of 5 moles of C6H12O6, glucose? 61. How many grams are in 3 moles of HCl, hydrochloric acid? 62. How many moles are in 1 gram sample of gold? Page 9

10 63. A scientist has a sample of magnesium that has a mass of 5 grams. How many atoms of magnesium does the scientist have? Chapter 11 & Define product. 65. Define reactant. 66. Label the products and reactants in the following chemical reactions: H2 (g) + O2 (g) H2O (l) Copper sulfate (aq) + iron (s) iron sulfate (aq) + copper (s) H2O + CO2 + light C6H12O6 + O2 67. What is changed when balancing a chemical equation? 68. Identify the type of reaction and balance: a) Al + NaOH Na3AlO3 + H2 b) C12H22O11 + O2 CO2 + H2O c) Ca + H2SO4 CaSO4 + H2 d) CaCl2 + NaCO3 NaCl + CaCO3 Page 10

11 69. Predict the products and balance these chemical reactions: a) Cu(NO3)2 + K b) SiF4 + NaOH c) MnO2 + K d) K + Cl2 e) C2H8 + O2 70. Define limiting reactant. 71. Define excess reactant. 72. Mg(OH)2 + 2 HCl MgCl2 + 2 H2O If the reaction begins with 25.44g of Mg(OH)2, what mass of H2O is produced from the reaction? Page 11

12 73. 6 Li + Ca3(PO4)2 2 Li3PO4 + 3 Ca If the reaction above produced 84.3 grams of calcium, what mass of lithium was used to start the reaction? 74. Ba3(PO4)2 + 6 KCl 3 BaCl2 + 2 K3PO4 The above reaction begins with 3.71 g of Ba3(PO4)2 and 3.71 g of KCl. What is the limiting reactant? What is the excess reactant? N2H4 + N2O4 3 N2 + 4 H2O The above reaction begins with g of N2H4 and g of N2O4, what is the limiting reactant? If the reaction actually produces 190.4g of N2, what is the percent yield? Na + Cl2 2 NaCl If 17 grams of Na and 17 grams of Cl2 reaction, what mass of NaCl will be made? Page 12

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First Semester Final Exam Study Guide

First Semester Final Exam Study Guide Name: Chemistry 1 2 3 4 5 6 7 60 multiple choice questions Chapter 1 1. Define matter and list five examples. Chapter 2 2. Define pure substance. 3. Define element.