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Chemistry Final Exam Review Atomic Theory 1. What is the name of Dalton’s model of the atom? 10. a. Plum pudding model b. Quantum mechanical model c. Nuclear model d. Solid sphere model 2. What is the name of J.J. Thompson’s model of the atom? 11. a. Plum pudding model b. Quantum mechanical model c. Nuclear model d. Solid sphere model 3. What is the name of Schrodinger’s model of the atom? 12. a. Plum pudding model b. Quantum mechanical model c. Nuclear model d. Solid sphere model 4. What is the name of Rutherford’s model of the atom? a. Plum pudding model 13. b. Quantum mechanical model c. Nuclear model d. Solid sphere model 5. Which of the following are results of Rutherford’s model of the atom? a. b. c. d. Dense positively charged nucleus, a lot of empty space Dense negatively charged nucleus, no empty space Dense positively charged nucleus, no empty space Dense negatively charged nucleus, a lot of empty space 6. Which of the following shows the order of correct charge, mass, and location of a proton in an atom? a. +1, 0, outside the nucleus b. +1, 1, outside the nucleus c. +1, 1, inside the nucleus d. +1, 0, inside the nucleus 7. Which of the following shows the order of correct charge, mass, and location of an electron in an atom? a. -1, 0, outside the nucleus b. -1, 1, outside the nucleus c. -1, 1, inside the nucleus d. -1, 0, inside the nucleus 8. How many electrons, protons, and neutrons does a neutral atom of strontium – 89 have? a. e- = 38, p+= 38, n0= 51 b. e- = 38, p+= 51, n0= 51 c. e- = 51, p+= 38, n0= 51 d. e- = 38, p+= 51, n0= 38 9. How many electrons, protons, and neutrons does a neutral atom of fluorine – 19 have? a. e- = 19, p+= 19, n0= 10 b. e- = 19, p+= 9, n0= 9 c. e- = 9, p+= 9, n0= 10 d. e- = 10, p+= 10, n0= 9 14. 15. 16. 17. How many electrons, protons, and neutrons does an ion of 9 Be have? a. e- = 2, p+= 4, n0= 9 b. e- = 2, p+= 4, n0= 5 c. e- = 4, p+= 4, n0= 5 d. e- = 4, p+= 2, n0= 5 How many electrons, protons, and neutrons does an ion of 32 P have? a. e- = 15, p+= 15, n0= 17 b. e- = 15, p+= 17, n0= 15 c. e- = 18, p+= 15, n0= 16 d. e- = 18, p+= 15, n0= 17 When an atom of gallium forms a gallium ion, its charge becomes ______ because it _______ electron(s). a. -3, loses b. +3, gains c. +3, loses d. -3, gains When an atom of bromine forms a bromine ion, its charge becomes _______ because it _______ electron(s). a. -1, gain b. -1, lose c. +1, gain d. +1, lose Which of the following is not a property of a metal? a. Conducts electricity in the solid state b. High melting/boiling point c. Malleable and ductile d. Conducts electricity when dissolved in water Which of the following is not a property of a nonmetal? a. Most solids at room temperature b. Low melting/boiling point c. Brittle solids at room temperature d. Right of the metalloid (semi-metal) “staircase” line Which of the following is a correctly balanced nuclear equation? a. 23892U 23490Th + 42He b. 146C 147N + 42He c. 146C 147N + 0-1 Mg d. 23892U 23490Th + 42β Which of the following represents alpha decay of radon – 222? a. 22286Rn 21884Po + 42He b. 86222Rn 82220Th + 42He c. 22688Ra 22286Rn + 42He d. 22286Rn 22287Fr + 0-1β 1. What is the mass of 2 moles of propane gas, C3H8? a. 11 grams b. 44 grams c. 22 grams d. 88 grams 2. How many moles of propane gas, C3H8, are contained in 11 grams of C3H8? a. 11 × 1023 moles b. 4 moles c. 1.5 × 1023 moles d. 0.25 moles 3. What is the mass in grams of 3 moles of water molecules, H2O? a. 54 grams b. 0.166 grams c. 6 grams d. 21 grams 4. How many moles of water molecules, H2O, are present in a 27 gram sample of water? a. 9 × 1023 moles b. 1.5 moles c. 2 moles d. 2/3 mole 5. What is the mass of 10 moles of ammonia, NH3? a. 1.7 grams b. 27 grams c. 170 grams d. 0.587 grams 6. How many moles of methane, CH4, are in 80 grams of