Name: Date: Period: CP CHEMISTRY SPRING FINAL REVIEW SHEET NOTE: Below are concepts that we have covered in class throughout the year. Questions are organized by chapter/concept to help you study. You need to be able to understand the concepts listed AND apply this knowledge to answer the questions that are provided. Keep in mind the final has 100 questions total, approximately 50 questions from the first semester and 50 questions from the second semester. HAPPY STUDYING! Investigation and Experimentation 1. Be able to recognize which step of the scientific method is being described from different examples. Remember the basic steps are the following: a) observation, b) hypothesis, c) experimentation or testing, d) collecting or recording data, e) drawing conclusions. 2. As an example, which step of the scientific method would each of the following be? a) After years of testing and analyzing data a scientist determines that chemicals in a city s drinking water are causing an increase in the intelligence of children in the city. e) drawing conclusions b) A biology student notices many frog tadpoles in a particular pond have two tails. a) observation c) The biology student prepares tables and graphs of all the information she has collected.d) collecting or recording data d) The biology student guesses that waste water from a sewage treatment plant which empties into the pond is causing the two tailed tadpoles. b) hypothesis e) The biology student obtains some freshly laid frog eggs and keeps some in the pond water and others in distilled water while waiting for them to hatch. c) experimentation or testing 3. Know the difference between an observation and a conclusion. As an example, classify each of the following as an observation or conclusion. a) I see smoke coming from the engine so the car must be overheating. conclusion b) The eggs stink. observation c) The eggs are rotten. conclusion d) Vinegar tastes sour so it must be an acid. conclusion 4. Be able to infer data from a table or graph. For example, analyze the data in the tables below and answer the following questions. a) Salt A and salt B were dissolved separately in 100-milliliter beakers of water. The water temperatures were measured and recorded as shown in the table below. Is the dissolving of Salt A and Salt B exothermic or endothermic? SALT A SALT B Initial water temperature 25.1 C 25.1 C Final water temperature 30.2 C 20.0 C - Dissolving of Salt A: Temp. increase so EXOTHERMIC - Dissolving of Salt B: Temp. decrease so ENDOTHERMIC b) Given the reaction: A + B AB The table below shows student data obtained about the rate of reaction (time to react) when the concentration of solution A is kept constant and the concentration of solution B is diluted by adding H 2 O. How does reaction rate change as water is added? Trial Volume of Volume of Volume of Reaction Solution A Solution B H 2 O Added Time 1 10 ml 10 ml 0 ml 10.4 sec 2 10 ml 5 ml 5 ml 4.9 sec Trial 1 took longer so it had a SLOWER rate. Trial 2 had a smaller time so it had a FASTER rate. Trial 2 had the diluted solution so as you add water, the rate INCREASES. CHAPTERS 1 & 2: Metric Conversions 5. Metric Conversions: Give the value of the following in the units indicated. a. 12.3 m = 0.0123 km b. 0.123 L = 123 ml c. 1.23 cg = 12.3 mg d. 123 km = 0.123 Mm CHAPTERS 4 & 25: Structure of the Atom and Nuclear Chemistry 6. Complete the table: Mass (amu) charge Proton 1 +1 Neutron 1 0 Electron 0-1
7. The major portion of an atom s mass consists of which subatomic particles? Neutrons and protons (electrons have very small mass) 8. What is the relative size, charge and mass of the nucleus compared with the rest of the atom? The nucleus has most of an atom s mass but a very small volume (size). The charge of the nucleus is positive. 9. Which isotope is the standard for the atomic mass unit (amu)? 1 atom of carbon-12 has a mass of 12 amu 10. Isotopes of the same element have different masses because they have different number of neutrons. 