Name Period Date Semester 1 Review Chemistry Units & Unit Conversions Ch. 3 (p. 73-94) PART A SI UNITS What type of measurement is indicated by each of the following units? Choices are in the last column. 1. g/ml 2. s 4. g 5. cm 3 7. mg 8. L density length mass 3. km 6. mm 9. g/cm 3 time volume PART B DENSITY 10. A small gold nugget has volume of 0.87 cm 3. What is its mass if the density of gold is 19.3 g/cm 3? 11. What volume is occupied by 35.2 g of carbon tetrachloride if its density is 1.60 g/ml? PART B UNIT CONVERSIONS Perform the following SI prefix conversions: 12. 25 kg = g 13. 9.3 ml = L 14. 0.36 mm = m 15. 24 cm = km Atomic Structure 4,5 (100-153) STATES OF MATTER 1. Draw a diagram to represent 8 particles in each state of matter. (HINT: Recall that plasma is gaslike but with one major difference.) SOLID LIQUID GAS PLASMA PART A SUBATOMIC PARTICLES The table below contains information about several elements. In each case, enough information has been provided for you to fill in the blanks. Assume all atoms are neutral. Ch. 3 Uncertainty in Measurements CHEM
Isotope Name Nuclear Symbol Atomic Number Mass Number # of Protons # of Electrons # of Neutrons 1. calcium-40 2. 12 24 3. 1 2 197 4. 79 Au 5. 26 30 6. 201 80 7. 17 18 PART B AVERAGE ATOMIC MASS 8. Calculate the average atomic mass for neon if its abundance in nature is 90.5% neon-20, 0.3% neon-21, and 9.2% neon-22. 9. Calculate the average atomic mass of silver if 13 out of 25 atoms are silver-107 and 12 out of 25 atoms are silver-109. 10. Distinguish between mass number, relative atomic mass, and average atomic mass. Periodic Trends Ch. 6 (p. 154-185) 1. Put the following elements in order from smallest to largest atomic radius and explain why: C, O, Sn, Sr. 2. Put the following elements in order from lowest to highest first ionization energy and explain why: Al, Ar, Cs, Na. Ionic Compounds Ch. 2 & 13
3. Explain what happens to the radius of an atom as you move from left to right across a period. 4. Define Cation and Anion and describe what types of atoms form cations and anions. Ionic Compounds Ch. 7 (p. 186-201) PART A IONIC NAMES Name the following ionic compounds. Remember to consider polyatomic ions and Roman numerals. 1. ZnCl 2 PART B IONIC FORMULAS Write formulas for the following ionic compounds. Remember to consider polyatomic ions and Roman numerals. 9. potassium iodide 2. Cu 2 CO 3 3. Na 2 C 2 O 4 4. Fe 2 O 3 5. MgCO 3 6. NH 4 NO 3 7. LiF 8. SnBr 4 10. iron(iii) sulfide 11. potassium phosphate 12. magnesium chloride 13. tin(iv) chloride 14. sodium hydroxide 15. mercury(ii) oxide 16. aluminum acetate Ionic Compounds Ch. 2 & 13
Use Lewis electron dot diagrams to show the formation of the following molecular substances. Include partial charges if a polar bond is formed. 16. Br 2 17. PCl 3 PART C MOLECULAR NAMES Name the following compounds. 18. SiO 2 19. S 4 N 2 20. Cl 2 O 21. PF 3 PART D MOLECULAR FORMULAS Write formulas for the following compounds. 22. sulfur trioxide 23. carbon tetrachloride 24. diphosphorous pentoxide 25. sulfur hexafluoride Molar Conversions Ch. 10 (p. 286-319) ***FOR ALL CALCULATIONS, SHOW YOUR WORK AND INCLUDE UNITS & SIG FIGS*** PART A MOLAR MASS 1. Calculate the molar mass for each of the following compounds. Include units! calcium nitrate Formula: lead(ii) iodide Formula: Semester 1 Study Guide 2011-2012 CHEM
PART B MOLAR CONVERSIONS (Show your work and include units!) 2. How many moles of ammonia are in 1.20 10 25 molecules of ammonia? Formula: Answer: 3. You need 2.5 moles of aluminum for an experiment. How many atoms of aluminum is this? Formula: Answer: 4. There are 3.20 10 22 atoms of copper in the outer shell of pennies. How many grams of copper is this? Formula: Answer: Balancing Equations CH. 11 (320-341) Balance the following chemical equations. 1. N 2 + O 2 N 2 O 2. KI + Cl 2 KCl + I 2
3. C 2 H 6 + O 2 CO 2 + H 2 O 4. Mg(NO 3 ) 2 + K 3 PO 4 Mg 3 (PO 4 ) 2 + KNO 3 Stoichiometry Ch. 12 (p.352-382) 1. How many liters of ammonia gas are formed in this reaction after 25L of Nitrogen React? N 2 + H 2 NH 3 2. 50.0 ml of 2.00M H 2 SO 4 react, how many grams of Na 2 SO 4 will be formed? H 2 SO 4 + NaOH Na 2 SO 4 + H 2 O 3. If 6.57 g of iron reacted and 14.63 g of iron(iii) chloride are obtained. Calculate the theoretical yield and percent yield of FeCl 3. Fe + HCl H 2 + FeCl 3
4. How many grams of lithium sulfate will be formed when 25 grams of sulfuric acid react with an excess of lithium hydroxide in the following reaction? H2SO4 + LiOH Li2SO4 + H2O 5. How many grams of sodium acetate are required to react completely with 30 grams of iron (III) nitrate in the following reaction? NaC2H3O2 + Fe(NO3)3 NaNO3 + Fe(C2H3O2)3 6. How many grams of fluorine gas will it take to convert 100 grams of C2H4 to C2H4F2 using the following reaction? C2H4 + F2 C2H4F2 The Nature of Solutions Ch. 15,16 (p.444-503) 5. What is the difference between a strong electrolyte and a weak electrolyte? 6. Define the following Terms; unsaturated, saturated and supersaturated. Use the solubility curve to answer questions 6 8. 7. How many grams of potassium nitrate can dissolve in 100 g of water at 50 C? 8. At 20 C, a solution contains 120 g of NaNO 3 in 100 g of water. Is this solution saturated, unsaturated, or supersaturated? 9. You need to make a solution containing 150 g of potassium chloride in 300 g of water. What temperature is required?
Molarity Solve the following molarity problems. Show all work and include correct units and sig figs! 10. Find the molarity of a solution in which 58 g of NaCl are dissolved in 2.5 L of solution. 11. How many grams of KMnO 4 should be used to prepare 2.00 L of a 0.500M solution? 12. What volume of 0.25M solution can be made from 5.0 g of KCl? 13. Find the molarity of a 450 ml solution containing 13.7 g of ZnSO 4. 14. How many grams of CuCl 2 are required to make 75 ml of a 0.20M solution?