methane? a. 6.022 × 1080 moles b. 5 moles c. 80 × 1023 moles d. 0.201 moles MOLES 7. How many molecules are contained in 3 moles of water, H2O? a. 6 molecules b. 54 molecules c. 1.8 × 1024 molecules d. 3 × 1023 molecules 8. A sample of carbon dioxide gas (CO2) contains 6.022× 1023 molecules. How many moles of carbon dioxide does this represent? a. 1 mole b. 440 moles c. 44 moles d. 10 moles 9. How many molecules of ethane gas, C2H6, are in 15 grams of the compound? a. 0.5 moles b. 2 moles c. 3 × 1023 moles d. 45 moles 10. What is the mass, in grams, of 3 × 1023 atoms of helium? a. 2 grams b. 1.2 × 1024 grams c. 3 × 1023 grams d. 8 grams 11. Approximately how many atoms of carbon are present in a 120 gram sample of carbon? a. 10 atoms b. 6 × 1022 atoms c. 6 × 1024 d. 1440 atoms 1 mole of a compound/element = _________________________ particles 1 mole of a compound/element = _________________________ in grams Molar mass of a compound = 6.022× 1023 particles ELECTRON CONFIGURATIONS & LIGHT 1. What element has the noble gas configuration [Ne]3s23p1? __________________ 2. What element has the electron configuration notation 1s22s22p63s1? ___________ 3. Which of the following is the correct noble-gas notation for the element strontium? a. [Kr]5s1 b. [Xe]5s2 c. [Kr]6s2 d. [Kr]5s2 4. The above orbital notation is used to represent which element? a. Boron b. Sulfur c. Oxygen d. fluorine 5. The above orbital notation is used to represent which element? a. Phosphorus b. Arsenic c. Nitrogen d. Silicon 6. Which of the following is the correct configuration notation for the element titanium (Ti)? a. 1s22s22p63s23p64s23d2 b. 1s22s22p63s23p63d24s2 c. 1s22s22p63s23p64s24d2 d. 1s22s22p63s23p64s21d2 7. Which of the following is the correct electron configuration notation for the element nitrogen, N? a. 1s22s2 b. 1s22s22p6 c. 1s22s22p3 d. 1s22s21p3 8. Which of the following is the correct electron configuration for an aluminum ion? a. 1s22s22p63s23p1 b. 1s22s22p63s23p6 c. 1s22s22p6 d. 1s22s22p63s2 9. Which of the following electron dot notations is correct for the element phosphorus when it is in the ground state? a. I b. II c. V d. III 10. Which of the following electron dot notations is correct for the element oxygen when it is in the ground state? 11. 12. 13. 14. 15. 16. a. I b. II c. III d. V Which of the following elements has the same number of valence electrons as the element sodium? a. Ar b. Cs c. Ca d. Mg Which of the following elements has the same number of valence electrons as the element selenium? a. Fe b. K c. P d. O Which of the following elements will have similar physical and chemical properties as lithium? a. Rubidium b. Carbon c. Nitrogen d. neon Which of the following elements will have similar physical and chemical properties as Iodine? a. Te b. Xe c. F d. Po When an electron gets excited and goes up an energy level, it has ____________ energy and when an electron goes down an energy level, it has ____________ energy. a. Absorbed, released b. Released, absorbed c. Ground, excited d. Excited, absorbed When an electron transitions from n = 3 to n = 4, it has ____________. a. Absorbed energy b. Released energy c. Excited state d. Ground state 17. When an electron transitions from n = 6 to n = 2, it has ___________. a. Absorbed energy b. Released energy c. Ultraviolet radiation d. Infrared radiation 18. What color light is produced when an electron transitions from the n = 6 to n = 2 energy level? a. Blue b. Red c. No color d. violet 19. Given the representation of a chlorine atom, which circle might represent an atom of bromine? a. b. c. d. 20. Circle B Circle D None of these Circle C Given the representation of a chlorine atom, which circle might represent an atom of sulfur? a. b. c. d. None of these Circle B Circle D Circle C 21. As one moves from left to right ( → ) within a period across