11. When given the isotope symbol for an element be able to give its number of protons, electrons and neutrons. As an example: Give the number of each of these particles for. 30 protons, 30 electrons 40 neutrons (70 is the mass number = # protons + # neutrons. 30 is the atomic number = # protons.) 12. Complete the table below for the three types of natural radioactivity: symbol Mass (amu) charge Alpha α 4 +2 Beta β 0-1 Gamma γ 0 0 13. Why does natural radioactivity occur? Spontaneous decay (breakdown) of unstable nucleus (nuclei) 14. Know how to balance a nuclear reaction selecting the correct particle. As an example: What particles would you use to balance the following nuclear reactions? a) (a) (b) (c) (d) b) (a) (b) (c) (d) CHAPTER 5 Electrons in Atoms 15. Be able to write the electron configurations and draw the corresponding orbital filling diagrams for the first twenty elements. As an example: Draw the orbital filling diagrams (boxes) and electron configurations (1s 2 2s 2.) for the following elements. (NOTE: Z = Atomic Number) a) O (Z= 8) 1s 2 2s 2 2p 4 b) K (Z= 19) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 16. The rule that atoms follow when forming bonds in order to attain a noble gas electron configuration: OCTET RULE CHAPTER 6 The Periodic Table 17. In the modern Periodic Table, the elements are arranged in order of increasing atomic number (number of protons). 18. Know where each of the following groups are found on the periodic table and which elements belong to the group. a) Alkali metals GROUP 1A b) Alkaline earth metals GROUP 2A c) Halogens GROUP 7A d) Noble gases GROUP 8A 19. Elements that are in the same GROUP (COLUMN) OR PERIOD (ROW) on the periodic table share similar chemical properties. 20. How does electronegativity vary within a group and within a period on the periodic table (i.e., the trend)? Know which elements on the periodic table have the highest electronegativity and the lowest electronegativity. Element with highest electronegativity: Fluorine (remember Group 8A has zero electronegativity). Element with lowest electronegativity: Francium (bottom left corner) 2
Know which element has the highest atomic radius and which element has the smallest atomic radius. Smallest atomic radius: Helium Largest atomic radius: Francium (bottom left corner) 21. Know that elements in the same group (vertical column) on the periodic table have the same number of valence electrons in the outer energy level. Elements in the same period (horizontal row) have the same principle energy level. a. How many valence electrons do elements in Group 3A have? 3 valence electrons b. What is the outermost energy level for these elements: O n = 2, Al n = 3, Sr n = 5 22. If given the group number and the period number, be able to identify the element. As an example: What is the element in group 2A and period 3? Mg What is the element in group 7A and period 2? F CHAPTER 8/9 Ionic, Metallic and Covalent Bonds 23. Know the differences between ionic, covalent, and metallic bonds. a. What kinds of bonds are found between (1) metal cations and a sea of electrons, metallic bonds (2) between nonmetal atoms, covalent bonds (3) metal and nonmetal ions ionic bonds b. What phases (solid, liquid, gas) conduct electricity in ionic and covalent compounds and in metals? Ionic compounds: liquid (melted) and aqueous solutions conduct electricity Covalent compounds: are poor conductors of electricity Metallic compounds: solid and liquid metals conduct electricity c. Which bond type is found in materials with extremely high boiling points? Ionic compounds (crystal lattice structure) d. What bond type forms molecules? Covalent bonds e. In which bond type are electrons shared Covalent bonds and in which electrons are transferred. Ionic bonds 24. What types of bonds are found in the following compounds? NaCl, NaOH, BaSO 4, LiCl, CaS: are ionic CH 4, CO 2 : are covalent 25. Which group of elements (metals or nonmetals) tends to GAIN electrons when they form ionic bonds? nonmetals Which group of elements (metals or nonmetals) tends to LOSE electrons when they form ionic bonds? metals 3
26. Draw the Lewis Dot Structures for Br 2, H 2 O, and CO 2 The correct dot structure will have the correct number of valence electrons (add up the valence electrons from each element. Br has 7 [group 7A] so Br 2 has a total of 14 valence electrons). It will also have 8 electrons around each atom in the compounds (only 2 for Hydrogen). That is why CO 2 has double bonds. 27. Give the names of the following covalent compounds containing sulfur and oxygen a. SO 2 sulfur dioxide b. S 2 O 3 disulfur trioxide 29. Give the names of the compounds from the following formulas. a. MgCl 2 magnesium chloride b. Mg(NO 2 ) 2 magnesium nitrite 30. Give the formulas for the following compounds. a. Cobalt (III) Sulfate Co 2 (SO 4 ) 3 b. Cobalt (II) Sulfate CoSO 4 c. Cobalt (III) Sulfite Co 2 (SO 3 ) 3 d. Cobalt (II) Sulfite CoSO 3 e. Cobalt (III) Sulfide Co 2 S 3 f. Cobalt (II) Sulfide CoS CHAPTER 10/18 Chemical Reactions, Reaction Rates and Equilibrium 34. Balance the following chemical reactions and then answer the questions that follow. a. 1 C 3 H 8 + 5 O 2 3 CO 2 + 4 H 2 O b. 4 FeS + 7 O 2 2 Fe 2 O 3 + 4 SO 2 c. When a is correctly balanced using the smallest whole number coefficients, what is the sum of these coefficients? 1 + 5 + 3 + 4 = 13 d. When b is correctly balanced using the smallest whole number coefficients, what is the coefficient for SO 2? 4 35. Is this reaction exothermic or endothermic? 2H 2 O + energy H 2 (g) + O 2 (g) endothermic 36. Using the following reaction at equilibrium answer the following questions: A (g) + 3 B (g) 2 C (g) + heat a. What direction will the reaction shift if heat is added? left b. What direction will the reaction shift if pressure is increased? right (smaller # of moles on right side) c. What will happen to the concentration of A if the temperature decreases? right (remove heat) so decrease in A d. What will happen to the concentration of C if more A is added? shift right, so increase in C e. What direction will the reaction shift if more C is added? left 37. What conditions are equal when a reversible chemical reaction is at equilibrium? Rate of forward reaction equals the rate of the reverse reaction 38. How do heat, pressure, concentration and surface area affect the rate of reaction? Heat: increase T, increase rate (increase in molecule motion so more collisions). Pressure: Increase P, increase rate (molecules closer together). Concentration: increase concentration, increase rate (more molecules so more collisions). Surface area: increase surface area, increase rate (increased contact between reactants) 39. How does a catalyst change the rate of a chemical reaction? A catalyst lowers the activation energy which increases the FORWARD reaction rate 4
40. According to the energy diagram to the left, which curve (A or B) represents the catalyzed reaction? Curve B shows a decrease in activation energy caused by the catalyst 41. According to the diagram to the left, is the forward chemical reaction endothermic or exothermic? Exothermic. Products have less energy than reactants, so energy was released. CHAPTER 11 The Mole 42. What is the definition of the mole? What is Avogadro s Number? The mass of which isotope is used to define the mole? One mole is 6.02x10 23 particles. It is the number of particles for which the atomic mass in amu is equal to the molar mass in grams (1 carbon atom has a mass of 12 amu; 1 mole of carbon atoms has a mass of 12 grams). 12 grams of carbon-12 is exactly equal to 6.02x10 23 particles. 43. Be able to calculate the molar mass for a compound. As an example: Calculate the molar masses of Mg(NO 3 ) 2 and HNO 3. Mg(NO 3 ) 2 : 148.3 g HNO 3 : 63.0 g 44. Make the following conversions: a. What is the number of moles of CaS in 120 grams of the compound? 120 g CaS 1 mol = 1.7 mol 72.2 g CaS b. What is the mass in grams of 18.06 10 23 molecules of CO 2? 18.06 x 10 23 molecules 1 mol 44.0 grams = 132 g CO 2 6.02 x 10 23 molecules 1 mol c. How many molecules of CO 2 are in 0.50 moles? 0.5 mol CO 2 6.02x10 23 particles = 3.01x10 23 particles CO 2 1 mol d. How many liters do 7.87 10 23 molecules of dihydrogen sulfide (H 2 S) occupy? 