the periodic table, the electronegativity of the elements encountered tends to: a. stay the same b. increase c. decrease 22. The elements with the largest atomic radii are found in the: a. b. c. d. lower right-hand corner of the periodic table lower left-hand corner of the periodic table upper right-hand corner of the periodic table upper left-hand corner of the periodic table 23. Of the following elements, which one would have the largest radius? a. b. c. d. Cesium (Cs, atomic #55) Potassium (K, atomic #19) Hydrogen (H, atomic #1) Sodium (Na, atomic #11) 24. The energy required to remove an electron from an atom is known as: a. b. c. d. radioactivity electron affinity ionization energy electronegativity 25. Of the following elements, which one would have the largest ionization energy? a. b. c. d. Cesium (Cs, atomic #55) Hydrogen (H, atomic #1) Sodium (Na, atomic #11) Potassium (K, atomic #19) BONDING AND INTERMOLECULAR FORCES 1. Which of the following gases does not exist in a. 1 nature as a diatomic molecule? b. 3 a. Nitrogen c. 4 b. Helium d. 2 c. Hydrogen 11. In the correct Lewis structure for the methane d. oxygen molecule, how many unshared electron pairs 2. Ionic compounds generally form: surround the carbon? a. Liquids a. 2 b. Gases b. 0 c. Crystals c. 8 d. molecules d. 4 3. In metallic bonding, the valence electrons of all 12. In nonpolar covalent bonds, valence electrons are atoms are shared in: a. Equally shared a. A nonpolar covalent bond b. Unequally shared b. An electron sea c. Destroyed c. A polar covalent bond d. transferred d. Transferred to metallic ions 13. Which of the following is an acceptable Lewis 4. The metalloids possess properties of metals and structure for chloromethane (CH3Cl)? nonmetals and are also known as a. Semimetals b. Halogens c. Gases d. Liquids 5. The seven elements that occur as diatomic a. elements are a. H2,N2,O2,He2,Ne2,C2,Na2 b. H2,N2,O2,He2,Ne2,Cl2,Br2 c. H2,N2,O2,F2,I2,Cl2,Br2 d. Fe2,Rn2,O2,He2,Ne2,C2,Br2 6. The bond between sodium and oxygen is expected to be b. a. Gaseous b. Nonpolar covalent c. Ionic d. Polar covalent 7. When compared to single bonds, double bonds are generally c. a. Shorter and stronger b. Longer and stronger c. Longer and weaker d. Shorter and weaker 8. The bond between lithium and fluorine is a. Polar covalent d. b. Ionic 14. In a diatomic molecule of an element, the bond c. Nonpolar covalent between the atoms must be d. metallic a. Nonpolar covalent 9. In the ionic compound magnesium fluoride, what b. Polar covalent is the ratio of the two elements necessary so that c. Metallic each element obtains its octet from the transfer of d. ionic electrons? 15. In polar covalent bonds, valence electrons are a. 3 magnesium: 1 fluorine a. transferred b. 1 magnesium: 1 fluorine b. unequally shared c. 2 magnesium: 1 fluorine c. destroyed d. 1 magnesium: 2 fluorine d. equally shared 10. In the correct Lewis structure for water, how many unshared pairs of electrons will oxygen 16. In ionic bonds, valence electrons are have? a. Equally shared b. Transferred c. Unequally shared d. destroyed 17. How many atoms are needed to provide the electrons necessary to complete the valence octet of an oxygen atom? a. Three sodium atoms b. Two sodium atoms c. Four sodium atoms d. One sodium atom 18. The measure of the attraction that an atom has for electrons involved in chemical bonds is known as a. Ionization energy b. Radioactivity c. Electronegativity d. Electron affinity 19. Which of the following is the correct Lewis structure for ammonia? 22. Which of the following is the correct Lewis structure for formaldehyde, CH2O a. b. c. d. 23. Which of the following is the correct Lewis structure for phosphorus tribromide? a. b. c. d. 20. In drawing Lewis structures, a single line (single bond) between two elements represents a. An octet of electrons b. An unshared pair of electrons c. A shared electron d. A shared pair of electrons 21. Which of the following is a correct Lewis structure for hydrogen cyanide, HCN? a. b. c. d. a. b. c. d. 24. Which of the diatomic elements has a double bond between its atoms? a. Fluorine b. Nitrogen c. Oxygen d. Hydrogen NOMENCLATURE, EMPIRICAL AND MOLECULAR FORMULAS 1. What is the correct name for ClO2? c. Copper nitrogen a. Chlorine dioxide d. Copper (I) nitride b. Monochlorine dioxide 12. The correct formula when calcium and nitrogen c. Dichlorine monoxide bond together would be d. Chlorine oxide a. Ca2N3 2. What is the name of SiCl4? b. Ca3NO3 a. Silicon chloride c. Ca3(NO3)2 b. Monosilicon tetrachloride d. Ca3N2 c. Silicon tetroxide 13. The correct formula when nickel (II) and bromine d. Silicon tetrachloride combine would be 3. What is the correct formula for nitrogen a. NiBr monoxide? b. Ni2Br a. NO2 c. NiBr2 b. No d. NiB2 c. NO 14. The formula that would be made when aluminum d. N2O and the phosphate ion bond would be 4. What is the correct formula for the compound a. AlPO4 tetraphosphorus trisulfide? b. Al3(PO4)3 a. P3S4 c. Al3(PO4) b. P5S4 d. Al(PO4)3 c. 4PS2 15. If an element in group 1 was to combine with an d. P4S3 element in group 7 (17), the resulting compound 5. What is the correct name of CO? would have a ratio of: a. Monocarbon monoxide a. 1 : 1 b. Carbon dioxide b. 1 : 2 c. Carbon monoxide c. 2 : 1 d. Carbon oxide d. 3 : 3 6. What is the correct formula for boron trifluoride? 16. If an element in group 2 was to combine with an a. B3F element in group 5 (15), the resulting compound b. BF3 would have a ratio of: c. 3BF a. 2 : 3 d. B3F3 b. 3 : 2 7. What is the correct formula for carbon c. 2 : 5 tetrabromide? d. 5 : 2 a. CB4 17. If a molecular formula for some compound is b. C4Br X2Y6, a possible empirical formula could be c. C4Br a. XY2 d. CBr4 b. X2Y3 8. What is the correct formula for dinitrogen c. XY3 monoxide? d. X3Y a. N2O2 18. What is a possible molecular formula for a b. N2O compound with the empirical formula of CH3? c. NO2 a. C2H3 d. 2NO b. C3H6 9. The correct name for the compound Fe2S3 is c. C3H a. Diiron trisulfide d. C3H9 b. Iron (III) sulfide 19. Analysis shows that some compound has the c. Iron sulfide following percent composition: 40.05% S and d. Iron (II) sulfide 59.95 % O. What is its empirical formula? 10. The correct formula for chromium (III) oxide is a. S3O a. Cr3O3 b. SO3 b. Cr3O2 c. SO c. Cr2O2 d. S3O6 d. Cr2O3 11. The correct name for the compound Cu3N is a. Copper (II) nitrate b. Copper (III) nitride 20. Propane is a hydrocarbon. it is 81.82 % carbon and 18.18 % hydrogen. What is the empirical formula? a. CH8 b. C8H3 c. CH3 d. C3H8 21. Analysis shows that some compound has the following percent composition: 48.64% C, 8.16 % H, and 43.20 % O. What is its empirical formula? a. C1.5H3O b. C2H3O c. CHO d. C3H6O2 22. The hydrocarbon used in manufacture of foam plastics is called styrene. Analysis of styrene indicates the compound is 92.25 % C and 7.75 % H and has a molar mass of 104 g/mol. What is the molecular formula for styrene? a. C8H b. CH c. C8H8 d. CH4 23. A colorless liquid composed of 46.68 % N and 53.32 % O has a molecular mass of 60.01 g/mol. What is the molecular formula? a. NO b. N2O c. N2O2 d. N4O2 24. The empirical formula of a compound is C3H3O. The molecular mass is 110.0 g/mol. What is its molecular formula? a. C3H3O b. CHO3 c. C6H6O d. C6H6O2 25. The empirical formula of a compound is PNCl2. The molecular mass is 695 g/mol. What is its molecular formula? 26. Calculate the % composition of sodium sulfate. 27. What is the percent composition of H3PO4? 