7.87 x 10 23 molecules 1 mol 22.4 L = 29.3 L H 2 S 6.02 x 10 23 molecules 1 mol e. What is the number of liters that 27 g of oxygen gas occupies? 27 g O 2 1 mol 22.4 L = 18.9 L O 2 32.0 g 1 mol 45. a. If 40 grams of magnesium (Mg) are reacted with excess nitric acid (HNO 3 ), how many grams of hydrogen are produced. Mg + 2 HNO 3 Mg(NO 3 ) 2 + H 2 40 g Mg 1 mol Mg 1 mol H 2 2.0 g H 2 24.3 g Mg 1 mol Mg 1 mol H 2 = 3.3 g H 2 b. If 1.7 grams of hydrogen was actually produced, what was the percent yield? (Hint: Use your answer from part a) 1.7 g is the actual yield, 3.3 g is the theoretical yield (from part a). Percent Yield = 5
CHAPTER 12 - Stoichiometry 46. Know how to calculate mole to mole, mole to mass (or vice versa), and mass to mass stoichiometry problems. 47. Given the equation: Zn (s) + 2 HCl (aq) ZnCl 2 (aq) + H 2 (g) a. How many moles of hydrogen will be formed when 4 moles of HCl are reacted? 4 mol HCl 1 mol H 2 = 2 mol H 2 2 mol HCl b. How many moles of Zn will be reacted to form 3 moles of ZnCl 2? 3 mol ZnCl 2 1 mol Zn = 3 mol Zn 1 mol ZnCl 2 c. How many moles of HCl will be needed to react completely with 3 moles of Zn? 3 mol Zn 2 mol HCl = 6 mol HCl 1 mol Zn 48. Given the equation: 2 Na (s) + 2 H 2 O (l) 2 NaOH (aq) + H 2 (g) a. Calculate the number of grams of sodium hydroxide that will be produced if 8.215 grams of sodium are reacted with excess water. 8.215 g Na 1 mol Na 2 mol NaOH 40.0 g NaOH 23.0 g Na 2 mol Na 1 mol NaOH = 14.29 g NaOH b. Calculate the number of grams of water that are needed to produce 6.7 g hydrogen gas. 6.7 g H 2 1 mol H 2 2 mol H 2 O 18.0 g H 2 O = 120.6 g H 2 O 2.0 g H 2 1 mol H 2 1 mol H 2 O 50. Given the equation: 2 CO (g) + O 2 (g) 2 CO 2 (g) a. How many moles of carbon monoxide (CO) are required to produce a total of 44 grams of oxygen (O 2 )? 44 g O 2 1 mol O 2 2 mol CO = 2.75 mol CO 32.0 g O 2 1 mol O 2 b. How many grams of carbon dioxide (CO 2 ) are produced if 3.6 mol O 2 react? 3.6 mol O 2 2 mol CO 2 44.0 g CO 2 1 mol O 2 1 mol CO 2 = 316.8 g CO 2 CHAPTER 13 States of Matter and Kinetic Molecular Theory 51. What are the three common states of matter in terms of particle motion? Solids particles vibrate about a fixed point. Liquids particles are close together and move freely. Gases particles are far apart and move in straight line unless colliding with something. 52. What does STP stand for and what are its numerical values? Standard Temperature and Pressure. Standard T is 0 C or 273K. Standard pressure is 1 atm, 101.3 kpa, 14.7 psi, 760 mm Hg, or 760 torr 53. What are the three assumptions of the Kinetic Theory? 1. All matter is made of small particles. 2. These particles are in constant motion. 3. Collisions between particles are elastic (no loss of energy). 54. At what temperature does all molecular motion theoretically stop? What is the name for this temperature? Zero Kelvin (0 K). Absolute zero. 6
55. Convert the following temperatures from Celsius to Kelvin and Kelvin to Celsius: a. 22 C 295 K c. 220 K -53 C b. 474 C 747 K d. 390 K 117 C 56. Convert the following pressure units to the indicated pressure units: 2.6 atm = 1976 mmhg d. 120 mmhg = 16 kpa 57. The average kinetic energy of the molecules of an ideal gas is measured by what? Temperature in Kelvin CHAPTER 14 Gas Laws R = 0.0821 R = 8.315 At STP for gases only: 1 mole = 22.4 L Dalton s Law of Partial Pressures: P total = P 1 + P 2 + Graham s Law of Diffusion: = Ideal gas law: PV = nrt Combined Gas Law: = 58. There is a mixture of three gases in a closed container. Gas A exerts a pressure of 15 psi, Gas B a pressure of 6.5 psi, and the total pressure is 25.7 psi. What is the partial pressure of Gas C? 25.7 psi 15 psi 6.5 psi = 4.2 psi 59. Which of the following gases would diffuse most rapidly: N 2, O 2, He, Cl 2? He (smallest molar mass) 60. What is the ratio of the rate of diffusion of H 2 to He? = 61. A sample of gas occupies 0.200 ml at STP. Calculate the volume this same amount of gas would occupy at pressure of 0.75 atm and a temperature of 10 C. = V 2 = 0.28 ml 62. If the pressure on 26.0 milliliters of a gas at STP is changed to 0.50 atm at constant temperature, what is the new volume of the gas? 1 atm x 26.0 ml = 0.5 atm x V 2 V 2 = 52 ml 63. The pressure inside the container at 20.0 C was 1.50 atm. What is the pressure of the fluorine gas after its temperature is increased to 40.0 C? (Volume is constant) = P 2 = 1.6 atm 64. If 25.0 L of CO 2 is produced at a temperature of 1100 C and allowed to reach room temperature (25.0 C) without any pressure changes, what is the new volume of the carbon dioxide? 