28. Determine the % by mass (% composition) of sucrose, C12H22O11. a. 42.10 % C, 6.480 % O, 51.42 % H b. 58.23 % C, 7.65 % O, 34.12 % H c. 44.0 % C, 6.75 % H, 53.98 % O d. 42.10 % C, 6.480 % H, 51.42 % O Define these terms: Single replacement reaction Combination reaction Decomposition reaction Precipitate Reactants Products Coefficient Chemical equation Balanced chemical equation CHEMICAL REACTIONS Double replacement reaction Diatomic molecule Catalyst 1. What are the 5 types of chemical reactions? 2. Know how to identify the different types of chemical reactions. Examples: a) FeCl3 + NaOH → Fe(OH)3 + NaCl b) Al + O2 → Al2O3 c) C2H2 + O2 → CO2 + H2O d) Na + H2O → NaOH + H2 e) KClO3 → KCl + O2 3. Know how to balance equations. Example: a) __ FeCl3 + __ NaOH → __ Fe(OH)3 + __ NaCl b) __ Al + __ O2 → __ Al2O3 c) __ C2H2 + __ O2 → __ CO2 + __ H2O d) __ Na + __ H2O → __ NaOH + __ H2 e) __ KClO3 → __ KCl + __ O2 Yield sign Activity series Law of conservation of mass Subscript Combustion reaction Aqueous 4. Know how to predict products as in the problems below. Example: a) Al + N2 → b) H2O → c) Ca + H2O → d) Cl2 + NaBr → e) FeS + HCl → 5. What is an activity series chart? What type of reaction do you use it for? a) Using the activity chart, why can sodium replace hydrogen? 6. What are 5 indicators/observations of a chemical reaction? 7. List the chemical formulas for the 7 diatomic molecules. 8. Know how to translate chemical equations and balance them appropriately. Example: a) ammonium chloride reacts with calcium hydroxide to form calcium chloride and nitrogen trihydride (ammonia) and water b) sodium oxide and water yield sodium hydroxide STOICHIOMETRY Mole ratio 1. What do the coefficients mean in a chemical equation? 2. Know how to calculate the mole ratio between reactants and products in a chemical formula. a) What is the mole ratio for calcium and oxygen in 2Ca + O2 → 2CaO 3. Know how to solve mole to mole, mole to mass, mass to mole, mass to mass problems. a) How many moles of lithium hydroxide are required to react with 20. mol of carbon dioxide? CO2 + 2LiOH → Li2CO3 + H2O b) What mass, in grams, of glucose is produces when 3.00 mol of water react with carbon dioxide? 6CO2 + 6H2O → C6H12O6 + 6O2 c) How many moles of NO are formed when 824 g of ammonia reacts with an excess of oxygen? (balance the equation first) NH3 + O2 → NO + H2O d) How many grams of SnF2 are produced from the reaction of 30.0 g HF with Sn? Sn + 2HF → SnF2 + H2 GASES Ideal gas law directly proportional inversely proportional molar volume STP Ideal gas constant Partial pressure Dalton’s law kinetic molecular theory elastic collision law of combining volumes 1. Know the 5 assumptions of the kinetic molecular d) The volume of a sample of oxygen gas is 300.0 ml theory. when the pressure is 1.00 atm and the temperature is 2. Know the difference between an ideal gas and a real 27.0 C. At what temperature would the volume gas. change to 1.00 L and the pressure change to 0.500 3. Explain a gas based on the following properties: atm? density, compressibility, diffusion, effusion, fluidity, e) A sample of gas at 25.0 C has a volume of 11.0 L and shape , IMF, particle arrangement, and volume exerts a pressure of 660.0 mmHg. How many moles expansion. of gas are in the sample? 4. Define pressure. What are some common pressure f) A sample of gas in a closed container at a temperature units? of 100. C and a pressure of 3.0 atm is heated to 300. C. 5. Know how to convert pressure units: What pressure does the gas exert at the higher a) convert .200 atm to mmHg temperature? b) convert 345.8 kPa to atm 8. Use the law of combining volumes, Avogadro’s law, c) convert 760 mmHg to kPa and molar volume to solve 6. What is standard temperature and standard pressure? these problems. 