65. Given the reaction at STP: = V 2 = 5.4 L 1 N 2 + 3 H 2 2 NH 3 How many liters do 56 grams of N 2 occupy? 56 g N 2 1 mol N 2 22.4 L 28.0 g N2 1 mol = 44.8 L N 2 OR 56 g N 2 1 mol N 2 28.0 g N2 = 2 moles N 2 Using the ideal gas law (at STP): 1.0 atm x V = 2.0 moles x 0.0821 x 273 K V = 44.8 L How many liters of H 2 will be required to react with 56 grams of N 2? (Use the answer from above: 56 g occupies 44.8L; and the coefficients from the chemical reaction) 44.8 L N 2 3 L H 2 = 134.4 L H 2 1 L N 2 You can use the mole ratios for gases because all gases occupy the 7
same volume at the same temperature. OR You could also do a mole-mole ratio and convert to liters: 2 moles N 2 3 mol H 2 22.4 L 1 mol N2 1 mol 66. Given the reaction at STP: 2 H 2 (g) + O 2 (g) 2 H 2 O (g) What is the total volume water vapor produced if 40 liters of hydrogen gas are consumed in the reaction above? 40 L H 2 2 L H 2 O = 40 L H 2 O 2 L H 2 67. Know that at STP, one mole of any gas occupies a volume of 22.4 L. Perform the following molar volume conversions: a. What is the volume of 3.00 moles of CO 2 gas at STP? = 134.4 L N 2 3 mol CO 2 22.4 L = 67.2 L CO 2 1 mol b. How many moles are equal to 44.8 liters of a N 2 gas at STP? 44.8 L N 2 1 mol = 2.0 mol N 2 22.4 L 68. According to the Combined Gas Law, be able to identify the relationships between variables. For example: a. If temperature is constant, what is the relationship between pressure and volume? If you increase the pressure you decrease the volume (inversely proportional) b. If volume is constant, what is the relationship between pressure and temperature? If you increase the temperature you increase the pressure (directly proportional) c. If pressure is constant, what is the relationship between temperature and volume? If you increase the temperature you increase the volume (directly proportional) 69. Be able to solve ideal gas law problems. If I have an unknown quantity of gas held at a temperature of 753 K in a container with a volume of 21 liters and a pressure of 6.5 atm, how many moles of gas do I have? 6.5 atm x 21 L = n x 0.0821 x 753 K n = 2.2 mol 70. If 8.1 moles of a gas at a pressure of 111.2 kpa have a temperature of 473 K, what is the volume? 111.2 kpa x V = 8.1 mol x 8.314 x 473 K V = 286 L 71. What version of the Combined Gas Law does the graph at the left best represent? (Hint: start with the Combined Gas Law, what is constant?) = If volume is constant this equation becomes: = (P increases as T increases) CHAPTER 15 Solutions and Solubility 72. Know the definition of a solution, solute and solvent. Solution: A homogeneous mixture (i.e., salt water) Solute: The substance dissolved into a solution (i.e., the salt is the solute in salt water) Solvent: The substance into which a solute is dissolved (or the substance that does the dissolving). (i.e., the water is the solvent in salt water) 8
73. What is the concentration (molarity) of a solution that contains 12.5 g NaCl in a 250 ml solution? 250 ml = 0.25 L 12.5 g NaCl 1 mol = 0.214 mol NaCl 58.5 g NaCl Molarity: 0.214 moles NaCl / 0.25 L = 0.86 M CHAPTER 19 Acids and Bases 74. Know the ph scale. ph ranges from 0 to 14 Also, know the ph range for an acid and for a base. 0-6 = acidic; 7 = neutral; 8-14 = basic. 75. If given the hydrogen ion concentration, be able to determine the ph of an aqueous solution. As an example: What is the ph of an aqueous solution with a hydrogen ion concentration of 0.0001 ph = -log[h + ] ph = -log(0.0001) ph = 4 9