7. Know how to solve problems using Boyle’s law, a) 3O2 → 2O3 Both gases are measured at the same Charles law, Gay-Lussac, Combined, Ideal, Density temperature and pressure. How many liters of O2 are and Molar mass using the Ideal gas law and Dalton’s required to make 24 L of O3 ? law of partial pressure. b) How many liters of O3 are formed from 12 mol of O2 a) A gas occupies a volume of 200. ml at 100. mmHg. at STP? What volume will the gas occupy at 300. mmHg? 11. Know these answers: b) Air has a total pressure of 20.6 atm and contains a) As the temperature of a gas decreases, the volume of a carbon monoxide, oxygen, and nitrogen. If air is made gas will ____________. up of 0.6 atm of carbon monoxide, 12.6 atm of b) As the temperature of a gas decreases, the pressure of oxygen, what would be the partial pressure of the gas will ____________. nitrogen? c) As the volume of the gas decreases, the pressure of the c) If a sample of gas occupies 15.9 L at 34 C, what will its gas will ____________. volume be at 27 C if the pressure does not change? Know these terms Vaporization Freezing point Freezing Endothermic reaction temperature THERMOCHEMISTRY/SOLIDS & LIQUIDS Condensation Sublimation Deposition Exothermic reaction heat 1. State the 6 phase changes of state and which ones work in opposition to each other. i.e. sublimation and deposition 2. Explain how a solid melts into a liquid using kinetic energy in your explanation. 3. What 2 temperatures measure the same amount during a phase change of a liquid pure solvent to a solid? 4. Know how to read phase diagrams. Sketch a quick diagram locating the triple point, critical point, the melting point /freezing point line and the boiling point/condensation point line. Also label the 3 sections as solid , liquid, and gas. 5. Know how to read a heating and cooling curve. What do the plateaus tells you? What do the slopes tell you? Where is the KE of the substance constant? 6. Sketch an endothermic reaction graph, labeling the reactants, products, activation energy, activated complex, and the heat of reaction. 7. What is the sign of an endothermic reaction and exothermic reaction? 8. Using the specific heat values for water and iron, which one would have the largest temperature change if they have the same mass? Evaporation Melting point Triple point Melting Phase Diagram Boiling Heating &Cooling Curve Specific heat capacity 9. Know how to calculate the heat released or absorbed during a physical change. a. Calculate the heat absorbed when 15.0 g of ice melts to liquid. See reference sheet for Hfus b. Calculate the heat released when 75.4 g of vapor condenses into liquid. See reference sheet for Hvap 10. Know how to calculate the heat released or absorbed in a chemical reaction? a) What is the specific heat of a metal that releases 2500 J of energy. The metal has a mass of 25 g and had a temperature change of 5C. b) How much heat is released when iron is dropped in a beaker of water. The mass of the metal was 43 g and the initial temperature of the metal was 78 C. The water temperature changed from 25 C to 32 C. The specific heat of the metal is .45J/gC. c) What is the amount of heat absorbed by water if 23.4 g of water is heated from 34C to 78 C. See reference sheet for specific heat of water. Know these terms Rate Collision theory Transition state KINETICS & EQUILIBRIUM Activation energy Activated complex Catalyst 1. Explain the three criteria of the collision theory. 2. On the pathway below, label the activated complex, activation energy with catalyst, and activation energy without catalyst 3 What are the five factors that affect the rate of a reaction? 4. Which of the five factors change collision frequency? 5 Which factor changes collision frequency and the energy of the collisions? 6. How does rate change if you increase the concentration of the reactants? 7. How does rate change if you increase the surface area? 8. How does rate change if you decrease the temperature? 9. How does rate change if you add a catalyst? 10. Write the equilibrium expression for the following reaction. a) H2(g) + Cl2(g) 2HCl(g) + heat 11. In the process of chemical equilibrium, what stays constant at equilibrium? 12. In the process of equilibrium, are the rates equal to each other? 13 Using the reaction above, answer the following questions regarding Le Chatelier’s principle. a) Which direction does the reaction shift if temperature increases? b) Which direction does the reaction shift if hydrogen gas is increased? c) Which direction does the reaction shift if HCl is removed? d) Which direction does the reaction shift if the volume is decreased? e) Which direction doe the reaction shift if temperature is decreased? 14. If K = .00045, what side of the reaction will be favored? SOLUTIONS Know these terms: Solution solute soluble insoluble miscible immiscible electrolytes (strong and weak) non-electrolytes solubility solvent supersaturated solution aqueous solution unsaturated solution saturated solution Henry’s law molarity Know the following: 1. Explain the like dissolves like rule and give an example following the rule. 2. Name 3 factors that increase the rate of dissolution of a substance. 3. Describe solution equilibrium. 4. Name substances that are considered electrolytes and non-electrolytes. 5. What is the effect of temperature and pressure on gas solubility? 6. What is the effect of temperature on the solubility for most ionic solids? 7. Know how to calculate molarity a) What is the molarity of 4.5 moles of Ba(OH)2 in 10.0 L? b) A solution has a molarity of 2.8 M and a volume of 250 ml. How many moles of solute are in the solution? 8. Know how to read a solubility graphs. a. Using the solubility graph from the notes, how much of NaCl can be dissolved at 45C b. Using the solubility graph from the notes, 50 g of KClO3 is dissolved in 100 g of water at 45C. Is the solution saturated or unsaturated? 9. Know how to solve dilution problems. a) How many ml of a 2.0 M NaBr solution are needed to make 200 ml of a 0.50 M solution? 10. Which types of substances produce electrolytes? 11. Which type of substances produce non electrolytes? ACIDS & BASES Know these terms: Arrhenius acid Arrhenius base Bronsted-Lowry acid Bronsted-Lowry base pH conjugate acid Conjugate base hydroxide ion hydronium ion neutralization titration equivalence point 1. List some common properties of an acid. 2. List some common properties of a base. 3. Define self-ionization of water. 4. Know how to predict the products and balance neutralization (double replacement) reactions. a) H2CO3 + Fe(OH)3 → 5. Know how to calculate the pH from hydrogen and hydroxide ion concentrations a) What is the pH of a [OH-] = 1 x 10-5 M? b) What is the pH of a [H+] = 1 x 10-5 M? c) What is the pOH of a [H+] = 1 x 10-1 M? d) What is the pOH of a [OH-] = 1 x 10-12 M? 6. What is the hydrogen ion concentration of 0.001 M HNO3? What is the [OH-]? 7. What is the hydrogen ion concentration of [OH-] = 3.0 x 10-2 M? What is the pH? 8. What is the pH of a solution if the [H+] = 3.4 x 10-5 M? What is the hydroxide concentration? 9. Determine the pH of a 2.0 x 10-2 M Sr(OH)2? 10. The pH of a solution is measured and determined to be 7.52? What is the hydrogen ion concentration? Is the solution acidic or basic? 11. Know how to look at an equation and predict Bronsted-Lowery acids and bases and conjugate acids and conjugate bases. a) NH4+ + H2O → NH3 + H3O+ What is the base? What is the conjugate base? What is the acid? What is the conjugate acid? 12. What are the products of neutralization? 13. Know how to name acids and bases a) HF b) H2SO4 c) NaOH d) HNO2 e) Fe(OH)2 14. In a titration, how much of .15 M NaOH is needed to neutralize 20 ml of .500M HCl solution? HCl + NaOH H2O + NaCl 15. In a titration, what is the molarity of HNO3 if 25 ml of it neutralized 15 ml of .60M Ca(OH)2 2 HNO3 + Ca(OH)2 2 H2O + Ca(NO3)2 16. What is the difference between end point and equivalence point?